Basic II-Chemistry (BIO102) book HANDOUTS in pdf

<b><span style="font-size: large;">Basic II-Chemistry (BIO102) book HANDOUTS in pdf</span></b><br />Handouts | Lectures | Contents | Books<div> <br /></div> <div> <a href="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEja-2uxqLK0E-EzAwx4kZsCzgg3fA5S64HEo90qmlFMB4Y5o7d1gAdxOhLSsRBnqRVDXC8yoeqPG8O0LADzCrtq3l9OGsjFvrTRJ_L37vRnCiqdnjRm0sYPGKLFmpoQEKuZYdpDj0oYbgQ/s1600/9781256328278-us.jpg" imageanchor="1"><img border="0" height="640" src="https://blogger.googleusercontent.com/img/b/R29vZ2xl/AVvXsEja-2uxqLK0E-EzAwx4kZsCzgg3fA5S64HEo90qmlFMB4Y5o7d1gAdxOhLSsRBnqRVDXC8yoeqPG8O0LADzCrtq3l9OGsjFvrTRJ_L37vRnCiqdnjRm0sYPGKLFmpoQEKuZYdpDj0oYbgQ/s640/9781256328278-us.jpg" width="538" /></a></div> <div> <br /></div> <div> <br /></div> <div> <div> ©Copy right Virtual University of Pakistan</div> <div> 1</div> <div> Lecture 1</div> <div> Atomic Structure</div> <div> Man has known the fact that matter is composed of very very small particles since ancient times.</div> <div> It is generally accepted that the concept that matter is discontinuous and composed of tiny</div> <div> particles was first given by the Greek philosophers of the fifth century BC, Leucippos and</div> <div> Democritos. It is notable that Plato and Aristotle did not believe in the existence of atoms and</div> <div> regarded the matter as continuous. Later in 100 BC a Roman poet Lucretius illustrated the idea of</div> <div> atom by his poetry. The word atom is from Greek words ‘a’ meaning ‘not’ and ‘tomos’ meaning</div> <div> divisible. Thus, the meaning of atom is the particle that is indivisible.</div> <div> In 1808, John Dalton (1766-1844), an English schoolteacher put forward the concept of atom on</div> <div> the scientific grounds. He proposed that the elements of atoms that are indivisible.</div> <div> Today, the atom is defined as the smallest particle of an element that can enter into chemical</div> <div> reactions.</div> <div> Atom is not indivisible. It is, in fact, a tiny organized system consisting of smaller particles. These</div> <div> smaller particles that make up the atom are called subatomic particles. Subatomic particles of</div> <div> fundamental importance are electrons, protons and neutrons. These particles combine together in</div> <div> varying number to form atoms of different elements.</div> <div> The number of protons in its atoms is the characteristic property of an element. All the atoms of</div> <div> an element have the same number of protons. The atoms of different elements have different</div> <div> number of protons. A neutral atom has as many electrons as protons. Number of neutrons is not</div> <div> fixed. Generally, it is equal to the number of protons in the lighter atoms, and much higher than</div> <div> the number of protons in the heavier atoms. Atoms of the same element can have different</div> <div> number of neutrons. Such atoms are called isotopes. Most elements have two or more isotopes.</div> <div> DISCOVERY OF SUBATOMIC PARTICLES</div> <div> Now, the question is how were electrons, protons, and neutrons discovered? There are a number</div> <div> of experiments that ultimately led to the discovery of these particles and their arrangement in the</div> <div> atom. These experiments include:</div> <div> · Passage of electricity through solutions</div> <div> · Discharge tube experiments</div> <div> · Radioactivity</div> <div> · Chadwick’s experiment</div> <div> · Spectroscopic studies</div> <div> PASSAGE OF ELECTRICITY THROUGH SOLUTION –ELECTROLYSIS</div> <div> There are numerous substances that pass electricity in solution or in their molten state. They are</div> <div> called electrolytes, and the process is called electrolysis. Electrolysis was thoroughly studied by</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 2</div> <div> an English scientist Michael Faraday (1791-1867), who in 1832-1833 found out the relationship</div> <div> between the quantity of electricity passed through the solution and the quantity of matter</div> <div> deposited at the electrodes as a result of the passage of electricity, and put forward his two famous</div> <div> laws.</div> <div> Faraday’s experimental results led George John Stone Stone Stoney in 1874 to conclude</div> <div> that electricity like matter is discontinuous and consists of particles. In 1891, he gave the name</div> <div> electron to this particle of electricity. (Electron is a Greek word for ‘amber’ which is electrified</div> <div> when rubbed with wool).</div> <div> Discharge Tube Experiments</div> <div> (Crookes tube Experiments)</div> <div> Discharge tube experiments provided strong evidence for the existence of subatomic particles. A</div> <div> discharge tube is a glass tube having two electrodes sealed in at each end. It is connected to a high</div> <div> voltage battery to provide required voltage and to a vacuum pump to evacuate air or gas from the</div> <div> tube.</div> <div> Since William Crookes, a British physicist, was the first of several scientists to construct</div> <div> discharge tubes, these are also known as Crookes tubes.</div> <div> Working: At ordinary pressure, electricity does not pass through the tube since gases are nonconductor</div> <div> of electricity. As the gas or air from the tube is gradually evacuated with the vacuum</div> <div> pump, pressure decreases in the tube; a number of interesting phenomena are observed one after</div> <div> the other.</div> <div> When the pressure is reduced to a very small value, around 0.01-0.001 torr, an electric</div> <div> discharge takes place producing a uniform glow inside the tube. Experiments have shown that it is</div> <div> due to a radiation emitted from the cathode, which travels towards the anode. Because the rays are</div> <div> emitted from cathode they are given the name cathode rays. Now the question is what are these</div> <div> cathode rays? The cathode rays are the streams of electrons emitted from the cathode. All</div> <div> electrons have the same charge and mass; they all are identical.</div> <div> CONCLUSION</div> <div> All atoms contain smaller particles called electrons. All electrons are identical. An</div> <div> electron is a negatively charged particle. Its charge is –1.6 x 10-19 coulombs. An electron is an</div> <div> extremely small material particle with a mass of 9.1 x 10-28 grams.</div> <div> DISCOVERY OF PROTON</div> <div> In 1886, E. Goldstein, a German physicist discovered that in addition to cathode rays,</div> <div> another type of rays is also produced in the discharge tube which travel in the direction opposite</div> <div> to that of cathode rays. The rays were given the name canal rays as they passed through the canals</div> <div> of the cathode. These rays were made up of positively charged particles. These particles are</div> <div> formed due to the ionization of the atoms or molecules of the gas present in the tube. In the</div> <div> discharge tube, cathode rays, the electrons, collide with the atoms or molecules of the gas in the</div> <div> tube and knock out one or more electrons from these atoms or molecules leaving them with a</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 3</div> <div> positive charge. When an atom or molecule loses one or more of its electrons, it acquires a net</div> <div> positive charge equal to the number of electrons lost. For example, hydrogen and helium may</div> <div> produce H+2 and He+ ions. They may also produce H+ and He2+ ions.</div> <div> Thus, positive rays are the stream of positive ions of the gas present in the tube that are produced</div> <div> due to the ionization of their atoms or molecules.</div> <div> Out of all positive ions obtained from different gases, the lightest ion was that of hydrogen. This</div> <div> ion, H+, was named as proton (Greek protos, ‘first’) by Rutherford in 1920. It carries a charge</div> <div> equal to that of an electron but opposite in sign. The ions produced from all gases in the discharge</div> <div> tube have been found to have charge that is integral multiple of the charge of a proton (H+ ion). It</div> <div> shows that atoms contain one or more positively charged particles, the protons.</div> <div> J.J. THOMSON’S MODEL OF ATOM</div> <div> By the end of the nineteenth century it became all clear that atoms consist of negatively charged</div> <div> electrons and a positively charged part, but it was not clear how atoms are constructed. In 1898,</div> <div> J.J. Thomson suggested that an atom might be a positively charged sphere in which negatively</div> <div> charged electrons are embedded. This model of the atom is sometimes called the “plum pudding”</div> <div> model. Of course, it was based on insufficient information and soon found to be incorrect.</div> <div> Today, it is known that neutron is a neutral particle with mass 1.008665 amu or 1.675x10-24 g.</div> <div> The nucleus of every atom contains neutrons alongwith protons except ordinary hydrogen, which</div> <div> has no neutron.</div> <div> Rutherford’s Experiment</div> <div> In 1909, Ernest Rutherford, with Hans Geiger and Ernest Marsden, performed a very important</div> <div> experiment. In their experiment they shot a stream of high-speed a-particles at very thin sheets of</div> <div> various materials such as gold. They observed that most of the a-particles (about 99.9%) passed</div> <div> through the sheet without any deflection. However, some particles deflected through various</div> <div> angles, and a few almost completely reversed their direction.</div> <div> Rutherford’s atomic model: To explain the results of his experiment, Rutherford in 1911 proposed</div> <div> an atomic model. The main points of his model are as follows:</div> <div> An atom consists of a nucleus which contains almost all mass and all positive charge of the atom.</div> <div> Around the nucleus, there are electrons which revolve around the nucleus just as the planets</div> <div> revolve around the sun.</div> <div> Light</div> <div> Atomic spectra have greatly helped us in understanding the electronic structure of the atom. The</div> <div> light emitted from an atom gives us information about the electronic structure of that atom.</div> <div> Light has many forms such as visible or white light, ultraviolet light, infrared light, X-rays and yrays</div> <div> etc. all these forms of light propagate through space as electromagnetic waves. Like all other</div> <div> waves, electromagnetic waves are described in terms of their velocity, frequency, wavelength,</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 4</div> <div> and amplitude. The wavelength, λ(lambda), is the distance between two successive crests or</div> <div> troughs of the wave. The frequency, v (nu), is the number of wave crests or troughs that pass a</div> <div> given point in one second. Light travels with a constant speed of 3.00 x 108 m/s, c. Frequency is</div> <div> inversely proportional to wavelength at constant velocity. Thus,</div> <div> v = c/λ</div> <div> All electromagnetic waves have the same speed in space but their wavelengths and frequencies</div> <div> vary. The wavelength of the visible or white light is about 400 to 750 nm, (4000 to 7500A0)</div> <div> Continuous spectrum</div> <div> A spectrum is a display of component colours (or wavelengths) of radiation on a screen or</div> <div> photographic plate. Visible light is composed of a mixture of electromagnetic waves of all</div> <div> wavelengths (frequencies) ranging from about 400 nm to 750 nm. When a narrow beam of visible</div> <div> light is passed through a glass prism it is spread out into a band of colours ranging from longwavelength</div> <div> red light to short-wavelength violet light. A prism refracts light to different</div> <div> wavelengths through different angles and thus disperses the beam of light into its component</div> <div> wavelengths to form a continuous spectrum. A continuous spectrum contains all the wavelengths</div> <div> present in that range. All the colours are present in continuity and there is no demarcation or dark</div> <div> band between them.</div> <div> Atomic Spectrum</div> <div> When an element in its atomic form is heated, it emits radiation. If a beam of light emitted by a</div> <div> sample of an excited element is passed through a prism, a spectrum containing only a few lines</div> <div> (wavelengths) is obtained. This spectrum is known as a line spectrum. Since the light is emitted</div> <div> by atoms of an element, the spectrum is also called an atomic emission spectrum (or simply</div> <div> emission spectrum or atomic spectrum). The dark regions of the spectrum correspond to the</div> <div> wavelengths for which light is not emitted. On the other hand, if we pass a beam of white light</div> <div> through a sample of unexcited gas (any element in gaseous state) and the transmitted light is</div> <div> passed through a prism, the spectrum obtained is called an absorption spectrum. The absorption</div> <div> spectrum of an element is also a line spectrum. It is the photographic negative of the emission</div> <div> spectrum. The absorption spectrum of an element lacks those wavelengths that are present in its</div> <div> emission spectrum.</div> <div> Significantly, each element displays a characteristic set of lines in its emission or absorption</div> <div> spectrum. An element emits the same wavelengths that it absorbs. Thus, the emission or</div> <div> absorption spectra can serve as “fingerprints” on the basis of which the elements can be</div> <div> identified.</div> <div> Atomic Spectra of Hydrogen</div> <div> Hydrogen is the simplest element. A number of scientists studied the emission of radiation from</div> <div> hydrogen atom. The atomic spectrum of hydrogen consists of several series of lines.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 5</div> <div> Planck’s Quantum Theory of Light:</div> <div> In 1099, Max Planck proposed a revolutionary theory about the nature of light. The theory</div> <div> successfully explained several observations concerning the interaction of light and matter, which</div> <div> could not be explained by the wave theory of light. The theory is called Plank’s quantum theory</div> <div> of light. It states:</div> <div> 1) Light is emitted or absorbed by matter discontinuously in the form of discrete packets or</div> <div> bundles of energy called quanta (singular quantum). These quanta of light are also named</div> <div> as photons.</div> <div> 2) Each photon of light has a particular amount of energy which is proportional to the</div> <div> frequency of light.</div> <div> Here E is the energy, v is the frequency, ? is the wavelength, c is the velocity of light, and h is the</div> <div> constant of proportionality called the Planck’s constant. Its value is 6.63 x 10-34 joule second.</div> <div> The Planck’s quantum theory, thus, says that light is discontinuous and consists of discrete</div> <div> particles called photons. The light (energy) is absorbed or emitted by a substance in the form of</div> <div> these photons.</div> <div> Bohr’s Atomic Model</div> <div> Keeping in view the Rutherford’s atomic model, the Planck’s quantum theory of light, and the</div> <div> results of the studies of atomic spectra, Neils Bohr (1885 – 1962), a Danish physicist, in 1913 put</div> <div> forward his model of atomic structure. The postulates of this model are as follows:</div> <div> 1. Electrons revolve around the nucleus in certain fixed orbits, or stationary states. It means</div> <div> that the energy of the electrons is quantified, that is, an electron can have only certain</div> <div> discrete quantities of energy and none in between. The orbits are also called energy levels.</div> <div> 2. Only those orbits are possible for which the angular momentum (mvr) of an electron is the</div> <div> integral multiple of h/2?. That is,</div> <div> mvr = n x h/2? Where n = 1,2,3,4,…</div> <div> Where m is the mass, v the velocity, r the radius of the orbit of an electron, h is the</div> <div> Planck’s constant. n specifies the orbits. When n = 1, this is orbit number 1, when n = 2,</div> <div> this is orbit number 2, and so on.</div> <div> 3. As long as an electron moves in its own orbit, no emission or absorption of radiation</div> <div> occurs.</div> <div> 4. Emission or absorption of radiation takes place when an electron jumps from one orbit to</div> <div> another. An electron jumps from an orbit of lower energy (an inner orbit) to an orbit of</div> <div> higher energy (an outer orbit) as a result of the absorption of radiation. When an electron</div> <div> jumps from an orbit of higher energy to an orbit of lower energy, radiation is emitted. The</div> <div> energy of the light absorbed or emitted is equal to the difference of the energies of the</div> <div> orbits between which the electronic transition occurs.</div> <div> :. ΔE = E2 – E1 = hv</div> <div> Radii or Orbits: Different orbits have different radii. On the basis of Bohr’s atomic model, radii of</div> <div> the electronic orbits can be determined.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 6</div> <div> Ground State and Excited State</div> <div> The ground state of an electron is the lowest energy state that an electron normally occupies.</div> <div> When energy is absorbed, the electron jumps to a higher energy state. This higher energy state is</div> <div> known as the excited state (higher energy state). The ground state of an electron is the stable state</div> <div> while the excited states are unstable states. As a result of absorption of energy when an electron</div> <div> jumps to an excited state, it immediately jumps back to the ground state with the emission of</div> <div> energy.</div> <div> Wave Nature of Matter:</div> <div> In 1924, a French physicist Louis de Broglie, using Planck’s equation (E = hv) and Einstein’s</div> <div> equation (E = mc2), put forward a hypothesis that electron (and hence matter) possesses wave</div> <div> properties. That is, electron is a wave. Later, Davison and Germer in 1927 experimentally</div> <div> discovered the wave nature of electron. Subsequently, other particles, such as protons, neutrons</div> <div> and helium nuclei (a-particles) were also shown to exhibit wave properties.</div> <div> The wave behaviour of matter can only be observed for very small particles. For big bodies the</div> <div> associated wavelength is much too small to be observed.</div> <div> In short, under certain conditions, an electron resembles a particle while under some other</div> <div> conditions, it behave as if it is a wave. The electron and hence the matter possesses dual nature.</div> <div> However, it has been found that electrons can be treated as wave more effectively than as small</div> <div> particles.</div> <div> Heisenberg Uncertainty Principle</div> <div> In 1927, Werner Heisenberg (1901–1976), a German physicist, put forward a principle called the</div> <div> Heisenberg uncertainty principle (or Heisenberg indeterminacy principle). It states:</div> <div> “It is impossible to determine with exact accuracy both the momentum and the position of an</div> <div> electron (or any other body) simultaneously.”</div> <div> It means the more accurately we know the momentum of an electron, the less accurately can we</div> <div> know its position and vice versa.</div> <div> Thus, Bohr’s model of an electron moving with a known velocity (or momentum, as it is mv), in a</div> <div> well-defined orbit (that is, with known position) cannot be correct. In other words, we cannot say</div> <div> where an electron lies around the nucleus.</div> <div> Quantum Mechanics</div> <div> The concepts given by Planck, Bohr, de Broglie, Heisenberg, and others ultimately led to the</div> <div> development of a new theory of the structure of atom, known as quantum mechanics or wave</div> <div> mechanics. Quantum mechanics replaced the classical mechanics (the Newtonian mechanics).</div> <div> If the electron is a wave, then it should be possible to describe it with the help of some</div> <div> mathematical equations like the other waves. In 1926, Erwin Schrodinger (1887–1961) succeeded</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 7</div> <div> to develop an equation. This equation is the foundation stone of the quantum mechanics. It offers</div> <div> a mathematical approach to understanding the wave nature of matter.</div> <div> The quantum mechanical treatment of an electron is highly mathematical as the Schrodinger</div> <div> equation is a very complex equation. Hydrogen is the only atom for which it has been solved.</div> <div> Each solution of this equation is described by a set of constant numbers called quantum numbers.</div> <div> Quantum Numbers</div> <div> The quantum mechanics explains the structure of atom with the help of four quantum numbers.</div> <div> These numbers are used to describe the distribution of electrons in space around the nucleus. The</div> <div> quantum numbers and their significance are given below:</div> <div> 1. Principal Quantum Number, n</div> <div> The principal quantum number (n) describes the main energy level, or shell, that an electron</div> <div> occupies in an atom. n may be any positive integer:</div> <div> n = 1, 2, 3, ……</div> <div> When n = 1, it is the first main energy level, or the first shell. When n =2, it is the second shell.</div> <div> And, so on. A main energy level or shell is the main region around the nucleus in which an</div> <div> electron may be present. Shells 1, 2, 3, 4,….. are also sometimes called K, L, M, N, …….shells.</div> <div> 2. Azimuthal (or subsidiary) Quantum Number, l</div> <div> Azimuthal quantum number (l) describes subshells or sublevels in shells or main energy levels. It</div> <div> denotes the shape of the region in space an electron occupies. Its values are given by:</div> <div> l = 0 to (n – 1)</div> <div> Thus, when n = 1, then</div> <div> l = 0</div> <div> When n = 2, then l = 0, 1</div> <div> When n = 3, then l = 0, 1, 2</div> <div> And, so on.</div> <div> Thus, the values of l designate subshells or sublevels that electrons can occupy in a main shell.</div> <div> Each value of l is given a letter notation. Thus:</div> <div> l = 0, 1, 2, 3, 4, 5, …….</div> <div> s, p, d, f, g, h,..........</div> <div> Thus, in the first shell (n = 1, l = 0), there is one subshell and it is s. In the second shell there are</div> <div> two subshells, s and p. in the third shell (n = 3), there are three subshells, s, p and d. In the fourth</div> <div> shell (n = 4), there are four subshells (s, p, d, and f). And, so on. Notice that the number of</div> <div> subshells in any given shell is simply equal to the value of n for that shell.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 8</div> <div> (The first four letters (s, p, d, f) find their historical origin in the study of the atomic spectra of the</div> <div> alkali metals. In these spectra four series of lines are observed and they are termed as the sharp,</div> <div> principal, diffuse and fundamental series. For l = 4, 5, …… we just continue with the alphabet.)</div> <div> Notice that the electrons in their ground state in atoms of all elements occupy only s, p, d and f.</div> <div> 3. Magnetic Quantum Number, m</div> <div> It designates the spatial orientation of the region an electron occupies. This region is called an</div> <div> orbital. An orbital is the region around the nucleus in which the probability of finding an electron</div> <div> is maximum. Each subshell is composed of one or more orbitals. The values of m are given by:</div> <div> m = -l, …….0…….+l</div> <div> Thus, when l = 0, m = 0</div> <div> When l = 1, m = -1, 0, +1</div> <div> When l = 2, m = -2, -1, 0, +1, +2</div> <div> When l = 3, m = -3, -2, -1, 0, +1, +2, +3</div> <div> Values of m serve to determine the orientation of an orbital in space relative to the other orbitals.</div> <div> In subshell s, there is only one orbital s. In subshell p, there are three orbitals, px, py ,pz. In</div> <div> subshell d, there are five orbitals, and in subshell f, there are seven orbitals.</div> <div> 4. Spin Quantum Number, s</div> <div> It refers to the spin of an electron. Electron has a spin, that is, it rotates about its own axis. An</div> <div> electron can spin only in either of two directions, clockwise or anticlockwise. Therefore, s has</div> <div> only two values +1/2 and –1/2.</div> <div> Atomic Orbitals:</div> <div> From the above discussion we can conclude that in an atom there are various main energy levels</div> <div> or shells. Each shell has one or more sublevels or subshells. Each subshell consists of one or more</div> <div> orbitals. An orbital is the actual region where an electron is found. It can be defined as follows:</div> <div> “An orbital is the region around the nucleus in which the probability of finding an electron is</div> <div> maximum.”</div> <div> Note that an orbital can have a maximum of two electrons.</div> <div> SHAPES OF ATOMIC ORBITALS</div> <div> Each s orbital in an atom is spherically symmetrical with respect to the nucleus. That is, an s</div> <div> orbital in shape is like a symmetrical sphere such as a football. Each p orbital is like a dumbbell.</div> <div> The three p orbitals (px, py and pz) are present perpendicular to one another. Subscripts x, y and z</div> <div> indicate the axis along which each of the three two-lobed p orbitals is directed. The d and f</div> <div> orbitals have complex shapes.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 9</div> <div> ELECTRONIC CONFIGURATION</div> <div> The arrangement of electrons in shells, subshells and rbitals in an atom is called the electronic</div> <div> configuration. The electronic configuration of an atom is governed by the following rules and</div> <div> principles:</div> <div> 1. Pauli Exclusion Principle,</div> <div> Since electrons are described by four quantum numbers, any two electrons in an atom must differ</div> <div> in at least one quantum number. Pauli Exclusion Principle states:</div> <div> “ In an atom no two electrons can have the same values of the four quantum numbers.”</div> <div> It can also be stated as:</div> <div> “An orbital can accommodate no more than two electrons and these electrons must have opposite</div> <div> spins.”</div> <div> Take the example of helium, He, which has two electrons, both the electrons are in 1s orbital. We</div> <div> express this as 1s2. These two electrons have the same three quantum numbers, for both</div> <div> n = 1</div> <div> l = 0</div> <div> m = 0</div> <div> But, they differ in the fourth quantum number. If for one electron</div> <div> s = + 1/2</div> <div> then for the other s = -1/2</div> <div> The two electrons having opposite spins are said to be “paired”, while if there is only one</div> <div> electron in an orbital, it is called “unpaired”.</div> <div> (Pauli exclusion principle was put forward by a Swiss physicist Walgang Pauli (1900-1958) in</div> <div> 1925.)</div> <div> 2. Aufbau Principle</div> <div> ( Aufbau is from German aufbauen meaning “to build up”). This principle states:</div> <div> “In an atom an electron tends to occupy the lowest energy orbital available.”</div> <div> As we go away from the nucleus, the energy of the shells increases. In writing electronic</div> <div> configuration, we start filling electrons from 1s. When it is completely filled, we go to 2s the to</div> <div> 2p and son on. Consider the following:</div> <div> H (Z =1) 1s1</div> <div> He (Z = 2) 1s2</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 10</div> <div> Li (Z= 3) 1s2, 2s1</div> <div> Be (Z=4) 1s2, 2s2</div> <div> B (Z=5) 1s2, 2s2, 2p1</div> <div> C (Z=6) 1s2, 2s2, 2p2</div> <div> N ( Z =7) 1s2, 2s2, 2p3</div> <div> O ( Z =8) 1s2, 2s2, 2p4</div> <div> F (Z=9) 1s2, 2s2, 2p5</div> <div> Ne (Z=10) 1s2, 2s2, 2p6</div> <div> Na (Z=11) 1s2, 2s2, 2p6 , 3s1</div> <div> 3. (n+l) Rule</div> <div> This rule states: “The orbital with the lowest energy is the one for which the sum of n and l is</div> <div> lowest. When two orbitals have the same ( n + l ) value, the orbital with the lower n value has the</div> <div> lower energy.</div> <div> Take the example of potassium (K) that has 19 electrons. The first eighteen electrons have the</div> <div> following arrangement:</div> <div> 1s2, 2s2, 2p6, 3s2, 3p6</div> <div> Where will go the 19th electron, to orbital 3d or 4s? Let us apply (n+l) rule, for 3d</div> <div> n + l</div> <div> 3 + 2 =5</div> <div> and for 4s</div> <div> n + l</div> <div> 4 + 0 =4</div> <div> Thus, 4s has lower energy than 3d; the 19th electron will go to 4s.</div> <div> K (Z = 19) 1s2, 2s2, 2p6, 3s2, 3p6 ,4s1</div> <div> 4. Hund’s Rule</div> <div> All the orbitals in a subshell have equal energy. How are then electrons filled in a subshell?</div> <div> Hund’s rule states:</div> <div> “As long as the orbitals of same energy available, electrons occupy the separate orbitals and</div> <div> have parallel spins.”</div> <div> Atomic Radius</div> <div> The size of an atom is measured in terms of radius. It is not possible to determine the precise size</div> <div> of an atom since the outermost shell of an atom has no precise outer boundary. As we go away</div> <div> from the nucleus, the electron density decreases to a very small value but it does not become</div> <div> exactly zero even at a large distance from the nucleus. Moreover, we cannot isolate a single atom</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 11</div> <div> and measure its diameter the way we can measure the diameter of a football. We can estimate</div> <div> only the radius of a combined atom. Moreover, the size of the electronic cloud also depends upon</div> <div> its environment. In spite of all these difficulties, we talk about the atomic size because it is very</div> <div> useful in understanding many properties of elements. The atomic radius is defined as follows.</div> <div> “The half of the bond length of a homonuclear diatomic molecule is called the atomic radius of</div> <div> the bonded atoms.”</div> <div> That is, if two atoms of the same element are covalently bonded, the half of the bond length will</div> <div> be the radius of each atom.</div> <div> For Example, H-H bond length is 0.78 A. The one-half of this distance is 0.37 A. This is the</div> <div> atomic radius of hydrogen. Similarly, the CI-CI bond length is 1.98 A, the atomic radius of CI is</div> <div> 0.99 A. Other atomic radii can be found in a similar way. This atomic radius is the covalent</div> <div> radius since it is the radius of an atom involved in a covalent bond, and not the radius of a free</div> <div> atom.</div> <div> If we know the length of a bond between two dissimilar atoms (A-B), and the radius of one of</div> <div> them, we can calculate the atomic radius of the other:</div> <div> rA + rB = A-B bond length</div> <div> Change of atomic radius in the periodic table</div> <div> 1. Atomic radius decreases as we go from left to right in the periodic table. It is because as</div> <div> we go from left to right in a period, the nuclear charge increases. This increasing nuclear</div> <div> charge pulls the valence electrons with increasing force and as a result the atomic radius</div> <div> decreases from left to right in the periodic table.</div> <div> 2. Atomic radius increases as we go from top to bottom in a group. It is because the number</div> <div> of shells increases in going from top to bottom. The increasing number of shells increases</div> <div> the atomic radius.</div> <div> Ionic Radius</div> <div> When one or more electrons are added to an atom, a negative ion, anion, is formed. While, when</div> <div> one or more electrons are removed from an atom, a positive ion, cation, is formed .In general,</div> <div> cations are always smaller than the neutral atoms from which they are formed, whereas, anions</div> <div> are larger than the neutral atoms from which they are formed.</div> <div> The decrease in the size of a cation is due to two factors:</div> <div> 1) The removal of electron(s) from the outermost shell. Often all the electrons from the</div> <div> outermost shell are removed which means the removal of a whole shell.</div> <div> 2) The decrease in the total electron-electron repulsion. For example, radius of Na = 1.86A</div> <div> and that of Na+ = 0.95 A.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 12</div> <div> The increase in the size of an anion is due to the added repulsion between the new electrons and</div> <div> the electrons already present. For example, radius of CI=0.99 A and of CI- = 1.81 A.</div> <div> In a series of ions of different elements having the same number of electrons (isoelectronic ions)</div> <div> the size decreases with the increase in the nuclear charge.</div> <div> Ionization Energy</div> <div> Since electrons are attracted by the nucleus, energy must be supplied to remove an electron from</div> <div> an atom. This energy is called ionization energy or ionization potential, and defined as follows:</div> <div> “The minimum amount of energy required to remove the most loosely bound electron from an</div> <div> isolated gaseous atom in its ground state is called the ionization energy,”</div> <div> For example, the ionization energy of Na is 495.8 kJ/mol.</div> <div> An atom has as many ionization energies as it has electrons. The energy required to remove the</div> <div> first electron is called the first ionization energy, the energy required to remove the second</div> <div> electron is known as the second ionization energy, and so forth. For example, for calcium the first</div> <div> ionization energy (IE) is 590 kJ/ mol.</div> <div> The second ionization energy is 1145 kJ/ mol.</div> <div> The ionization energy depends on the following factors</div> <div> 1. Atomic radius</div> <div> The greater the atomic radius, the farther away is the outermost electron from the nucleus; the</div> <div> easier it is to remove it. Thus, lower will be the ionization energy.</div> <div> 2. Nuclear Charge</div> <div> The greater the nuclear charge, the more attracted will be the electrons by the nucleus, and the</div> <div> higher will be the ionization energy.</div> <div> 3. Shielding effect</div> <div> The intervening shells between the nucleus and the outermost electrons reduce the</div> <div> attraction between the nucleus and the outermost electrons. This effect is known as the shielding</div> <div> or screening effect. The greater the shielding effect, the easier it will be to remove the electron,</div> <div> and the lower will be the ionization energy. The shielding effect increases from top to bottom in</div> <div> the periodic table with the increase in the number of shells.</div> <div> 4. Type of orbital</div> <div> The order of ionization energy with respect to orbitals is</div> <div> s > p> d > f</div> <div> It is most difficult to remove an electron from an s orbital because it is least diffused. The f is the</div> <div> most diffused orbital, so it is the easiest to remove an electron from f orbital.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 13</div> <div> Change of IP in the period table.</div> <div> In the periodic table the ionization energy increases from left to right due to decrease in the</div> <div> atomic size and increase in the nuclear charge.</div> <div> From top to bottom in a group, the ionization energy decreases due to increase in the atomic size</div> <div> and increase in the shielding effect.</div> <div> Electron Affinity</div> <div> Electron affinity can be defined as follows. When an electron enters an atom, energy is either</div> <div> released or absorbed. This energy is called electron affinity. Thus:</div> <div> “The amount of energy evolved or absorbed when an electron is added to a gaseous atom in its</div> <div> ground state is called electron affinity.”</div> <div> Electron affinity shows the ability of an atom to accept one or more electrons. When a chlorine</div> <div> atom accepts an electron 348kJ/mol energy is released.</div> <div> Cl(g) + e ----------------Cl(g) EA= -348 kJ/mol</div> <div> The halogens that have the most negative electron affinities readily form negative ions. By</div> <div> accepting one electron a halogen atom assumes the stable electronic configuration of the noble</div> <div> gas immediately following it.</div> <div> In general, electron affinity becomes more and more negative in going from left to right in the</div> <div> periodic table as the nuclear charge increases and the atomic radius decreases.</div> <div> Some elements gain more than one electrons to form anions. For example, oxygen gains two</div> <div> electrons to form a bivalent anion.</div> <div> O(g) + e ---------------- O (g) EA = -142 kJ/mol</div> <div> O(g) + e ---------------- O(g) EA = +780 kJ/mol</div> <div> In the second step energy is absorbed. Electron affinities of anions are always positive. This is</div> <div> because of the repulsion between the negative ion and the incoming electron.</div> <div> Problem: (a) Which element should easily form the cation: K (IP= 419kJ/mol or F (IP=</div> <div> 1680kJ/mol)?</div> <div> (b) Which element should easily form anion: Na (EA= -53kJ/mol) or Br (EA= -324kJ/mol)?</div> <div> Answer: (a) K, (b) Br.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 14</div> <div> Lecture 2</div> <div> Matter and Energy</div> <div> Fundamental Concepts of Chemistry: Matter and Chemistry</div> <div> Chemistry is a branch of science that deals with matter and changes that occur in it. In chemistry</div> <div> we also study theories, laws and principles related to changes in matter. More precisely, in</div> <div> chemistry we study chemical and physical changes and properties of elements and compounds.</div> <div> In this lecture our focus is on the basic concepts of chemistry. There are many concepts that are</div> <div> used in chemistry. The fundamental concepts include atom, molecule, ion, element, compound,</div> <div> mixture, symbols, formulas, equations, atomic number, mass number, atomic mass, molecular</div> <div> mass, mole, etc.</div> <div> Matter:</div> <div> Matter is any substance that has physical existence. That is, anything that has mass and volume.</div> <div> All that we see around us, or we can feel by touch are material substances. Mass is the quantity of</div> <div> matter that a substance contains in it. Volume is the space that a body occupies.</div> <div> Physical states:</div> <div> Matter exists in three physical states. These are: solid, liquid and gas. In solids, the particles are</div> <div> very close to one another. In gases, the particles are very away from one another. In liquids, the</div> <div> situation is intermediate. A substance can be changed from one physical state to another. Ice,</div> <div> liquid water and steam are three physical states of water.</div> <div> Atom:</div> <div> The atom is the smallest particle of an element that represents the elements in all chemical</div> <div> reactions. There are three types of fundamental particles in an atom.</div> <div> 1. Electrons</div> <div> 2. Protons</div> <div> 3. Neutrons</div> <div> The electron is a very small particle. It has negative electric charge. The proton is a positively</div> <div> charged particle which is much heavier than electron. Neutron is an electrically neutral particle. It</div> <div> is the heaviest among the three particles. Protons and neutrons are found in the center of atom and</div> <div> collectively constitute a body called nucleus. Electrons revolved around the nucleus.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 15</div> <div> Atomic Mass:</div> <div> An atom is a very small particle. But it also has mass. The quantity of this mass is very small.</div> <div> Atomic mass is the mass of one atom of an element and is defined as follows. The relative</div> <div> average mass of one atom of an element which tells us how heavy or light is that atom with</div> <div> respect to one atom of carbon-12. This unit is called amu or atomic mass unit.</div> <div> One atomic mass unit is one-twelfth of the mass of carbon-12.</div> <div> 1 amu = 1/12 x(mass of one atom of C-12)</div> <div> In this unit carbon-12 is chosen as standard. Carbon-12 is the most abundant isotope of carbon.</div> <div> For example: the atomic mass of hydrogen is 1.008amu.</div> <div> Molecular mass:</div> <div> By molecular mass we mean the mass of a molecule. It is the sum of atomic masses of all the</div> <div> atoms of all elements present in a molecule. We can calculate the molecular masses of water,</div> <div> sulfuric acid, and glucose as follows.</div> <div> 1. Water, H2O; 1x2 + 16 = 18 amu</div> <div> 2. H2SO4: 1 x 2 + 32 +16 x 4 = 98 amu</div> <div> 3. C6H12O6: 12 x 6 + 1 x 12 + 16 x 6 = 180 amu</div> <div> Mole:</div> <div> The amount of a substance that contains Avogadro’s number of particles is called a mole. It is the</div> <div> atomic mass, molecular mass or formula mass expressed in grams. For example, 18 grams of</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 16</div> <div> water is one mole of water. If you drink 9 grams of water, it means you have taken a half more of</div> <div> water. The Avodagro’s Number is the number of particles that is found in one mole of a</div> <div> substance. Its value is: 6.02 x 1023. It is a huge number. If a person has eaten 180 g of glucose, he</div> <div> has eaten 6.02 x 1023 of glucose molecules.</div> <div> Formula mass:</div> <div> The Formula mass of a substance is the sum of the atomic masses of all the atoms present in the</div> <div> formula of that substance. For example, the formula mass of NaCl is 23+ 35.5 = 58.5.</div> <div> Molecular mass:</div> <div> The molecular mass of a substance is the sum of the atomic masses of all the atoms present in a</div> <div> molecule of that substance. For example, molecular mass of water, H2O, is 1x2+16= 18.</div> <div> Ion:</div> <div> An electrically charged atom or molecule is called an ion. It has two types: cations or positive</div> <div> ions, and anions or negative ions. H+ is a cation whereas Cl- is an anion. Ions may be simple or</div> <div> compound. Simple ion consists of a single atom, but a compound ion is comprised of two or more</div> <div> atoms. For example, Na+ is a simple ion but SO4</div> <div> 2- is a compound ion.</div> <div> Molecule:</div> <div> An atom or group of bonded atoms which can exist independently. One molecule of helium (He)</div> <div> consists of one atom of helium. It is a monoatomic molecule. The molecules of other noble gases</div> <div> are also monoatomic. One molecule of water (H2O) consists of three atoms. So it is a polyatomic</div> <div> molecule.</div> <div> Atomic number:</div> <div> The total number of protons in the nucleus of an atom is called atomic number. The atomic</div> <div> number of hydrogen is 1. The atomic number of oxygen is 8, sodium is 11 and iron is 26. All the</div> <div> atoms of an element have same atomic number.</div> <div> Mass number:</div> <div> The total number of protons and neutrons present in the nucleus of an atom is called the mass</div> <div> number of that atom. The mass number of carbon-12 is 12. It has 6 protons and 6 neutrons.</div> <div> Atoms of an element may differ in the mass number.</div> <div> Isotopes:</div> <div> The atoms of the same element having different mass number are called isotopes. Isotopes of an</div> <div> element have equal number of protons but different number of neutrons. Hydrogen element has</div> <div> three isotopes. They are called:</div> <div> 1. Ordinary hydrogen or protium,</div> <div> 2. Deuterium or heavy hydrogen, and</div> <div> 3. Tritium or radioactive hydrogen.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 17</div> <div> Ordinary hydrogen has no neutron. Deuterium has one neutron. Tritium has two neutrons.</div> <div> Element:</div> <div> An element is the simplest form of matter which cannot be converted into simper forms by any</div> <div> physical or chemical means. If in a collection of atoms, all atoms have the same atomic number,</div> <div> they all belong to the same element.</div> <div> For example, Na, Fe, O etc. are elements. If you look at a periodic table, you will see elements in</div> <div> it. Elements are of many types. They have different physical states. Today, more than 110</div> <div> elements are known. Of these 90 elements occur naturally. Other elements have been formed in</div> <div> nuclear reactors through nuclear reactions. Majority of the elements are metals, but many</div> <div> elements are non-metals. There are some elements which have properties in-between metals and</div> <div> non-metals. They are given the name metalloids.</div> <div> Metals are elements which have high electropositivity. They easily lose electrons and have</div> <div> shining surfaces. For example: Cu, Ag, Au are metals. Non-metals are generally less</div> <div> electropositive and more electronegative.</div> <div> Compound:</div> <div> A compound is a substance that is formed by the chemical combination of two or more elements</div> <div> in a definite ratio. For example, water (H2O) and common salt (NaCl) are compounds. At present,</div> <div> many millions of compounds are known.</div> <div> Both the elements and compounds are pure substances. Elements are represented by symbols and</div> <div> compounds are shown by chemical formulas.</div> <div> Mixture:</div> <div> When two or more substances are mixed in any ratio in a manner that no chemical combination</div> <div> occurs is called a mixture. A mixture is a mixed substance formed by the physical combinations</div> <div> of two or more substances (elements or compounds). For example, air is a mixture of nitrogen,</div> <div> oxygen and some other gases. The mixtures are of two types.</div> <div> 1. Homogeneous mixtures, and</div> <div> 2. Heterogeneous mixtures</div> <div> A homogeneous mixture has uniform composition. It is also known as a solution. For example, a</div> <div> mixture of salt (NaCl) and water is a homogeneous mixture. In a homogenous mixture, one thing</div> <div> is completely “dissolved” in another substance. On the other hand, a heterogeneous mixture does</div> <div> not have uniform composition. For example, a mixture of sand and water is a heterogeneous</div> <div> mixture.</div> <div> Chemical reaction:</div> <div> A chemical reaction is a chemical change that involves one or more elements or compounds. It is</div> <div> process in which one or more chemical bonds are broken or formed. In a chemical reaction, one</div> <div> or more substances undergo chemical change and form one or more new substances. The</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 18</div> <div> substances before the reaction are called reactants and the substances that are formed as a result of</div> <div> the reaction are called products.</div> <div> A chemical reaction is shown by a “chemical equation”. A chemical equation consists of</div> <div> reactants, products and an arrow sign.</div> <div> Reactants Products</div> <div> A chemical equation must be “balanced”. A balanced chemical equation is that in which the</div> <div> number of atoms of all elements involved in the reaction are equal on both sides of equation. That</div> <div> is, the number of atoms of an element in the reactants must be equal to the number of atoms of</div> <div> that element in the product. Look at the following equation. Is it balanced?</div> <div> H2 + Cl2 HCl (not balanced)</div> <div> No. it is not balanced. Why? Because the number of hydrogen atoms in the reactants side is 2</div> <div> while in the product side there is only 1 H atom. Similarly, number of Cl atoms is not equal on</div> <div> both the sides. How can you balance it? Write 2 before HCl in the product side. The equation</div> <div> now will be balanced.</div> <div> H2 + Cl2 2HCl (balanced)</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 19</div> <div> Lecture 3</div> <div> Nature of Chemical Bonding</div> <div> How atoms join with each other?</div> <div> The existence of solids and liquids is a clear proof that atoms combine with one another. But, the</div> <div> question is why atoms combine with one another and how?</div> <div> Isolated atoms are unstable and to gain stability they combine with one another. This combination</div> <div> is expressed by the term chemical bond (the word bond means linkage, attachment, connection,</div> <div> etc.). The chemical bond can be defined as follow:</div> <div> “The chemical bond is a force which holds atoms together in a molecule or a compound”</div> <div> It was in 1916 that W. Kossel (Germany) and G.N. Lewis (USA) independently proposed that</div> <div> atoms join together through the involvement of electrons. The noble gases (He, Ne, Ar, etc.) are</div> <div> stable, and have eight electrons in their outermost shells (except helium which has two electrons).</div> <div> It was thought that since the noble gases are stable, all elements tend to achieve the noble gas</div> <div> electronic configuration. This is known as the octet rule. (Octet: group of eight.) Now, how can</div> <div> atoms achieve eight electrons in their outermost shells? Kossel suggested that this is done by</div> <div> actual transfer of one or more electrons from one atom to another. Lewis proposed that atoms</div> <div> acquire noble gas electronic configuration by mutual sharing of electrons between two atoms.</div> <div> These views gave the two concepts of bonds: ionic bond, and covalent bond.</div> <div> Ionic Bond or Electrovalent Bond</div> <div> The concept of ionic bond was suggested by W. Kossel in 1916. Ionic bond is formed due to</div> <div> transfer of electrons. When electrons are transferred from one atom to another, the atom that loses</div> <div> electrons becomes a positive ion and the atom that gains electrons becomes a negative ion. These</div> <div> oppositely charged ions then attract each other and are attached with each other through the</div> <div> opposite electric charges. Thus, we can define the ionic bond as follows:</div> <div> “The ionic bond is the bond which is formed between oppositely charged ions due to electrostatic</div> <div> force of attraction.”</div> <div> Take the example of NaCl. Sodium atom has one electron in its valence (outermost) shell.</div> <div> Chlorine atom has seven electrons in its outermost shell.</div> <div> Sodium atom loses its valence electron and becomes a positive ion, while chlorine atom accepts</div> <div> one electron to form a negative ion. These ions attract each other due to their opposite changes to</div> <div> form ionic bond:</div> <div> Na Na+ + e-</div> <div> Cl + e- Cl-</div> <div> Na+ + Cl- Na+Cl- NaCl</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 20</div> <div> Energetics of Ionic Bond</div> <div> Why an ionic bond is formed? The natural principle is that every system tends towards more</div> <div> stability. Atoms form bonds to acquire stability. For most cases we can understand stability as a</div> <div> state of less energy. Thus, in general, we can say that every system tends to acquire a state of</div> <div> lower energy. Let us study the energy changes in the formation of NaCl from gaseous sodium and</div> <div> chlorine atoms.</div> <div> To remove the outermost electron from sodium atom the energy (the ionization energy) required</div> <div> is 494 kJ/mol. When a chlorine atom gains an electron, the energy (electron affinity) released is</div> <div> 348 kJ/mol.</div> <div> This shows that the ionization energy of sodium is so high that the energy released when a</div> <div> chlorine atom accepts the electron is not sufficient to provide the energy needed to remove the</div> <div> electron from sodium. An additional 495-348 = 147 kJ energy is required to make possible the</div> <div> transfer of one mole of electrons from isolated sodium atoms to isolated chlorine atoms.</div> <div> These sodium ions and the chloride ions combine together to form solid crystalline sodium</div> <div> chloride. During this step a high amount of energy (788 kJ/mol) is released.</div> <div> Na+(g) + Cl-(g) NaCl(s) ΔH = - 788 kJ/mol</div> <div> This energy is known as the lattice energy.</div> <div> “The lattice energy is the amount of energy evolved when one mole of a crystalline ionic</div> <div> compound is formed from its gaseous ions.”</div> <div> The lattice energy here is high because one mole of sodium ions and one mole of chloride ions are</div> <div> arranged in the formation of crystal in such a way that one sodium ion is surrounded by six</div> <div> chloride ions and vice versa. Thus, it is the large lattice energy, which makes the formation of</div> <div> sodium chloride energetically favorable. The lattice energy is, in fact, the driving force behind the</div> <div> formation of sodium chloride from its elements.</div> <div> In the same way we can understand the formation of any ionic compound. Ionic bond is formed</div> <div> between the elements having large difference of electronegativity, (usually higher than 1.7). In</div> <div> general the elements of group VIIA (halogens) form ionic bond with metals.</div> <div> Characteristics of Ionic Compounds</div> <div> 1. Ionic Compounds are Non-molecular</div> <div> An ion attracts the oppositely charged ions from all directions, so the ionic bond is nondirectional.</div> <div> When a large number of cations and anions come closer to one another, they arrange</div> <div> themselves in such a way that one ion is surrounded by a number of oppositely charged ions.</div> <div> Each ion is surrounded by the greatest possible number of oppositely charged ions depending on</div> <div> the relative ionic radii of cations and anions. Thus, ionic bonds do not form freely existing</div> <div> particles called molecules. Therefore, ionic compounds do no consist of molecules.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 21</div> <div> 2. Ionic Compounds are Solids</div> <div> Under ordinary conditions, all ionic compounds are solids. It is because in ionic compounds, ions</div> <div> are held together by strong electrostatic forces of attraction.</div> <div> 3. Ionic compounds have high melting and boiling points</div> <div> In ionic compounds, the oppositely charged ions are held in the crystal by strong electrostatic</div> <div> forces. Therefore, high energy is needed to overcome the strong electrostatic forces. Thus, the</div> <div> ionic compounds have high melting and boiling points.</div> <div> 4. Ionic compounds are soluble in polar solvents</div> <div> Polar solvents like water have high dielectric constants (the ability of a solvent to decrease the</div> <div> attractive force between the ions dipped in it). Therefore, such solvents break the forces of</div> <div> attraction between the ions and get them dispersed in them. In nonpolar solvents ionic compounds</div> <div> are generally insoluble.</div> <div> 5. Ionic compounds are electrolytes</div> <div> Ionic compounds conduct electricity in their solutions and in molten sates. This is because in</div> <div> solution or in molten state the ions become free from one another and can move under the effect</div> <div> of electric field.</div> <div> Covalent Bond</div> <div> The concept of covalent bond was given by G.N. Lewis in 1916. Covalent bond can be defined as</div> <div> follows:</div> <div> “Covalent bond is the bond which is formed due to</div> <div> sharing of electrons between tow atoms.”</div> <div> Take the example of a hydrogen molecule. Hydrogen atom has one electron in its valence shell:</div> <div> H (Z = 1) 1s1</div> <div> When two hydrogen atoms come close to each other, they share their electrons to form a covalent</div> <div> bond:</div> <div> H. +H. H:H</div> <div> This shared pair of electrons holds the two atoms together; each electron of the pair is attracted</div> <div> by both the nuclei. The shared pair of electrons is also known as bonding pair of electrons.</div> <div> Covalent bond is a unidirectional bond and holds the bonded atoms to form a freely exiting</div> <div> particle called a molecule. Therefore, the covalent bond is generally shown by a small line drawn</div> <div> between the symbols of the bonded atoms. The line represents the bonding pair of electrons. Thus</div> <div> H:H is represented as H–H</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 22</div> <div> Type of Covalent Bond</div> <div> A covalent bond between two atoms may be formed by mutual sharing of one, two or three pairs</div> <div> of electrons. Therefore, the covalent bonds are of three types.</div> <div> Single Covalent Bond: The bond that is formed due to sharing of one pair of electrons between</div> <div> two atoms is called the single covalent bond. The H-H bond is a single covalent bond. Similarly,</div> <div> H-C1, Cl-C1, F-F Br-Br are single covalent bonds.</div> <div> Double Covalent Bond: The bond that is formed due to sharing of two pairs of electrons between</div> <div> two atoms is called a double covalent bond. Each atom contributes two unpaired electrons that are</div> <div> paired with the electrons of the other atom. Take the example of O2. An oxygen atom has the</div> <div> following electronic configuration.</div> <div> O 1s2, 2s2, 2px</div> <div> 2 , apy</div> <div> 1, 2pz</div> <div> 1</div> <div> It has two unpaired electrons, one in orbital 2py and other in 2pz. Two oxygen atoms share two</div> <div> pairs of electrons to form a double covalent bond:</div> <div> O: + :O O::O</div> <div> or</div> <div> O=O</div> <div> Similarly, ethene has a carbon-carbon double bond:</div> <div> H2C=CH2</div> <div> Triple Covalent Bond</div> <div> When a bond is formed due to sharing of three pairs of electrons between two atoms, it is called a</div> <div> triple covalent bond. Take the example of N2. A nitrogen atom (z = 7) has the following</div> <div> electronic configuration:</div> <div> N 1s2, 2s2, 2px</div> <div> 1 , apy</div> <div> 1, 2pz</div> <div> 1</div> <div> It has three unpaired elections. Two nitrogen atoms form a triple bond by sharing three pairs of</div> <div> electrons. Similarly, ethyne has a carbon-carbon triple bond.</div> <div> Coordinate Covalent Bond</div> <div> Coordinate covalent bond is a special type of covalent bond, which differs from the simple</div> <div> covalent bond in the way of formation:</div> <div> “The covalent bond between two atoms in which both of the bonding electrons are donated by</div> <div> one of the atoms, is known as the coordinate covalent bond.”</div> <div> The H+ ion has no electron, yet it forms a bond. For example, it forms a bond with the nitrogen of</div> <div> ammonia. It is because the nitrogen of ammonia has a non-bonding, or lone, pair of electrons</div> <div> which it donates to form a bond with hydrogen ion:</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 23</div> <div> H+ + :NH3 NH4</div> <div> +</div> <div> A coordinate covalent bond is also represented by an arrow sign showing the direction of the</div> <div> donation of electron pair.</div> <div> The atom that donates the pair of electrons is called the donor, and the atom that accepts the pair</div> <div> of electrons is called the acceptor.</div> <div> Coordinate covalent bond is also known as dative bond or co-ionic bond.</div> <div> -----------------------------------------------------------------------------</div> <div> Characteristics of Covalent Compounds</div> <div> 1. Covalent compounds are molecular: The atoms bonded by covalent bonds form</div> <div> independently existing units known as molecules.</div> <div> 2. Covalent compounds may be solids, liquids or gases: The covalent compounds consist</div> <div> of molecules. In gases these molecules are free from one another and exert practically no force of</div> <div> attraction under ordinary conditions. In liquids and solids intermolecular forces of attraction, van</div> <div> der Waals forces, or hydrogen bonds hold the molecules together.</div> <div> 3. Non-electrolytes: In general, the covalent compounds are non-electrolytes. Exceptions</div> <div> include acids and bases.</div> <div> 4. Insoluble in water: In general, the covalent compounds are insoluble in polar solvents</div> <div> like water, and soluble in nonpolar solvents like organic liquids.</div> <div> Electronegativity</div> <div> A shared pair of electrons is attracted by both of the bonded atoms. For example in H:C1 both H</div> <div> and C1 attract the bonding pair. This power of an atom to attract the bonding pair of electrons is</div> <div> called electronegativity, which can be defined as follows:</div> <div> “The ability of an atom to attract the shared pair of electrons towards it is known as</div> <div> electronegativity.”</div> <div> Fluorine is the most electronegative element. Its electronegativity is 4. Then comes oxygen whose</div> <div> electronegativity is 3.5.</div> <div> Polar and Nonpolar Bonds</div> <div> If two bonded atoms have equal electronegativity, the bonding electrons pair will be shared</div> <div> equally between the two atoms. In other words, the electron density of the bonding electron pair is</div> <div> distributed equally and symmetrically between the two atoms. Such a bond is known as a</div> <div> nonpolar bond. Thus;</div> <div> “The bond between two atoms having equal electronegativity is a nonpolar bond.”</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 24</div> <div> For example, C1-C1, O=O, N=N bonds are nonpolar bonds.</div> <div> On the other hand, if the atoms forming a bond have different electronegativities, the bonding pair</div> <div> is shared unequally. The more electronegative atom acquires a greater share of the bonding</div> <div> electrons as compared to the less electronegative atom. Due to this the more electronegative atom</div> <div> acquires a partial negative charge and the less electronegative atom acquires a partial positive</div> <div> charge. These charges are of equal magnitude, and the overall charge is zero. Such a bond is</div> <div> known as a polar bond. It has negative and positive poles. Thus:</div> <div> “The polar bond is the bond between two atoms having different electronegativities.”</div> <div> Take the example of H–C1. Chlorine is more electronegative than hydrogen. Thus, the bonding</div> <div> pair of electrons is shared unequally, and chlorine has greater share of electrons than hydrogen.</div> <div> As a result, chlorine is partially negatively charged and hydrogen is partially positively charge.</div> <div> ---------------------------------------------------------------------------------</div> <div> Problem:</div> <div> Which of the following molecules are polar and which are nonpolar?</div> <div> BeCl2, H2S, CS2, BF3, NH3, PCl3</div> <div> Answers:</div> <div> BeCl2 is a linear molecule, and is nonpolar.</div> <div> H2S is an angular molecule (like H2O) with polar bonds, so it is polar.</div> <div> CS2 is a linear molecule (like CO2), so it is nonpolar.</div> <div> BF3 is a trigonal planar molecule. Its bond moments cancel one another so the molecule is</div> <div> nonpolar.</div> <div> PC13 is also a trigonal pyramidal molecule (like NH3) and polar.</div> <div> Bond length</div> <div> “The distance between the nuclei of two bonded</div> <div> atoms is called the bond length”.</div> <div> The bond length of a single bond is longer than a double bond, which in turn is longer than a</div> <div> triple bond.</div> <div> Bond Energy</div> <div> When a bond is formed between two atoms a certain amount of energy is released during this</div> <div> process. When that bond is broken the same amount of energy is required to do so. This energy is</div> <div> known as bond energy</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 25</div> <div> “The bond energy is the energy evolved when a bond is formed between two atoms.”</div> <div> Or</div> <div> “The bond energy is the energy required to break a bond between two atoms.”</div> <div> The bond energy is also known as bond dissociation energy.</div> <div> The Modern Concept of covalent Bonding</div> <div> There are two important modern concepts of chemical bonding both are based on the quantum</div> <div> mechanics: valence bond theory (VBT), and molecular orbital theory (MOT). According to VBT</div> <div> “a covalent bond between two atoms is formed due to the overlap of two half-filled orbitals, one</div> <div> from each atom.”</div> <div> Molecular orbital theory (MOT) treats a molecule like an atom. An atom consists of atomic</div> <div> orbitals surrounding a nucleus. Likewise, a molecule consists of molecular orbitals surrounding</div> <div> two or more nuclei present in the molecule. Atomic orbitals are monocentric, they surround only</div> <div> one nucleus while the molecular orbitals are polycentric, and surround more than one nuclei.</div> <div> Summary: A chemical bond is the force which holds atoms together. The bond may be ionic or</div> <div> covalent. Ionic bond is formed by the transfer of 1 or more electrons and covalent bond is formed</div> <div> by the sharing of electrons. In ionic bond there is complete transfer of electrons which result in</div> <div> formation of positive and negative ions which are called cations and anions these cations and</div> <div> anions are held together by the strong electrostatic force of attraction. So the ionic bond is very</div> <div> strong bond. The electrons which are present in the outer most shell of an element are called</div> <div> valence electrons. It the valence electrons which take part in bonding, for example there are one</div> <div> valence electrons in Na and seven in Cl.</div> <div> Differentiation between ionic and covalent bond:</div> <div> · Ionic compounds have ionic compounds while covalent compounds have covalent bonds.</div> <div> · Since in ionic compounds, ions are held together through very strong electrostatic forces,</div> <div> ionic compounds are solids at room temperature. Whereas covalent compounds consist of</div> <div> molecules which are held together through weak inter molecular force, there for they may</div> <div> be solids, liquids or gases.</div> <div> · Ionic compounds have higher melting and boiling points than covalent compounds.</div> <div> · In general, ionic compounds are soluble in water and insoluble in organic solvents. While</div> <div> covalent compounds are generally insoluble in water and soluble in organic solvents.</div> <div> · Ionic compounds don’t conduct electricity as such but do so in molten state. Covalent</div> <div> compounds don’t conduct electricity in general.</div> <div> Explain bond formation in H2 and NaCl?</div> <div> H2 In H2 two hydrogen atoms are held together through a covalent bond. Each hydrogen atom</div> <div> contributed 1 electron. So there are two shared electrons this is known as a shared pair of</div> <div> electrons. It is this shared pair of electrons which holds two H atoms together.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 26</div> <div> H. + H. H__H</div> <div> NaCl In NaCl there is ionic bond Na+ is positive ion while Cl- is negative ion, these ions are</div> <div> held together through very strong electrostatic force of attraction.</div> <div> Na+ + Cl- NaCl</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 27</div> <div> Lecture 4</div> <div> Chemical Reactions</div> <div> Chemical Reactions:</div> <div> A chemical reaction is a chemical change in which one or more substances change into one or</div> <div> more other substances. In other words, a chemical reaction is a process in which one or more</div> <div> bonds are broken and one or more bonds are formed. For example, rusting of iron is a chemical</div> <div> reaction. Similarly burning of natural gas is a chemical reaction.</div> <div> Chemical and physical change:</div> <div> During a chemical change a new substance is formed while during a physical change no new</div> <div> substance is formed. For example, burning of wood is a chemical change while change of water</div> <div> into ice is a physical change. During a chemical change a chemical reaction occurs whereas</div> <div> during a physical change no chemical reaction occurs. Chemical change is generally a permanent</div> <div> change, while physical change is only a temporary change and can be reversed easily.</div> <div> Chemical property of a substance is a property in which that substance undergoes a chemical</div> <div> change. For example, it is a chemical property of an acid to react with a base and form salt and</div> <div> water.</div> <div> Exothermic and endothermic reactions:</div> <div> When a chemical reaction takes place, change in energy also occurs. The reaction during which</div> <div> heat is evolved is called an exothermic reaction. On the other hand, the reaction during which</div> <div> heat is absorbed is called an endothermic reaction.</div> <div> Heat of reaction:</div> <div> When a chemical reaction occurs, certain amount of heat is either evolved or absorbed. This</div> <div> amount of heat which is evolved or absorbed during a chemical reaction is known as heat of</div> <div> reaction. It is also known as enthalpy of reaction and is denoted by ΔH. Its units are joules per</div> <div> mole or Kcal per mole.</div> <div> Types of Reactions:</div> <div> There are a number of types of reactions.</div> <div> 1. Addition reactions: the reactions in which two or more substance react to form a single</div> <div> substance. Ethene reacts with hydrogen to form ethane.</div> <div> C2H4 + H2 C2H6</div> <div> 2. Substitution reaction: the reactions in which one substance is replaced by another</div> <div> substance from a molecule:</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 28</div> <div> CH3-Cl + NaOH CH3-OH + NaCl</div> <div> This reaction is also known as double displacement reaction.</div> <div> 3. Elimination reaction: the reactions in which one or more atoms are eliminated from a</div> <div> molecule. For example, removal of H2O from ethanol to form ethene.</div> <div> CH3-CH2-OH CH2=CH2 + H2O</div> <div> 4. Oxidation-Reduction Reactions</div> <div> Many reactions involve transfer of electrons from one atom to another atom. These reactions are</div> <div> known as oxidation-reduction reactions. The loss of electrons from an atom is known as</div> <div> oxidation, while the gain of electrons by an atom is called reduction. Oxidation and reduction</div> <div> always take place simultaneously and the extent of oxidation is equal to the extent of reduction.</div> <div> That is, the loss of electrons is equal to the gain of electrons. The substance that undergoes</div> <div> oxidation is known as the reducing agent, and the substance that undergoes reduction is known</div> <div> as the oxidizing agent.</div> <div> 5. Combustion reactions: combustion means burning, and it is the reaction of a substance</div> <div> with oxygen. In a combustion reaction heat is produced. For example, the reaction of methane</div> <div> with oxygen is a combustion reaction.</div> <div> CH4 + O2 CO2 + H2O</div> <div> Generally in a combustion reaction, carbon dioxide and water are produced.</div> <div> 6. Neutralization reactions:</div> <div> When an acidic solution is added to a basic solution in the right proportion, the properties of each</div> <div> are lost, and a salt and water are formed. This process is called neutralization.</div> <div> “The reaction between an acid and a base that forms a salt and water is called neutralization”.</div> <div> For Example,</div> <div> HCI (aq) + NaOH (aq) NaCI (aq) + H2O</div> <div> HNO3 (aq) + NaOH (aq) NaNO3 (aq) + H2O</div> <div> Chemical reactions may occur in solution or gaseous form. Generally, the chemical reactions are</div> <div> carried out in solution. Water and organic solvents are used for this purpose.</div> <div> Rate of Reaction:</div> <div> When a reaction occurs, one or more bonds are broken and then one or more bonds are formed.</div> <div> This process of bond breakage and formation takes some time. How much a reactant changes into</div> <div> a product in a given time is important. By the rate of reaction we mean the amount of reactant that</div> <div> is converted to product per unit time. It can also be defined on the bases of the appearance of a</div> <div> product. So, the rate of a reaction is the amount of product formed per unit time.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 29</div> <div> The rate of reaction is also called speed or velocity of reaction. It is measured in moles/seconds.</div> <div> Thus, the rate of a reaction is the amount of product which is formed in one second.</div> <div> Rate of Reaction = concentration of product formed/time</div> <div> On the basis of the rate of reactions, we can classify the reactions into the following categories.</div> <div> · Very slow reactions: some reactions are very slow. For example, conversion of graphite</div> <div> into diamond in the earth crust is a very slow process.</div> <div> · Slow reactions: some reactions occur with slow speed, e.g. rusting of iron is a slow</div> <div> process.</div> <div> · Moderate reactions: most organic reactions occur with moderate speed.</div> <div> · Fast reactions: many reactions occur with a fast speed. Strong acid strong base reactions</div> <div> are fast reactions.</div> <div> · Very fast reactions: some reactions are very fast such as explosions.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 30</div> <div> Lecture 5</div> <div> Organic Chemistry</div> <div> Organic Chemistry:</div> <div> Organic chemistry is a branch of chemistry that involves the study of carbon compounds. It deals</div> <div> with the structures, composition, and synthesis of organic compounds. Organic compounds are</div> <div> obtained from natural sources like plants, animals and microbes. They are also synthesized in</div> <div> laboratory. There are millions of organic compounds. They have wide variety of application such</div> <div> as in medicine, fertilizer, food, clothing, communication, transportation etc.</div> <div> Catenation:</div> <div> Carbon has a very unique ability that its large number of atoms can bond with each other to make</div> <div> chains of different lengths. This property of carbon to make chains is called catenation.</div> <div> C-C C-C-C C-C-C-C C-C-C-C-C C-C-C-C-C-C-C</div> <div> Hydrocarbons:</div> <div> Hydrocarbons are the compounds of carbon and hydrogen. There are different classes of</div> <div> hydrocarbons such as aliphatic, aromatic, cyclic, acyclic, saturated and unsaturated, etc. Aromatic</div> <div> compounds are benzene and those which resemble benzene. They burn with black soot. They</div> <div> possess aromaticity and are comparatively more stable than expected. They are unsaturated</div> <div> compounds but are resistant to addition reactions. Their characteristic reactions are electrophilic</div> <div> substitution reactions. The structure of benzene is given below.</div> <div> Classification of organic compounds:</div> <div> There are millions of organic compounds and this makes it physically impossible to study each</div> <div> individual compound. To facilitate their study, organic compounds are classified into various</div> <div> groups and sub-groups.</div> <div> They may be broadly classified into the following classes:</div> <div> · Open chain or acyclic compounds</div> <div> · Closed chain or cyclic compounds</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 31</div> <div> Open chain or acyclic compounds:</div> <div> These compounds contain an open chain system of carbon atoms. The chains may be either</div> <div> straight chains (non-branched) or branched. The open chain compounds are also called aliphatic</div> <div> compounds. The name aliphatic is derived from the Greek word aleiphar meaning fats, as the</div> <div> earlier known compounds of this class were either obtained from animal or vegetable fats, or had</div> <div> fat like character.</div> <div> CH3-CH2-CH2-CH3 , CH3-CH2-CH2-CH2OH, H2C = CH2</div> <div> Closed chain or cyclic compounds</div> <div> These compounds contain one or more closed chains (rings) of atoms and are known as cyclic or</div> <div> ring compounds.</div> <div> The aliphatic hydrocarbons are the organic compounds which contain no benzene ring. They may</div> <div> be open chain or closed chain. The closed chain compounds are also called cyclic compounds.</div> <div> The major classes of aliphatic hydrocarbons are alkanes, alkenes, alkynes and cycloalkanes, etc.</div> <div> Functional groups:</div> <div> An atom or group of atoms which imparts some specific properties to a compound is called a</div> <div> functional group. For example, -OH is functional group. It is known as hydroxyl group. The</div> <div> compounds which have this functional group show specific properties. They are phenols and</div> <div> alcohols.</div> <div> Structural formula:</div> <div> It is very common in organic chemistry that two or more compounds have same molecular</div> <div> formula. For example, both ethanol and dimethyl ether have same molecular formula, C2H6O.</div> <div> Why they are different? They are different because they have different arrangements of atoms in</div> <div> their molecules. The arrangement of atoms in molecule is represented by a formula which is</div> <div> known is the structural formula. The structural formula of a compound is the formula which</div> <div> shows the arrangement of atoms in its molecule. The structural formula of a compound shows</div> <div> which atom is attached with which atom.</div> <div> Isomers and isomerism:</div> <div> Two or more different compounds which have same molecular formula are called isomers. This</div> <div> phenomenon is known as isomerism. Ethanol and dimethyl ether are isomers. Similarly, n-butane</div> <div> and isobutene are isomers. They are two different compounds, and have same molecular formula</div> <div> (C4H10) but different structures. Isomerism is very common in organic compounds. There are</div> <div> many types of isomerisms.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 32</div> <div> Sources of organic compounds:</div> <div> There are many sources of organic compounds. Plants, animals, petroleum, natural gas, coal are</div> <div> common sources of naturally occurring organic compounds. A huge number of compounds have</div> <div> been synthesized in laboratories by chemists. A great number of new compounds are being added</div> <div> continuously.</div> <div> All organic compounds must contain carbon. Almost organic compounds also possess hydrogen.</div> <div> Many compounds also contain oxygen and/or nitrogen. Other common elements found in organic</div> <div> compounds are halogens, sulfur, and phosphorus.</div> <div> Oxygen containing organic compounds:</div> <div> Oxygen-containing organic compounds are very numerous. Many functional groups contain</div> <div> oxygen. Some examples are given below.</div> <div> Alcohols:</div> <div> Alcohols are common organic compounds. They contain –OH functional group. This is called</div> <div> hydroxyl group. Examples: Methanol, CH3-OH ; ethanol, CH3-CH2-OH.</div> <div> Both methanol and ethanol are liquids and are used as solvents.</div> <div> Phenols:</div> <div> Phenols also contain –OH functional group, but in a phenol the –OH group is attached to an</div> <div> aromatic carbon.</div> <div> Aldehydes:</div> <div> Aldehydes are those organic compounds which have aldehydic carbonyl functional group,-CHO.</div> <div> Formaldehyde, acetaldehyde, benzaldehyde are common aldehydes.</div> <div> Ketones:</div> <div> Ketones are those organic compounds which contain ketonic carbonyl functional group. This</div> <div> group is –CO–. There is a double bond between carbon and oxygen, C=O. Acetone is a common</div> <div> ketone (H3C-CO-CH3). It is a liquid and used as a solvent.</div> <div> Carboxylic acid:</div> <div> Carboxylic acids are compounds which contain carboxylic group, –COOH. Carboxylic acids are</div> <div> important organic compounds. Examples: formic acid, acetic acid, benzoic acid. Carboxylic acids</div> <div> are common in living organisms, proteins are composed of amino acids. Citric acid is found in</div> <div> citrus fruits. Milk contains lactic acid. Gapes have tartaric acid.</div> <div> Carboxylic acids are weaker than mineral acids. Oils and fats are esters of glycerol with fatty</div> <div> acids.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 33</div> <div> Esters:</div> <div> Esters are also common organic compounds. They are formed by combination of carboxylic acids</div> <div> and alcohols. Ethyl acetate is a common ester. Its formula is: CH3-COOCH2-CH3. Fats and oils</div> <div> are esters of fatty acids with glycerol. Glycerol is a triol, i.e., it has three –OH groups.</div> <div> Ethers:</div> <div> Ethers are also a common class of organic compounds. They have C-O-C bonding. Dimethyl</div> <div> ether, and diethyl ether are common ethers. Diethyl ether is commonly called ether. It is</div> <div> anesthetic. It is used as a solvent in laboratory and industry.</div> <div> Nitrogen containing organic compounds:</div> <div> Nitrogen containing organic compounds are also very common. Amino acids contain nitrogen.</div> <div> Similarly DNA and RNA also contain nitrogen. Another important class of natural products</div> <div> called alkaloids also contains nitrogen. Common nitrogenous compounds have amino groups.</div> <div> Amines:</div> <div> Amines are nitrogenous organic compounds which contain amino group. Amines may be primary,</div> <div> secondary and tertiary. Examples: CH3-NH2 (Methyl amine), CH3-NH-CH3 (Dimethyl amine) and</div> <div> (CH3)3N (Trimethyl amine).</div> <div> Carbohydrates, proteins and lipids (Fats and oils) are naturally occurring common organic</div> <div> compounds. They are also part of our food.</div> <div> Organic Reactions:</div> <div> A chemical reaction is a chemical change. It is a process in which one or more bonds and broken</div> <div> and one or more bonds are formed. Coutless number of reactions occur involving organic</div> <div> compunds. But all these reaction belong to only a few types of reactions. Most common types of</div> <div> organic reactions are given below.</div> <div> · Substitution reactions</div> <div> · Elimination reactions</div> <div> · Additions reactions</div> <div> In substitution reactions, one group is replaced by another group.</div> <div> OH- + CH3-Cl CH3-OH + Cl-</div> <div> In the reaction Cl- is replaced by OH- group.</div> <div> In elimination reactions, groups are eliminated from a molecule.</div> <div> CH3-CH2-Cl H2C=CH2 + HCl</div> <div> In this reaction HCl is eliminated from chloroethane and ethene is formed.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 34</div> <div> Addition reaction is reverse of elimination reaction. In an addition reaction, a molecule adds to</div> <div> another molecule. So, two molecules of reactants combine to form one molecule of product.</div> <div> H2C=CH2 + HBr H3C-CH2-Br</div> <div> Applications of Organic Compounds:</div> <div> Organic compounds are very important chemical substances. They have a wide variety of</div> <div> applications in modern human life. Organic compounds are essential part of our food. The fibers</div> <div> and dyes are organic compounds. Similarly, fragrances and flavours are organic substances.</div> <div> Plastic and other polymers are also organic compounds.</div> <div> Numerous organic compounds are used as medicines for different diseases. The number of such</div> <div> compounds is rapidly increasing.</div> <div> Many fertilizers are also organic compounds. Similarly, soaps and detergents and other household</div> <div> materials are organic compounds.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 35</div> <div> Lecture 12</div> <div> Metals</div> <div> The modern classification of elements is based on the modern periodic law.</div> <div> The modern periodic law states that the physical and chemical properties of elements are the</div> <div> periodic function of their atomic numbers.</div> <div> When elements are arranged on the basis of the modern periodic law, we get modern periodic</div> <div> table.</div> <div> In the modern periodic table the elements are arranged in the order of increasing atomic number.</div> <div> When the elements are arranged in the order of increasing atomic number, the elements with</div> <div> similar properties appear at definite intervals. When elements with similar properties are placed in</div> <div> the same column, we get groups or families of elements having similar properties. Thus, in the</div> <div> modern periodic table, we have vertical columns and horizontal rows. The horizontal rows are</div> <div> also called periods.</div> <div> Periods in the periodic table:</div> <div> Rows of elements are called periods. There are 7 period in the periodic table. The number of</div> <div> elements in a period increases as you move down the periodic table.</div> <div> Groups in the periodic table:</div> <div> Columns of elements in the periodic table are called groups. Elements within a group show</div> <div> similarity in properties. In a group, the elements have the same outer electron arrangement. The</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 36</div> <div> outermost electrons are called valence electrons. Why the elements belonging to the same group</div> <div> have similar properties? It is because they have the same number of valence electrons.</div> <div> Representative elements: In the periodic table there are two sets of groups. The group A</div> <div> elements are called the representative elements. The group B elements are the non-representative</div> <div> elements. Representative elements are named after the first element of the group. So we have,</div> <div> lithium family, beryllium family, boron family, carbon family, nitrogen family, oxygen family,</div> <div> fluorine family or halogens, and helium family or noble gases. Group IA and IIA are also called</div> <div> alkali metals and alkaline earth metals respectively.</div> <div> Transition metals:</div> <div> The transition elements are located in groups IB to VIIIB. Copper, zinc, iron, silver, mercury and</div> <div> gold are examples of transition metals. These elements are very hard, with high melting points</div> <div> and boiling points. The transition metals are good electrical conductors and are very malleable.</div> <div> Periodic properties:</div> <div> There are a number of properties which change in a definite manner as we move from one end to</div> <div> the other end of periodic table. Such properties are called periodic properties. Periodic properties</div> <div> include: atomic and ionic radii, melting points and boiling points, metallic character, ionization</div> <div> potential, electron affinity, electronegativity, etc. take the example of electronegativity. The</div> <div> electronegativity of an element is the ability of the element with which it attracts the shared pair</div> <div> of electrons towards it. The electronegativity increases as we go from left to right in the periodic</div> <div> table. It decreases from top to bottom in a group.</div> <div> Physical states of elements:</div> <div> Most elements in the periodic table are solids. At room temperature all metals are solids except</div> <div> mercury. Only a few elements are liquids such as bromine and mercury. Some elements are gases.</div> <div> These are nonmetals. Nitrogen, oxygen, fluorine, chlorine and noble gases are found in gases state</div> <div> at room temperature.</div> <div> Ores: An ore is a material that contains a metal in such quantities that it can be mined and</div> <div> worked commercially to extract that metal.</div> <div> Metals are not found in free state. They are found in the form of minerals and ores.</div> <div> The metal is usually contained in chemical combination with some other element in addition to</div> <div> various impurities. For example, iron is found in the form of ores such as hematite (Fe2O3),</div> <div> magnetite (Fe3O4), limonite (FeO(OH).n(H2O)).</div> <div> Properties of metals and nonmetals</div> <div> (ref. (http://chemistry.about.com/od/periodictableelements/a/Metals-And-</div> <div> Nonmetals.htm)</div> <div> Metal Physical Properties</div> <div> · lustrous (shiny)</div> <div> · good conductors of heat and electricity</div> <div> · high melting point</div> <div> · high density (heavy for their size)</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 37</div> <div> · malleable (can be hammered)</div> <div> · ductile (can be drawn into wires)</div> <div> · usually solid at room temperature (an exception is mercury)</div> <div> · opaque as a thin sheet (can't see through metals)</div> <div> · metals are sonorous or make a bell-like sound when struck</div> <div> Metal Chemical Properties</div> <div> · have 1-3 electrons in the outer shell of each metal atom</div> <div> · corrode easily (e.g., damaged by oxidation such as tarnish or rust)</div> <div> · lose electrons easily</div> <div> · form oxides that are basic</div> <div> · have lower electronegativities</div> <div> · are good reducing agents</div> <div> Nonmetals</div> <div> Nonmetal Physical Properties</div> <div> · not lustrous (dull appearance)</div> <div> · poor conductors of heat and electricity</div> <div> · non-ductile solids</div> <div> · brittle solids</div> <div> · may be solids, liquids or gases at room temperature</div> <div> · transparent as a thin sheet</div> <div> · nonmetals are not sonorous</div> <div> Nonmetal Chemical Properties</div> <div> · usually have 4-8 electrons in their outer shell</div> <div> · readily gain or share valence electrons</div> <div> · form oxides that are acidic</div> <div> · have higher electronegativities</div> <div> · are good oxidizing agents</div> <div> Metalloids or semimetals:</div> <div> There are some elements which have properties between metals and nonmetals. They are called</div> <div> metalloids or semimetals. The metalloids or semimetals are located along the line between the</div> <div> metals and nonmetals in the periodic table. The metalloids are boron, silicon, germanium, arsenic,</div> <div> antimony, and tellurium.</div> <div> · Electronegativities between those of metals and nonmetals</div> <div> · Ionization energies between those of metals and nonmetals</div> <div> · Possess some characteristics of metals/some of nonmetals</div> <div> · Reactivity depends on properties of other elements in reaction</div> <div> · Often make good semiconductors</div> <div> Metals are generally good conductor of electricity. It is because metal has free electrons in it,</div> <div> according to modern bonding concepts metals have a special type of bond which is called metallic</div> <div> bond. The metal atom held together in a way there valence electrons are delocalized and free to</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 38</div> <div> move in the entire crystal. This is known as sea model. These free electrons help in conduction of</div> <div> electricity.</div> <div> For example cooper wires are commonly used to conduct electricity because they have free</div> <div> electrons which conduct electricity.</div> <div> On the other hand most nonmetals are poor conductor of electricity it is because they do not</div> <div> possess free electrons like metals. Some nonmetals do conduct electricity. For example, Graphite</div> <div> has layer structure and posses free electrons which conduct electricity.</div> <div> Chemical properties:</div> <div> Differences between Metals and Nonmetals</div> <div> Metals Nonmetals</div> <div> Lose their valence electrons easily. Gain or share valence electrons easily.</div> <div> Form oxides that are basic. Form oxides that are acidic.</div> <div> Are good reducing agents. Are good oxidizing agents.</div> <div> Have lower electro negativities Have higher electro negativities</div> <div> Usually have 1-3 electrons in their outer</div> <div> shell.</div> <div> Usually have 4-8 electrons in their outer</div> <div> shell.</div> <div> Physical properties</div> <div> How can you compare physical properties of Metals and Nonmetals? Some differences are given</div> <div> below:</div> <div> Metals Nonmetals</div> <div> Good electrical conductors and heat</div> <div> conductors.</div> <div> Poor conductors of heat and electricity.</div> <div> Malleable - can be beaten into thin sheets. Brittle - if a solid.</div> <div> Ductile - can be stretched into wire. Non-ductile.</div> <div> Possess metallic luster. Do not possess metallic luster.</div> <div> Opaque as thin sheet. Transparent as a thin sheet.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 39</div> <div> Lecture 13</div> <div> Metallurgy of Copper & Aluminum</div> <div> Almost all metals are found in the combined state as compounds on the crust of the earth. So we</div> <div> have to extract them. This process of metal extraction is called Metallurgy.</div> <div> Metallurgy is the science and art of extracting a metal from its ores and carrying out different</div> <div> processes to bring metals into their useful forms. Thus, the various processes involved in the</div> <div> extraction of metals from their ores and refining them are broadly known as metallurgy.</div> <div> Minerals: The natural materials in which the metals or their compounds occur in the earth are</div> <div> called minerals.</div> <div> Ores: Those minerals from which the metals can be extracted profitably are called ores.</div> <div> Gangue or Matrix: The rocky impurities, including silica and earthly particles, present in an ore</div> <div> are called gangue or matrix.</div> <div> Metallurgy of Aluminium:</div> <div> Aluminum is the most abundant metal. It is the third most abundant element in the earth crust</div> <div> (about 7%). Aluminium is a reactive metal and so does not occur free in nature. It occurs in clays</div> <div> in the form of aluminium silicates.</div> <div> Ores of aluminium:</div> <div> Aluminium has a number of ores. Some important ores are as follows:</div> <div> 1. Oxide Ores:</div> <div> Bauxite, Al2O3.xH2O</div> <div> Corundum, Al2O3</div> <div> 2. Fluoride Ore:</div> <div> Cryolite, 3NaF.AlF3 or Na3AlF6</div> <div> 3. Silicate Ores:</div> <div> Kaolin, Al2O3.2SiO2.2H2O</div> <div> Potash mica, K2O.3Al2O3.6SiO2.2H2O</div> <div> Extraction of Aluminium:</div> <div> (Ref: http://www.s-cool.co.uk/gcse/chemistry/extraction-of-metals/revise-it/the-electrolysisof-</div> <div> bauxite)</div> <div> Aluminium is commonly extracted from its ore bauxite. The bauxite (red-brown solid) -</div> <div> aluminium oxide mixed with impurities - is extracted from the earth. The extracted aluminium</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 40</div> <div> oxide is then treated with alkali, to remove the impurities. This results in a white solid called</div> <div> aluminium oxide or alumina. The alumina is then transported to huge tanks. The tanks are lined</div> <div> with graphite, this acts as the cathode. Also blocks of graphite hang in the middle of the tank, and</div> <div> acts as anodes. The alumina is then dissolved in molten cryolite - this lowers the melting point.</div> <div> Electricity is passed and electrolysis begins. Electrolysis is the decomposition of a compound</div> <div> using electricity.</div> <div> When dissolved, the aluminium ions and oxide ions in the alumina can move.</div> <div> Reaction at Cathode:</div> <div> 4Al+ + 12e- 4Al(s)</div> <div> Reaction at Anode:</div> <div> 3C + 6O2- 3CO2 (g) + 12e-</div> <div> Purification of Bauxite:</div> <div> Bauxite when obtained from the earth crust contains many ipurities such as ferric oxide, silica,</div> <div> and titanium oxide. These impurities are removed. Baeyer process is commonly used for the</div> <div> purification of bauxite. In this process, the ore is treated with concentrated NaOH solution called</div> <div> caustic soda. The following reaction takes place:</div> <div> Al2O3 + 2NaOH 2Na[Al(OH)4](aq)</div> <div> Then CO2 is passed.</div> <div> Reaction with CO2</div> <div> 2Al(OH)4</div> <div> - + CO2 2Al(OH)3(s) + CO3</div> <div> 2- + H2O</div> <div> The product is filtered and heated at 1000 °C. The process is called calcination. As a result of this</div> <div> process, pure alumina is obtained.</div> <div> Al(OH)3 + Heat Al2O3 (s) + 3H2O (g)</div> <div> Now we have pure alumina.</div> <div> Electrolysis of Alumina to obtain Aluminium:</div> <div> The process used for the electrolysis of alumina is called Hall’s process. Alumina is aluminium</div> <div> oxide, Al2O3. This is subjected to electrolysis. The cell used for electrolysis is a steel container</div> <div> whose inner surface is lined with graphite. The powdered alumina is mixed with fused cryolite</div> <div> and fluorspar. The melting point of alumina is very high, but it dissolves in molten cryolite at</div> <div> about 1000 °C. when electricity is passed, aluminium is deposited at the graphite cathode. From</div> <div> there it is drawn out in molten form. This aluminium is 99% pure. Further purification is carried</div> <div> out through a process called Hoope’s method.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 41</div> <div> Refining of Aluminium:</div> <div> In the Hoope’s process, aluminium is purified by electrolytic method. For this purpose a box-like</div> <div> cell is used. The cell is operated in a manner that in it three layers are maintained. The lower layer</div> <div> consists of impure aluminum alloyed with copper. The middle layer consists of a solution of</div> <div> cryolite and barium fluoride. The upper layer consists of pure aluminium. These layers remain</div> <div> separated due to difference is specific gravity. The lower layer acts as anode where aluminium is</div> <div> oxidized and Al3+ ions move toward the upper layer which is negatively charged. This layer acts</div> <div> as cathode and when Al3+ ions reach in this layer they gain electrons and are deposited as</div> <div> aluminium metal there.</div> <div> Reaction at Anode:</div> <div> Al Al3+ + 3e-</div> <div> Reaction at cathode:</div> <div> Al3+ + 3e- Al (pure aluminium)</div> <div> The aluminium obtained in this process is about 99.99% pure.</div> <div> Metallurgy of Copper:</div> <div> Copper is a common metal. It is used for a wide variety of purposes. It is a good conductor of</div> <div> electricity so its wires are used in electrical appliance and cables.</div> <div> Over 360 ores of Copper are known. Some important ores are as follows:</div> <div> Sulfide ores:</div> <div> Copper pyrite or chalcopyrite, CuS.FeS or CuFeS2</div> <div> Chalcocite, Cu2S</div> <div> Carbonate ores:</div> <div> Malachite, CuCO3.Cu(OH)2</div> <div> Azurite, 2CuCO3.Cu(OH)2</div> <div> Oxide ores: Cuprite, Cu2O</div> <div> Extraction of Copper:</div> <div> Extraction of copper is good example of pyrometallurgy. In the metallurgy of copper the</div> <div> following steps are commonly used:</div> <div> • Mining of ore from its mines</div> <div> • Crushing and grinding of the ore</div> <div> • Concentration of the ore and its purification</div> <div> • Chemical process to obtain metal from the ore</div> <div> • Purification of the metal</div> <div> • Conversion of the metal into alloys</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 42</div> <div> Important process employed in the extraction of copper:</div> <div> – Concentration through Froth-flotation</div> <div> – Roasting</div> <div> – Smelting</div> <div> – Reduction (Bessemerization)</div> <div> – Refining</div> <div> Concentration by Froth flotation method:</div> <div> By concentration here we mean to free the ore from impurities. Copper is commonly extracted</div> <div> from copper pyrite. Copper pyrite is concentrated by a process called froth-floatation. Ore is</div> <div> crushed and ground to make a powder. This is placed in a large tank containing water and oil</div> <div> (such as pine oil). The mixture is stirred by blowing air through it. The ore is wetted by oil and</div> <div> rises to the surface from where it is skimmed off. The impurities (called gangue) are wetted with</div> <div> water settle down at the bottom.</div> <div> Roasting:</div> <div> By roasting here we mean heating in air strongly but below the melting point of the ore. As a</div> <div> result, impurities are converted into silicate slag. The slag floats at the surface from where it is</div> <div> removed. The ore undergoes the following reaction:</div> <div> CuS.FeS + 4O2 Cu2S (s)+ 2FeO (s) + 3SO2 (g)</div> <div> Smelting:</div> <div> By smelting here we mean heating the product obtained from roasting strongly in a Reverberatory</div> <div> furnace. The product from the roasting is mixed with silica (SiO2) and heated at 1100 °C.</div> <div> Following reaction takes place in the furnace:</div> <div> FeO + SiO2 FeSiO3</div> <div> Ferrous silicate is formed which is a slag. It is less dense that molten copper matte, Cu2S, and so</div> <div> floats the surface from where it is removed.</div> <div> Reduction (Bessemerization):</div> <div> Molten copper(I) sulfide, Cu2S, is put in Bessemer converter and heated strongly. As a result of</div> <div> this elemental copper is produced. The copper obtained is called the blistered copper.</div> <div> 2Cu2S + 3O2 2Cu2O + 2SO2</div> <div> 2Cu2O(l) + Cu2S(l) 6Cu (l) + SO2(g)</div> <div> This copper is called blistered copper because it has blisters on it which are produced due to</div> <div> evolution of sulfur dioxide. When copper solidifies, sulfur dioxide escapes from it and produces</div> <div> blisters. Blistered copper is about 99% pure.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 43</div> <div> Refining of Copper:</div> <div> Copper is refined by electrolytic process. In the electrolytic cell CuSO4 solution is used as an</div> <div> electrolyte. Impure copper is used to act as anode, while pure copper is used to act as a cathode.</div> <div> At anode, copper atoms lose electrons and become copper ions. These copper ions move through</div> <div> the solution and reach at the cathode. At cathode these copper ions gain electrons and deposit</div> <div> there a atoms. So in this way refined copper is obtained which is almost 100% pure.</div> <div> Reaction at Anode:</div> <div> Cu Cu2+ + 2e-</div> <div> Reaction at Cathode:</div> <div> Cu2+ + 2e- Cu (s) pure copper</div> <div> --------------------------</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 44</div> <div> Lecture 14</div> <div> Metallurgy of Iron</div> <div> Ores:</div> <div> An ore is the naturally occurring compound of a metal from which that metal can be extracted. In</div> <div> the earth crust, metals are generally found in the form of compounds which are found mixed with</div> <div> other substances these naturally occurring minerals of metals are called ores. For example; halite</div> <div> is an ore of Na and hematite and magnetite are the ores of Fe.</div> <div> Alloys:</div> <div> A material that has metallic properties and is composed of two or more chemical elements of</div> <div> which at least one is a metal (i.e. steel is an alloy of carbon in iron; stainless steel is an alloy of</div> <div> carbon, chromium and sometimes nickel in iron.)</div> <div> Metallurgy:</div> <div> It is the science and art of obtaining a metal from its ores. There are number of types of</div> <div> metallurgy.</div> <div> In electrometallurgy, electricity is used to obtain a metal. Electrolysis is carried out of the molten</div> <div> or solution form of the ore and metal is obtained at the cathode. For example; Na is obtained by</div> <div> the electrolysis of NaCl by the Downs process.</div> <div> In pyrometallurgy strong heat is used to extract a metal from its ores. For example, Fe is extracted</div> <div> from its ores by heating at high temperature.</div> <div> Iron is extracted through pyrometallurgy in a furnace called Blast Furnace. Commonly hematite</div> <div> and magnetite are used to obtain iron. A number of reactions occurs at high temperature in the</div> <div> blast furnace.</div> <div> The reactions occurring in blast furnace are as follows.</div> <div> • Reactions in blast furnace:</div> <div> Fe2O3 + CO 2Fe3O4 + CO2 (g)</div> <div> Fe3O4 + CO 3FeO +CO2 (g)</div> <div> FeO + CO Fe(l) + CO2 (g)</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 45</div> <div> Iron is a common metal. It has been known since prehistoric ages. Iron is soft, malleable, and</div> <div> ductile.</div> <div> Occurrence:</div> <div> Iron is a common metal. How it occurs?</div> <div> · Metallic iron occurs in the free state in only a few localities in the world.</div> <div> · In chemical compounds the metal is widely distributed.</div> <div> · After aluminum, iron is the most abundant of all metals, and ranks fourth in abundance</div> <div> among all the elements in the earth's crust.</div> <div> · It is found in meteorites, usually alloyed with nickel.</div> <div> · Iron is found in the earth crust in the form of ores. Its principal ore is hematite and other</div> <div> important ores are magnetite, siderite, and limonite.</div> <div> · Pyrite, FeS, the sulfide ore of iron, is not processed as an iron ore because it is too difficult</div> <div> to remove the sulfur.</div> <div> · Small amounts of iron occur in combination in natural waters, in plants, and as a</div> <div> constituent of blood.</div> <div> Physical properties of Iron metal:</div> <div> What are the common physical properties of iorn?</div> <div> · The atomic number of iron is 26, and atomic weight is 55.847.</div> <div> · Pure iron melts at about 1535° C and boils at 2750° C.</div> <div> · The specific gravity of iron is 7.86.</div> <div> · It is a metal which is easily magnetized at ordinary temperatures, but at about 790° C the</div> <div> magnetic property of iron disappears.</div> <div> Chemical Properties:</div> <div> Describe some common properties of iron.</div> <div> Iron reacts with a number of elements. For example, it make compounds with the halogens</div> <div> (fluorine, chlorine, bromine, iodine), sulfur, phosphorus, carbon, and silicon.</div> <div> Iron reacts with acids and displaces hydrogen from dilute acids.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 46</div> <div> It reacts with oxygen to form ferrosoferric oxide, Fe3O4.</div> <div> Rust: When exposed to moist air, iron undergoes deterioration or corrosion and forms a</div> <div> reddish-brown, flaky, hydrated ferric oxide commonly known as rust.</div> <div> When iron is dipped into concentrated nitric acid, it forms a layer of oxide that renders it</div> <div> passive—that is, it does not react chemically with acids or other substances.</div> <div> Give some uses of iron:</div> <div> Iron has a wide variety of applications. It is used in processed forms, such as wrought</div> <div> iron, cast iron, and steel.</div> <div> Commercially pure iron is used for the production of galvanized sheet metal and of</div> <div> electromagnets.</div> <div> Many compounds of iron are used in different purposes. Ferrous sulfate (FeSO4), which is</div> <div> also called green vitriol or copperas. It is used as a mordant in dyeing, as a tonic medicine,</div> <div> and in the manufacture of ink and pigments.</div> <div> Metallurgy of Iron: what are the reactions that take place in the Blast Furnace?</div> <div> · Combustion of Coke</div> <div> C(s) + O2(g) ® CO2(g) + heat</div> <div> CO2(g) + C(s) ® 2CO(g)</div> <div> · Reduction of Fe2O3</div> <div> 2Fe2O3(s) + 3C(s) ® 4Fe(l) + 3CO2(g)</div> <div> Fe2O3(s) + 3CO(g) ® 4Fe(l) + 3CO2(g)</div> <div> · Calcination</div> <div> CaCO3(s) + heat ® CaO(s) + CO2(g)</div> <div> · Slag formation</div> <div> CaO(s) + SiO2(s) ® CaSiO3(l) (slag)</div> <div> CaO(s) + Al2O3(s) ® Ca(AlO2)2(l)</div> <div> What are the Products which are formed in the Blast Furnace:</div> <div> Pig iron - 93-95% Fe, 3-5% C, 1% Si, 0.1- 0.3% P, <1% S</div> <div> Waste gases – CO2 and CO</div> <div> Slag – CaSiO3 and Ca(AlO2)2</div> <div> Pig Iron:</div> <div> Pig iron usually refers to the iron alloy with carbon content of the 2 to 4.3%, also known as cast</div> <div> iron. Besides carbon, pig iron also contains silicon, manganese and small amounts of sulfur and</div> <div> phosphorus. It cannot be forged, but can be cast. According to the presence of carbon under the</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 47</div> <div> different forms, pig iron can be divided into steel making pig iron, foundry pig iron and ductile</div> <div> iron.</div> <div> What is cast iron?</div> <div> This is pig iron melted with scrap Iron. It is the least pure of all forms of iron containing 93% Fe</div> <div> and 5% C.</div> <div> The cast iron with carbon content of less than 0.2% is called wrought iron or pure iron. The cast</div> <div> iron with carbon content of 0.2-1.7% is called cast steel. So, cast steel is a kind of special cast</div> <div> iron. More than 2% of content is called pig iron.</div> <div> What is wrought iron?</div> <div> It is the purest form of iron produced when impurities are removed. It contains 0.5% impurities.</div> <div> Wrought iron is very soft, plastic and easily deformed, but its strength and hardness are lower, so</div> <div> not widely used.</div> <div> What is steel?</div> <div> Steel is a common form of iron which usually contains 0.1 to 2% carbon. Steel has many types.</div> <div> Since the pure iron metal is not very strong, impurities are added into it.</div> <div> For example, carbon is added in iron to make steel which is much more stronger than iron.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 48</div> <div> Lecture 15</div> <div> Corrosion and its Prevention</div> <div> We all are aware of the rusting of iron. In the same manner most other metals also deteriorate</div> <div> with time. This process is called corrosion.</div> <div> Corrosion occurs in all substances. It is the chemical or electrochemical reaction between a</div> <div> material, usually a metal, and its environment that produces a deterioration of the material and its</div> <div> properties. However, the term is commonly used with reference to metals. It can be broadly</div> <div> defined as follows.</div> <div> “The deterioration of a metal as a result of its reaction with the environment is called corrosion”.</div> <div> Chemically corrosion is an oxidation process that occurs at the surface of a metal. The rusting of</div> <div> iron, tarnishing of silver, the development of a green coating on copper are well known forms of</div> <div> corrosion. Corrosion is a great cause of the damage of metallic articles. Bridges, ships, vehicles,</div> <div> and other metallic things suffer from corrosion badly.</div> <div> Sources of Corrosion:</div> <div> Following are the common causes or sources of corrosion.</div> <div> · Contact with atmosphere</div> <div> · Submersion in water or water solutions</div> <div> · Underground soil attack.</div> <div> · Contact with chemicals.</div> <div> There are two major factors which increase the rate of corrosion rapidly. They are;</div> <div> · Temperature</div> <div> · Moisture</div> <div> Rate of corrosion:</div> <div> Corrosion is a chemical process. It is an oxidation-reduction reaction that occurs at the surface of</div> <div> a metal. In this process the metal undergoes oxidation. Different metal corrode with different</div> <div> speeds. Gold is so stable that it does not corrode. Alkali metals like sodium and potassium</div> <div> corrode rapidly when exposed to atmosphere. Iron corrodes slowly, and silver, copper and tin</div> <div> corrode very slowly.</div> <div> Rusting of Iron:</div> <div> The corrosion of major concern is the rusting of iron. It is the cause of huge economic loss every</div> <div> year. From common person to big industries all suffer from corrosion in one way or the other.</div> <div> Now the question is what the rust is.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 49</div> <div> Rusting is a chemical reaction. When iron cones in contact with oxygen and moisture it is</div> <div> oxidized to hydrated ferric oxidized, Fe2O3.xH2O. This is rust. This substance is porous and</div> <div> flaky. So it is removed away continuously. Thus the whole metal is eaten away.</div> <div> Iron + Oxygen + Water ==> Hydrated Iron(III) oxide</div> <div> Protection of iron form Rusting:</div> <div> The protection of corrosion can be classified as follows:</div> <div> · Active corrosion protection</div> <div> · Passive corrosion protection</div> <div> · Permanent corrosion protection</div> <div> · Temporary corrosion protection</div> <div> Active corrosion protection:</div> <div> The aim of active corrosion protection is to influence the reactions which proceed during</div> <div> corrosion. It is possible to control the corrosion reaction in such a manner that corrosion is</div> <div> avoided. Examples of such an approach are the development of corrosion-resistant alloys and the</div> <div> addition of such materials to the metal which can inhibit corrosion.</div> <div> Passive corrosion protection:</div> <div> In passive corrosion protection, the metal is mechanically isolated from the damaging</div> <div> environment. For example by using protective layers, films or other coatings rusting can be</div> <div> avoided.</div> <div> Permanent corrosion protection</div> <div> The purpose of permanent corrosion protection methods is mainly to provide protection at the</div> <div> place of use. Machines are located, for example, in factory sheds and are thus protected from</div> <div> extreme variations in temperature, which are frequently the cause of condensation.</div> <div> Prevention of Corrosion:</div> <div> There are a number of methods which are used to avoid or slow down corrosion in a metal and</div> <div> rusting of iron in particular.</div> <div> Protective metallic coating:</div> <div> In this method a metal is coated with a thin layer of another metal which is resistant to corrosion.</div> <div> Rusting of iron can be prevented by coating it with tin, zinc or chromium. Tin plating is a wellknown</div> <div> technique. In this technique, a layer of tin is applied on the surface of ion container. Tin</div> <div> cans are good example. When zinc is coated on iron, the process is called galvanizing, and the</div> <div> iron coated with zinc is called galvanized iron. However, zinc is not used in food cans because it</div> <div> readily reacts with acids found in food stuff and fruit juices.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 50</div> <div> Coating with paints:</div> <div> Metal can be protected against corrosion by applying acoating of paints on their surfaces. For this</div> <div> purpose, any paint can be used. However, red lead and zinc chromate are commonly used. Grease,</div> <div> plastic emulsions, and enamels can also be used for this purpose. Oiling is a common household</div> <div> and small scale industrial method to protect small tools from rusting.</div> <div> Cathode Protection:</div> <div> Corrosion is actually an electrochemical process. During the corrosion reaction the metal plays</div> <div> the role of anode where oxidation occurs which eats out the metal. If we connect iron through a</div> <div> conductor to a more reactive metal in a manner that the electrical circuit is established, then</div> <div> oxidation will occur on the other metal and iron will be acting as cathode. This method will then</div> <div> protect iron from corroding. For example, if an underground iron pipe is connected to a zinc rod.</div> <div> The zinc rod will undergo corrosion and iron will not. This method can be used to protect ship</div> <div> and underground foundations of bridges from rusting.</div> <div> Alloying:</div> <div> Many metals can be made more resistant to corrosion by making their alloys. Steel, bronze, brass</div> <div> are metal alloys and are more stable than their pure metals.</div> <div> Stain less steel is an alloy of iron. A wide variety of stainless steels are made today for different</div> <div> purposes. The common stain less steel has the following composition. C = 0.18 %, Cr = 18 %, Ni</div> <div> = 8 %, rest is iron or Fe. Stainless steel is used to make surgical instrument and cutlery etc.</div> <div> Applications of Alloys:</div> <div> Alloys are very important, as they have special properties and are more viable and useful than the</div> <div> pure metals. They have greater resistance to corrosion. For example, steel is mire resistance to</div> <div> corrosion than pure iron.</div> <div> Many alloys also have useful magnetic properties. For example, Alnico has a strong magnetic</div> <div> properties. It is Fe-Al-Ni-Co.</div> <div> Very hard steel is used to make jaws of rock crushers. It has Manganese and Vanadium.</div> <div> Some alloys also have attractive colours, For example, Au-Cu is an alloy which has red or yellow</div> <div> colour.</div> <div> Tin plating and methods of tin plating:</div> <div> The art of applying tin on the surface of a metal is called tin plating. The purpose of tin plating is</div> <div> to protect the metal from corrosion. Tin plating of iron is very common.</div> <div> There are a number of ways to coat tin on iron.</div> <div> Electroplating is commonly used for large scale tin plating. This is an electrolytic process. For</div> <div> this purpose, an electrolytic cell is made in the cathode is made of iron sheet and tin is used as</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 51</div> <div> anode. The electrolytic solution contains a salt of tin such as tin sulfate or tin chloride. When</div> <div> electric current is passed, tin ions move to the iron cathode and are deposited on it. In this way a</div> <div> layer of tin is coated on the surface of iron sheet.</div> <div> Another method to coat tin layer on iron surface is hot-dip method. In this method, sheets of iron</div> <div> are dipped in molten tin.</div> <div> Classical method of tin plating is a manual technique. A metal tool or article is polished with tin</div> <div> with the help of a rag or brush. Traditionally, brass and copper utensils are tin plated by this</div> <div> method.</div> <div> Stainless steel:</div> <div> Steel is an alloy of iron. There are many types of steel. A very useful type of steel is the stainless</div> <div> steel. It is so called because it is resistant to corrosion or rusting. The composition of common</div> <div> stainless steel is as follows: chromium 18%, nickel 8% and carbon 0.18 %.</div> <div> The stainless steel is hard and has a high tensile strength. Why it is resistant to rusting? This is</div> <div> due to the thin film of chromium oxide which is developed on the surface due to the reaction of</div> <div> chromium with oxygen. This is hard and tough film and does not corrode.</div> <div> Applications of Stainless Steel:</div> <div> o It is used to make cutlery</div> <div> o It is used to make surgical instruments</div> <div> o It is also used to make different household utensils and containers.</div> <div> o It is also used to make automobiles parts, decoration aricles and electronic appliances.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 52</div> <div> Lecture 16</div> <div> Periodic Table</div> <div> In this lecture we will discuss two important elements, namely hydrogen and oxygen.</div> <div> Hydrogen:</div> <div> Hydrogen is unique element. Its atom is the smallest atom and has only 1 proton and 1 electron. It</div> <div> is generally placed at the top of group IA along with alkali metals but it is not a metal and it is not</div> <div> fully justified to place it in group IA.</div> <div> Hydrogen is the most abundant element in the universe. The stars are mostly composed of</div> <div> hydrogen. In the earth crust hydrogen is found in the form of compounds. The most abundant</div> <div> compound of hydrogen is H2O. Petroleum and natural gas also contains hydrogen. Hydrogen is</div> <div> also a part of organic compounds.</div> <div> Isotopes of hydrogen: Hydrogen exists in the form of three isotopes. They are ordinary hydrogen</div> <div> or protium, heavy hydrogen or deuterium, and radioactive hydrogen or tritium.</div> <div> Industrial Preparation of hydrogen:</div> <div> On large scale hydrogen is prepared by a number of process. Some of them are discussed below.</div> <div> • Electrolysis of Water:</div> <div> In this method, electricity is passed through water containing a small amount of sulfuric acid.</div> <div> Hydrogen is obtained at cathode. At anode oxygen is obtained.</div> <div> 2H2O 2H2 + O2</div> <div> • Coke and Steam Process:</div> <div> Coke (a form of carbon) is mixed with steam and strongly heated. Water gas is produced which is</div> <div> a mixture of carbon monoxide and hydrogen gas.</div> <div> How to separate H2 from CO ?</div> <div> 1. Liquefaction process: the b.p. of CO = -192 °C and b.p. of H2 = -253 °C. When this</div> <div> mixture is liquefied, CO first liquefies and thus separated leaving behind hydrogen gas.</div> <div> 2. Bosch Process: Then more steam is added and the mixture is heated strongly in the</div> <div> presence of ferric oxide. A mixture of carbon dioxide and hydrogen gas is obtained. This</div> <div> is known as Bosch process.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 53</div> <div> How to separate CO2?</div> <div> Carbon dioxide is separated by passing the mixture through water under pressure. Carbon dioxide</div> <div> is dissolved in water and hydrogen is liberated.</div> <div> · Steam and hydrocarbon process:</div> <div> In this process a mixture of methane and steam is passed over a nickel catalyst at about 900 C</div> <div> under high pressure. A mixture of carbon monoxide and hydrogen gas is obtained.</div> <div> CH4 + H2O CO + 3H2</div> <div> Carbon monoxide is separated by the method discussed above.</div> <div> Laboratory Preparation:</div> <div> In the laboratory, hydrogen gas can be prepared by the action of a dilute solution of a strong acid</div> <div> on reactive metals like zinc.</div> <div> Action of dil. Acids on metals:</div> <div> • Zn + 2HCl ZnCl2 + H2 (g)</div> <div> • Zn + H2SO4 ZnSO4 + H2 (g)</div> <div> Action of Water on CaH2</div> <div> CaH2 + 2H2O Ca(OH)2 + 2H2 (g)</div> <div> Chemical reactions of hydrogen:</div> <div> 1. Reaction with Oxygen: Burning of H2</div> <div> 2H2 + O2 2H2O + Heat</div> <div> 2. Reaction with N2</div> <div> 3H2 + N2 2NH3</div> <div> 3. Reaction with Sulfur</div> <div> H2 + S H2S</div> <div> 4. Reaction with Phosphorus</div> <div> 3H2 +2P 2PH3</div> <div> 5. Reaction with Chlorine</div> <div> H2 + Cl2 2HCl</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 54</div> <div> Applications of H2</div> <div> · Preparation of NH3 (heating H2 and N2 at 500 °C using Fe2O3/Al2O3 as catalyst. The</div> <div> process is called Haber process.</div> <div> · Preparation of HCl: Hydrogen gas is reacted with chlorine gas to obtain HCl.</div> <div> · Preparation of methanol: heating water gas with ZnO2/Cr2O3 catalyst:</div> <div> CO + 2H2 CH3OH</div> <div> · Hydrogenation of vegetable oils to ghee</div> <div> Hydrides:</div> <div> The binary compounds of hydrogen with other elements are called hydrides. Hydrogen reacts</div> <div> with almost all elements and forms binary compounds. Common examples are H2O, NH3, HX</div> <div> (X is a halogen), CH4 and other hydrocarbons.</div> <div> Ionic Hydrides: With alkali metals and calcium, hydrogen forms ionic hydrides.</div> <div> 2Li + H2 2LiH</div> <div> 2Na + H2 2NaH</div> <div> Ca + H2 CaH2</div> <div> They are ionic compounds and show typical properties of ionic compounds.</div> <div> Covalent Hydrides: The hydrides of p-Block elements are covalent hydrides.</div> <div> Common examples: CH4, NH3, H2O, HCl, PH3, H2S, HBr, HI.</div> <div> They are molecular. Have low b.p., m.p.</div> <div> Metallic hydrides: The hydrides of transition metals are called metallic hydrides.</div> <div> They are non-stoichiometric. They are interstitial compounds and they retain properties of the</div> <div> original metal.</div> <div> Complex hydrides: common examples are LiAlH4 NaBH4</div> <div> Nascent Hydrogen: The hydrogen at the time of its generatioin in a reaction is called nascent</div> <div> hydrogen. It is atomic hydrogen. It may react with any reactant available, or atomic hydrogen</div> <div> may produce hydrogen gas. For example:</div> <div> Zn + 2HCl ZnCl2 + 2[H]</div> <div> 2[H] H2</div> <div> Oxygen: Oxygen is a common element. In the earth crust, it is the most abundant element.</div> <div> Oxygen atom has 8 protons and 8 electrons. Common isotope of oxygen has 8 neutrons.</div> <div> Molecular formula of oxygen gas is O2. Oxygen gas is about 20% of the air by volume. Oxygen is</div> <div> necessary for life. Most organic compounds also contain oxygen.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 55</div> <div> Allotropes of Oxygen: there are two allotropes of oxygen:</div> <div> • Ordinary oxygen, dioxygen, oxygen gas, O2</div> <div> • Ozone, O3</div> <div> Formation of O2- ions: oxygen has high electron affinity, so it absorbs electrons. Absorption of</div> <div> first electron is exothermic while that of the second electron is endothermic.</div> <div> O + 2e- O2-</div> <div> Preparation of Oxygen:</div> <div> Laboratory preparation: By heating potassium chlorate using MnO2 as catalyst: heated at</div> <div> 240°C</div> <div> 2KClO3 2KCl + 3O2 (g)</div> <div> Industrial preparation: On the industrial scale, oxygen is prepared by the following methods.</div> <div> · Electrolysis of water: Electrolysis of water produces hydrogen and oxygen gases.</div> <div> · Liquefaction of air: The major components of air are oxygen and nitrogen. Their boiling</div> <div> points are given below.</div> <div> b.p. of O2 = -183°C</div> <div> b.p. of N2 = -196°C</div> <div> So these gases can be separated by liquefaction.</div> <div> Chemical properties:</div> <div> Oxygen is a reactive element and reacts with almost all metals and non-metals, the compounds</div> <div> the oxygen formed with other elements are of different types.</div> <div> 1. Reaction with C, S, P, H2 oxides are formed.</div> <div> C + O2 CO2</div> <div> Characteristics of Oxygen:</div> <div> • Abundance: most abundant in Earth crust, 50%</div> <div> • In the air: 21% by volume</div> <div> • Most abundant compounds: H2O, SiO2</div> <div> • In the periodic table, group VIA,</div> <div> • Electronegativity: 3.4.</div> <div> • Electronic configuration:</div> <div> • O2 is paramagnetic</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 56</div> <div> S + O2 SO2</div> <div> P4 + 5O2 P4O10</div> <div> 2H2 + O2 2H2O</div> <div> 2. With sodium</div> <div> The following reaction occurs at room temperature:</div> <div> 4Na + O2 2Na2O sodium oxide</div> <div> At higher temperature:</div> <div> 2Na + O2 Na2O2 sodium peroxide</div> <div> 3. With potassium</div> <div> Potassium reacts with oxygen to for superoxide:</div> <div> K + O2 KO2 potassium superoxide</div> <div> 4. With Mg, Ca and Ba</div> <div> 2Ca + O2 2 CaO</div> <div> 5. Reaction with Fe</div> <div> 2Fe + O2 2FeO</div> <div> 4FeO + O2 2Fe2O3</div> <div> 6. Reaction with CH4 (Combustion)</div> <div> CH4 and other hydrocarbons burn:</div> <div> CH4 + 2O2 CO2 + 2H2O + Heat</div> <div> Oxides: The binary compounds of oxygen:</div> <div> The binary compounds of oxygen with other elements called oxides. Oxides are of three types.</div> <div> 1. Normal oxides; in which the oxidation number of oxygen is -2.</div> <div> 2. Peroxides; in which the oxidation number of oxygen.</div> <div> 3. Super oxides; in which the oxidation number of oxygen is -1/2</div> <div> Normal oxides are very conmen. H2O, CO2, NO2 are its examples. H2O2 is hydrogen per oxide.</div> <div> KO2 is potassium super oxide.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 57</div> <div> Types of Oxides:</div> <div> 1. Covalent oxides or acidic oxides. They are oxides of nonmetals, CO2, H2O, NO2, SO2,</div> <div> etc. When dissolved in water, they form acids:</div> <div> CO2 + H2O H2CO3</div> <div> SO2 + H2O H2SO3</div> <div> SO3 + H2O H2SO4</div> <div> 2. Basic oxides: The basic oxides are the oxides of metals. They react with water to form</div> <div> bases:</div> <div> Na2O + H2O 2NaOH</div> <div> CaO + H2O Ca(OH)2</div> <div> BaO + H2O Ba(OH)2</div> <div> Interesting chemistry of KO2: How O2 gas is produced in rescue masks? It is done by the reaction</div> <div> of K2O with water:</div> <div> 4KO2 + 2H2O 4KOH + 3O2</div> <div> 2KOH + CO2 K2CO3 + H2O</div> <div> Applications of Oxygen gas:</div> <div> 1. Oxidizing agent: it is used in steel making, and is used to burn impurities.</div> <div> 2. To bleach pulp and paper (oxidation of coloured compounds usually gives colourless</div> <div> products).</div> <div> 3. Medical use (to ease breathing difficulties)</div> <div> 4. To produce oxy-acetylene flame</div> <div> 2C2H2 + 5O2 4CO2 + 2H2O</div> <div> 5. To produce oxy-hydrogen flame</div> <div> Compounds of Oxygen: Sulfuric Acid</div> <div> Preparation of H2SO4 by Contact Process involves the following reactions:</div> <div> 1. S + O2 SO2</div> <div> 2. 2SO2 O2 2SO3 Catalyst: V2O5</div> <div> 3. SO3 + H2SO4 + H2S2O7</div> <div> 4. H2S2O7 + H2O 2H2SO4</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 58</div> <div> Lecture 17</div> <div> Properties and Trends in Periodic Table</div> <div> Nitrogen is a common element, it occurs in the form of element and compounds. Nitrogen gas is</div> <div> composed of N2 molecules. Nitrogen gas is found 78% by volume in air. Nitrogen molecule N2 is</div> <div> very stable because there is a triple bond which binds nitrogen atoms together. In the triple bond,</div> <div> 1 bond is sigma and other 2 bonds are pi bonds.Nitrogen compounds are numerous; examples</div> <div> include ammonia, urea and nitric acid. Many organic compounds also contain nitrogen. Naturally</div> <div> occurring minerals of nitrogen are saltpeters KNO3 and NaNO3.</div> <div> Laboratory preparation of nitrogen gas: By heating NH4NO2 which is prepared by the action</div> <div> of ammonium chloride with sodium nitrite:</div> <div> NH4Cl + NaNO2 NH4NO2 + NaCl</div> <div> NH4NO2 N2 + 2H2O</div> <div> Large scale obtaining: On large scale nitrogen gas is obtained by liquefaction of air:</div> <div> b.p. of O2 = -183°C</div> <div> b.p. of N2 = -196°C</div> <div> Chemical Properties:</div> <div> 1. Reaction with oxygen:</div> <div> N2 + O2 2NO heating at 1500-2000 °C</div> <div> 2. Heating with Mg, Ca, B</div> <div> Nitrides are formed: BN (boron nitride), Ca3N2 (calcium nitride)</div> <div> 3. Reaction with H2 (Haber Process) produces ammonia (NH3).</div> <div> Applications of Nitrogen gas:</div> <div> 1. Nitrogen is used to prepare ammonia which is then used to prepare nitric acid and</div> <div> fertilizer.</div> <div> 2. It is used as an inert gaseous blanket to exclude O2 to a processing and packaging of</div> <div> foods.</div> <div> 3. Liquid nitrogen is used as a coolant to freeze foods rapidly.</div> <div> Ammonia, NH3</div> <div> Ammonia is a base. It is because it has a lone pair of electrons which can be donated to a proton,</div> <div> thus ammonia is electron pair donor or it is an proton acceptor.</div> <div> N2 + H+ NH4</div> <div> +</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 59</div> <div> Ammonia is a binary compound of nitrogen and hydrogen. Its molecular formula is NH3. It is</div> <div> prepared on industrial scale by the method called Haber process. In this process nitrogen is</div> <div> reacted with hydrogen in the presence of catalyst Fe2O3/ Al2O3.</div> <div> N2 + 3H2 2NH3</div> <div> The reactants are heated at 500 0C.</div> <div> Laboratory preparation of ammonia:</div> <div> In the laboratory ammonia can be prepared by the action of base on ammonium chloride.</div> <div> 1. NH4Cl + NaOH NH3 + NaCl + H2O</div> <div> 2. NH4Cl + Ca(OH)2 NH3 + H2O + CaCl2</div> <div> Nitric Acid, HNO3</div> <div> Nitric acid is an important compound of nitrogen. It is one of the common acids used extensively</div> <div> in laboratories and industries.</div> <div> Preparation of Nitric acid:</div> <div> Nitric acid is prepared industrially by a method known as Ostwald’s method.</div> <div> In this method nitric acid is prepared by the oxidation of ammonia.</div> <div> There are three main steps:</div> <div> 1. Oxidation of ammonia to nitric oxide</div> <div> 2. Oxidation of nitric oxide to NO2.</div> <div> 3. Reaction of NO2 with water to produce nitric acid.</div> <div> Chemical reactions are shown by the chemical equations as follows:</div> <div> 4NH3 + 5O2 4NO + 6H2O</div> <div> 2NO + O2 2NO2</div> <div> 3NO2 + H2O 2HNO3 + NO</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 60</div> <div> Lecture 18</div> <div> Halogens (VII)</div> <div> The elements of group VII are called halogens, which mean salt formers. Halogens include the</div> <div> following elements:</div> <div> 1. Fluorine (F)</div> <div> 2. Chlorine (Cl)</div> <div> 3. Bromine (Br)</div> <div> 4. Iodine (I)</div> <div> 5. Astatine (At)</div> <div> Atomic Numbers of Halogens are given below along with the name of the element:</div> <div> 1. F Fluorine (9)</div> <div> 2. Cl Chlorine (17)</div> <div> 3. Br Bromine (35)</div> <div> 4. I iodine (53)</div> <div> 5. At Astatine (85</div> <div> Occurrence: halogens occur in the form of different compounds in the earth crust.</div> <div> 1. Salts e.g., NaCl, Br- , I-</div> <div> 2. Fluorspar, CaF2</div> <div> 3. Cryolite, Na3AlF6,</div> <div> 4. Fluorapatite, Ca5(PO4)3F</div> <div> -------------------------</div> <div> Halogens form diatomic molecules: F2 Cl2 Br2 I2</div> <div> Physical States: In their elemental forms, halogens occur as diatomic molecules. They have</div> <div> different physical states. Fluorine and chlorine are gases, bromine is a liquid and iodine is a solid.</div> <div> Electronegativity: The electronegativity of an element is the ability of the element to attract the</div> <div> shared pair of electrons toward it. The halogens are very electronegative elements. Fluorine has</div> <div> the highest electronegativity among all the elements. The electronegativity values of halogens are</div> <div> given below:</div> <div> 1. F 4.0</div> <div> 2. Cl 3.0</div> <div> 3. Br 2.8</div> <div> 4. I 2.5</div> <div> Electron affinity: The electron affinity of an element is the energy absorbed or evolved when an</div> <div> electron enters into an atom of the element. A halogen atom absorbs an electron to make a</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 61</div> <div> negative ion (anion). During this process energy is evolved. Electron affinity values of halogens</div> <div> are given below:</div> <div> Chlorine gas (Cl2)</div> <div> In this lecture we will study chlorine as a typical halogen. Chlorine is a yellowish green coloured</div> <div> gas.</div> <div> Laboratory preparation: In the laboratory, chlorine gas is prepared by the action of</div> <div> hydrochloric acid on manganese dioxide.</div> <div> 4HCl + MnO2 MnCl2 + 2H2O + Cl2(g)</div> <div> Industrial preparation:</div> <div> On the industrial scale chlorine gas is prepared by sodium chloride.</div> <div> 1. Electrolysis of molten NaCl (Downs Cell Process):</div> <div> It may be prepared by the electrolysis of molten NaCl. For this purpose a special electrolytic cell</div> <div> is used which is called Downs cell and the process thus is called the Downs cell process. The</div> <div> melting point of NaCl is 801°C, which is a very high temperature. To lower the melting point of</div> <div> NaCl, calcium chloride is added. The electrolytic mixture in the Downs cell contains 40% NaCl</div> <div> and 60% CaCl2. Chlorine gas is obtained at anode.</div> <div> 2NaCl 2Na + Cl2(g)</div> <div> 2. Electrolysis of Brine (Castner-Kellner Cell Process): In this method aqueous solution of</div> <div> sodium chloride (common salt) is electrolyzed. The aqueous solution of NaCl is called brine.</div> <div> Chlorine gas is obtained at the anode. In the Castner-Kellner cell Moving mercury cathode is</div> <div> used. Sodium is discharge at the mercury electrode and forms sodium amalgam. The sodium</div> <div> amalgam is taken into another compartment where it react with water to produce NaOH and</div> <div> hydrogen gas. Mercury is recirculated.</div> <div> 2NaCl + 2H2O 2NaOH + H2(g) + Cl2(g)</div> <div> Physical properties:</div> <div> 1. Chlorine is a greenish yellow gas.</div> <div> 2. It has pungent smell.</div> <div> 3. Chlorine is a toxic gas. If it comes in contact, it produces inflammation of lungs, throat</div> <div> and nose.</div> <div> 4. Its boiling point is -34.6 C, and melting point is -101.6C.</div> <div> 5. It is fairly soluble in water.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 62</div> <div> Chemical properties of Chlorine gas:</div> <div> Chlorine is a reactive gas. It reacts with many metals and non-metals. It also reacts with</div> <div> compounds.</div> <div> Reactions with metals:</div> <div> 2Na + Cl2 2NaCl</div> <div> 2K + Cl2 2KCl</div> <div> 2Sb + 3Cl2 2SbCl3</div> <div> 2Fe + 3Cl2 2FeCl3</div> <div> Sn + 2Cl2 SnCl4</div> <div> Reactions with non-metals</div> <div> P4 + 6Cl2 4PCl3</div> <div> P4 + 10Cl2 4PCl5</div> <div> 2S + Cl2 S2Cl2</div> <div> Reaction with hydrogen:</div> <div> H2 + Cl2 2HCl</div> <div> Reactions with Compounds</div> <div> Reaction with Ammonia:</div> <div> 2NH3 + 3Cl2 N2 + 6HCl(g)</div> <div> 6NH3 + 6HCl 6NH4Cl(s)</div> <div> Oxidation of FeCl2 to FeCl3</div> <div> 2FeCl2 + Cl2 2FeCl3</div> <div> Chlorine replaces bromine and iodine from their compounds. It is a stronger oxidizing agent than</div> <div> bromine and iodine.</div> <div> 2NaBr + Cl2 2NaCl + Br2</div> <div> 2KI + Cl2 2KCl + I2</div> <div> Photochemical reactions: chlorine also undergoes photochemical reactions. In these reaction a</div> <div> photon of light breaks the chlorine molecule into chlorine free radicals which are chlorine atoms.</div> <div> The chlorine radicals then react with other compounds such as methane.</div> <div> CH4 + Cl2 CH3Cl + HCl</div> <div> Chlorine as a bleaching agent:</div> <div> Chlorine reacts with water to form HCl and HOCl. HOCl then breaks up to form atomic or</div> <div> nascent oxygen. The atomic oxygen is very reactive and is a strong oxidizing agent. It reacts with</div> <div> the spots and removes them. This is called bleaching action.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 63</div> <div> H2O + Cl2 HCl + HOCl</div> <div> 2HOCl 2HCl + 2[O] Nascent Oxygen</div> <div> Uses of Chlorine:</div> <div> Chlorine is a useful element. It is used for a variety of purposes.</div> <div> · For sterilizing drinking water</div> <div> · To prepare PVC (Polyvinyl chloride) plastic</div> <div> · To prepare CCl4 and CHCl3</div> <div> · To prepare HCl</div> <div> · To prepare bleaching powder, CaOCl2.H2O</div> <div> Hydrochloric Acid (HCl)</div> <div> One of the important compounds of chlorine is HCl which hydrogen chloride. It is a gas. When it</div> <div> is dissolved in water, we obtain hydrochloric acid. Hydrochloric acid is a commonly used acid.</div> <div> Hydrogen chloride can be prepared by a number of methods.</div> <div> Laboratory preparation: In laboratory HCl can be prepared by the action of sulfuric acid on</div> <div> sodium chloride.</div> <div> NaCl + H2SO4 NaHSO4 + HCl(g)</div> <div> Industrial preparation: On industrial scale, HCl is prepared by the reaction of chlorine with</div> <div> hydrogen gas.</div> <div> H2 + Cl2 2HCl</div> <div> Chemical reactions of HCl:</div> <div> Reaction with a base, NaOH or KOH</div> <div> HCl + NaOH NaCl + H2O</div> <div> HCl + KOH KCl + H2O</div> <div> Reaction with ammonia</div> <div> HCl + NH3 NH4Cl</div> <div> Reaction with calcium oxide</div> <div> 2HCl + CaO CaCl2 + H2O</div> <div> Reaction with sodium carbonate</div> <div> 2HCl + Na2CO3 2NaCl + CO2 + H2O</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 64</div> <div> Lecture 20</div> <div> Properties of Solutions</div> <div> Mixture:</div> <div> A mixture is a “mixed” substance which obtained on mixing two are more substances without</div> <div> undergoing any chemical reaction.</div> <div> Types of mixtures: Mixtures are of two types:</div> <div> · Homogenous Mixtures</div> <div> · Heterogeneous. Mixtures</div> <div> Homogeneous mixture</div> <div> A Homogeneous mixture has uniform composition throughout its body. They are known as</div> <div> solutions. For example, when sugar or salt is dissolved in water it forms a solution in which the</div> <div> sugar or salt and water are uniformly mixed (the sugar or salt is dissolved uniformly in water).</div> <div> Another common example is air.</div> <div> Heterogeneous mixture</div> <div> A Heterogeneous mixture is that which does not have uniform composition. e.g. granite which</div> <div> consists of quartz, feldspar and mica and you can identify each of the substances separately in the</div> <div> mixture.</div> <div> Heterogeneous mixtures can be suspensions or colloids.</div> <div> Solution:</div> <div> “A solution is a homogenous mixture of atoms, ions or molecules of two or more substances.”</div> <div> A homogenous mixture is that which has uniform composition throughout its body. In other</div> <div> words, it has only one phase. Suppose, we stir some sugar into a glass of water. The grains of</div> <div> sugar disappear and we have a clear liquid that looks just like pure water. The sugar molecules</div> <div> have dispersed among the water molecules, and sugar and water thus form a solution.</div> <div> Types of solution</div> <div> Solutions can be classified on the basis of their state: solid, liquid, or gas. Thus there are nine</div> <div> types of solutions. The most common type of solutions is a solid dissolved in liquids. The liquid</div> <div> is mostly water.</div> <div> Solvent and Solute</div> <div> The solvent is the component of a solution that is visualized as dissolving another component</div> <div> called a solute. Usually the component present in the larger quantity is called the solvent, and the</div> <div> component present in the smaller quantity is called the solute.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 65</div> <div> Concentration and its units</div> <div> By concentration we mean the relative amounts of the components of a solution. It tells us the</div> <div> ratio of the quantity of one component to the quantity of the other or to the total quantity of</div> <div> solution. It has many units. Some common units are discussed below.</div> <div> Percentage</div> <div> The ratio of the mass of the solute to the mass of the solution multiplied by 100 is called mass</div> <div> percentage.</div> <div> Mass Percentage of a solute = Mass of the solute x 100</div> <div> Mass of solution</div> <div> Example</div> <div> A solution contains 25 grams of sugar in 150 grams of water. What is the %age by mass of sugar</div> <div> in the solution? (14.3%)</div> <div> For liquid-liquid solutions, it is sometimes more convenient to express the concentration in the</div> <div> units of percentage by volume:</div> <div> Volume Percentage of a liquid = Volume of the liquid x 100</div> <div> Total volume</div> <div> Example</div> <div> A solution of 10 ml acetone and 50 ml water was prepared. Calculate the percentage by volume of</div> <div> acetone. (16.67%)</div> <div> Molarity (M)</div> <div> Molarity is the most common unit of concentration used in chemistry.</div> <div> “Molarity or the molar concentration is the number of moles of solute dissolved per dm3 of</div> <div> solution”. (1 dm3 is equal to 1 Liter)</div> <div> Thus</div> <div> Molarity = Number of moles of solute</div> <div> Volume of solution in liters</div> <div> Example: 15g of NaOH were dissolved in enough water to prepare 100 cm3 solution. Calculate</div> <div> the molarity of NaOH. (3.75 mol/dm3)</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 66</div> <div> Molarity (m)</div> <div> “Molarity is the number of moles of solute dissolved per kilogram of solvent”.</div> <div> Example: 15 g of NaOH was dissolved in 100 g of water. Calculate the molality of NaOH.</div> <div> (3.75 mol/Kg)</div> <div> Molality is independent of temperature as it is mass-to-mass ratio, while molarity is temperature</div> <div> dependent as it is mass to volume ratio.</div> <div> Solubility and Saturation</div> <div> This is an important question that how much a solute at maximum can be dissolved in a</div> <div> given quantity of a solvent. The maximum amount of a solute dissolved in a given amount of a</div> <div> solvent is called solubility of that solute. It can be precisely defined as follows:</div> <div> “The solubility of a solute in a given solvent is the concentration of that solute in its saturated</div> <div> solution at a specific temperature”.</div> <div> Solubility can be expressed in any of the units of concentration. However, it is commonly</div> <div> expressed in grams per dm3.</div> <div> “A saturated solution is the solution which can dissolve no more amount of the solute, at the</div> <div> given temperature”.</div> <div> If we start slowly adding sugar (sucrose) to water taken in a beaker. At first it dissolves rapidly,</div> <div> but then gradually the dissolution process becomes slow. The continued addition of sugar</div> <div> eventually brings the solution to saturation, after which any additional sugar simply falls to the</div> <div> bottom of the container where it remains undissolved. At this point, there is an equilibrium</div> <div> between the excess undissolved solute and the dissolved one.</div> <div> This solution of sugar is said to be saturated, and the concentration of sugar in this solution is</div> <div> called its solubility.</div> <div> Dissociation of Solutes</div> <div> When sugar dissolves in water, the solute particles that are dispersed through out the</div> <div> resulting solution are sugar molecules. On the other hand many solids (electrolytes) undergo</div> <div> dissociation or ionization when dissolve. For example, NaCl forms sodium and chlorine ions in</div> <div> solution.</div> <div> The solubility of NaCl at 25oC have been estimated to be 5 mol/dm3 or 292.5 g/dm3.</div> <div> Sparingly Soluble Substances</div> <div> The substances that have very low solubility are known as sparingly soluble or slightly soluble</div> <div> substances. Example, AgCl.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 67</div> <div> Hydration</div> <div> When an ionizable substance (an acid, base or salt) is added to water to form a solution, it breaks</div> <div> up into ions. The ions are surrounded or encaged by water molecules due to attraction between</div> <div> the ions and water molecules. This process is known as hydration.</div> <div> Water is a polar molecule in which the hydrogen atoms are partially positively charged and the</div> <div> oxygen atom is partially negatively charged. Being polar, water molecules surround the ions</div> <div> dipped in it. Thus, in water solution all cations and anions are hydrated. The cations are attracted</div> <div> by the negative ends (oxygen) of water molecules and the anions are attracted by the positive ends</div> <div> of the water molecules (hydrogens).</div> <div> Water of hydration</div> <div> Water can become incorporated in solid compounds by occupying specific crystal sites in</div> <div> the solid without being strongly bonded to a specific atom. This water is called water of hydration</div> <div> or water of crystallization. The compounds having water of hydration are called hydrates. Some</div> <div> examples are CaSO4.2H2O, CuSO4.5H2O, CaC12.6H2O. Many hydrates lose their water of</div> <div> hydration upon heating and form anhydrous compounds. For example, crystalline copper sulfate</div> <div> pentahydrate (blue) changes into anhydrous copper sulfate (white) on heating.</div> <div> Hydrolysis</div> <div> “The reaction of water with an anion or a cation or both of a salt causing a change in the pH is</div> <div> called hydrolysis”.</div> <div> Colloids:</div> <div> Colloids form the dividing line between solutions and heterogeneous mixture. They are not true</div> <div> solution. In colloids the disperse particles are larger than typical molecules. In colloids we don’t</div> <div> have solute or solvent. Rather, we have dispersed phase and dispersion medium colloids scattered</div> <div> light. This is called Tyndal effect.</div> <div> Solution Colloid</div> <div> i. This is the homogeneous and stable</div> <div> mixture of a solute in solvent.</div> <div> i. This is the heterogeneous and stable</div> <div> mixture of particles in the solvent.</div> <div> ii. Particle sizes of solutes are very small</div> <div> does not visible under light microscope.</div> <div> ii. Particles are larger sized and visible</div> <div> through the light microscope.</div> <div> iii. Solute evenly distributed does not</div> <div> show. Brownian movement, tyndal effect</div> <div> etc.</div> <div> iii. Particles are remaining scattered does</div> <div> not steeled down longer perion. They</div> <div> show Brownian motion and exhibit</div> <div> tyndal effect.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 68</div> <div> What is a saturated solution?</div> <div> A saturated solution is that solution which cannot dissolve more solute at a given temperature. If</div> <div> in such a solution you add more solute at that temperature, it will not dissolve and will settle</div> <div> down at the bottom. Thus, the solution which can not dissolve more solute at a particular</div> <div> temperature is called a saturated solution.</div> <div> What is solubility?</div> <div> The amount of solute needed to form a saturated solution in a given quantity of solvent at specific</div> <div> temperature is called the solubility of that solute in that solvent at temperature. In other words</div> <div> solubility is the maximum amount of solute that will dissolve in a given amount of solvent at a</div> <div> particular temperature.</div> <div> How molarity is different from molality?</div> <div> Molarity is a unit of concentration. It is the number of moles of solute present per litter of</div> <div> solution.</div> <div> Molality is another unit of concentration. It is the number of moles of solute present per Kg of</div> <div> solvent. So, molarity is mass per volume whereas molality is mass per mass. Therefore, molarity</div> <div> varies with temperature but there is no effect of temperature on molality.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 69</div> <div> Lecture # 21</div> <div> Acids, Bases</div> <div> Arrhenius concept of acid and base:</div> <div> According to Arrhenius theory an acid is substance which produces H+ ions when dissolved in water,</div> <div> examples; HCl, HNO3 etc.</div> <div> HCl H+ + Cl-</div> <div> According to this theory a base is a substance which produces OH- when dissolved in water.</div> <div> Examples include NaOH and KOH.</div> <div> NaOH Na+ + OHWhen</div> <div> an acid and a base react, they cancel each other. This reaction is known as neutralization and in</div> <div> neutralization reaction a salt and water is formed. Example;</div> <div> HCl + NaOH NaCl + H2O</div> <div> Acid Base Salt Water</div> <div> Some bases Their uses</div> <div> Ammonia (NH3)</div> <div> Production of fertilizers (ammonium and nitrate salts), used in the</div> <div> manufacture of nitric acid, neutralize the acid (in the petroleum industry)</div> <div> and prevent premature coagulation in natural / synthetic latex.</div> <div> Aluminium hydroxide</div> <div> (Al(OH)3)</div> <div> Manufacture other aluminium compound and to make gastric medicine</div> <div> (antacid)</div> <div> Calcium hydroxide</div> <div> (Ca(OH)2)</div> <div> To make cement, limewater, neutralize the acidity of soil and application</div> <div> of sewage treatment.</div> <div> Sodium hydroxide</div> <div> (NaOH)</div> <div> Used in the manufacturing of soaps, detergents, and cleaners.</div> <div> Lowry-Bronsted concept of acid and base:</div> <div> According to Lowry-Bronsted concept, an acid is a proton donor and a base is a proton acceptor.</div> <div> When acid-base reaction occurs, another set of acid and base is formed. They are call d</div> <div> conjugate acids and bases.</div> <div> Example: When HCl is dissolved in water the following reaction occurs.</div> <div> HCl + H2O H3O+ + Cl-</div> <div> Hydronium (H3O+) ion is a conjugate acid whereasCl- ion is conjugate base.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 70</div> <div> Reaction of acids with metals:</div> <div> Al + HCl AlCl3 + H2</div> <div> Mg + H2SO4 MgSO4 + H2s</div> <div> Fe + H2SO4 FeSO4 + H2</div> <div> Acid-base reactions:</div> <div> NaOH + HNO3 NaNO3 + H2O</div> <div> KOH + CH3COOH CH3COOK + H2O</div> <div> Reactions of acids with metal oxides:</div> <div> Al2O3 + 6HCl 2AlCl3 + 3H2O</div> <div> 3K2O + 2H3PO4 2K3PO4 + 3H2O</div> <div> Acid and Base Strength</div> <div> According to Arrhenius definitions, an acid furnishes H+ ions and a base furnishes OH- ions in</div> <div> their aqueous solutions. For example:</div> <div> HCI(g) H+ (aq) + CI- (aq)</div> <div> NaOH (s) Na+ (aq) + OH- (aq)</div> <div> However, all acids, and all bases are not equally strong. The strength of an acid or a base is</div> <div> estimated by a number of parameters, such as:</div> <div> 1. Degree of dissociation</div> <div> 2. Percent degree of dissociation</div> <div> 3. Dissociation constant</div> <div> Degree of dissociation</div> <div> It tells us how fast an acid or base undergo dissociation of ionization when they are put in water.</div> <div> Degree of dissociation is the ratio of the number of the moles or molecules dissociated and the</div> <div> total number of the moles or molecules initially taken. Thus</div> <div> The higher the degree of dissociation of an acid or a base, the stronger they are. The maximum</div> <div> value of the degree of dissociation is 1.</div> <div> Percentage Degree of Dissociation</div> <div> Percent degree of dissociation is obtained by multiplying the degree of dissociation by 100.</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 71</div> <div> Strong acids and bases are almost 100% dissociated in water, while weak acids and bases are not</div> <div> 100% dissociated. HCI, HNO3, NaOH and KOH are almost 100% dissociated in dilute water</div> <div> solution, while CH3COOH is only 1.4% and NH4OH is again 1.4% dissociated.</div> <div> Dissociation Constant (Ka and Kb)</div> <div> According to Lowry-Bronsted concept an acid is a proton donor. Thus, in general, an acid, HA,</div> <div> forms the following equilibrium in water:</div> <div> HA + H2O A− + H3O+</div> <div> The equilibrium constant expression for this reaction is:</div> <div> Water is present in large amount so its concentration is almost constant. Thus, the product Kc</div> <div> [H2O] is also constant and known as the dissociation constant of an acid and is given the symbol</div> <div> Ka.</div> <div> Strong acids, such as HCI, HNO3 have very large values of dissociation constants. While weak</div> <div> acids have very small values of Ka, for example:</div> <div> CH3COOH Ka = 1.8 x 10-5</div> <div> HCN Ka = 4.9 x 10-10</div> <div> H2O Ka = 1.0x10-14</div> <div> According to Lowry-Bronsted concept a base is a proton acceptor. Thus, in general, a base, B,</div> <div> reacts with water as follows:</div> <div> B + H2O BH+ + OHpH</div> <div> and pOH</div> <div> Pure water is only about 2x10-7 percent dissociated (self-dissociation) into ions at 25oC.</div> <div> H2O H+ (aq) + OH- (aq)</div> <div> Then, the equilibrium constant, Kc, is</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 72</div> <div> Since water is in excess, [H2O] is constant:</div> <div> Kc [H2O] = [H+][OH-] = Kw</div> <div> Thus, Kw = [H+][OH-]</div> <div> Kw is known as ionic product of water. Its value has been experimentally found as:</div> <div> Kw = 1x10-14 mol2 dm-6</div> <div> Or Kw = 1 x 10-14 mol2 lit-2</div> <div> In pure water (or in a neutral aqueous solution)</div> <div> [H+] = [OH-] = 10-7 mol/dm3</div> <div> The H+ ion concentration in solution can range from more than 10 mol/dm3 to less than 1x10-15</div> <div> mol/dm3. The pH scale was devised as a convenient way of expressing such a wide range of</div> <div> concentration. The pH is defined as follows:</div> <div> “The pH of a solution is the negative logarithm to the base</div> <div> 10 of the hydrogen ion (or hydronium ion) concentration”.</div> <div> Mathematically,</div> <div> pH = −log10[H+]</div> <div> or pH = - log10 [H3O+]</div> <div> The concept of the pH was proposed in 1909 by the Danish chemist Soren Sorenson (1868-1939).</div> <div> In the same way we can also define pOH.</div> <div> pOH = - log [OH-]</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 73</div> <div> Table: 7.1 pH, pOH, (25oC), and H+ ion, oH ion concentration</div> <div> [H+], M pH [OH], M pOH</div> <div> More acidic 10</div> <div> 1</div> <div> 10-1</div> <div> 10-2</div> <div> 10-3</div> <div> 10-4</div> <div> 10-5</div> <div> -1</div> <div> 0</div> <div> 1</div> <div> 2</div> <div> 3</div> <div> 4</div> <div> 5</div> <div> 10-15</div> <div> 10-14</div> <div> 10-13</div> <div> 10-12</div> <div> 10-11</div> <div> 10-10</div> <div> 10-9</div> <div> 15</div> <div> 14</div> <div> 13</div> <div> 12</div> <div> 11</div> <div> 10</div> <div> 9</div> <div> More acidic</div> <div> Neutral 10-7 7 10-7 7 Neutral</div> <div> More basic 10-8</div> <div> 10-9</div> <div> 10-10</div> <div> 10-11</div> <div> 10-12</div> <div> 10-13</div> <div> 10-14</div> <div> 10-15</div> <div> 8</div> <div> 9</div> <div> 10</div> <div> 11</div> <div> 12</div> <div> 13</div> <div> 14</div> <div> 15</div> <div> 10-6</div> <div> 10-5</div> <div> 10-4</div> <div> 10-3</div> <div> 10-2</div> <div> 10-1</div> <div> 1</div> <div> 10</div> <div> 6</div> <div> 5</div> <div> 4</div> <div> 3</div> <div> 2</div> <div> 1</div> <div> 0</div> <div> -1</div> <div> More basic</div> <div> At 25 oC,</div> <div> [H+][OH-] = 1 x 10-14</div> <div> pH + pOH = 14</div> <div> Thus,</div> <div> The pH of a neutral solution is 7, the pH of acids is less than 7 while that of bases is higher than 7.</div> <div> ------------------------------------------------</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 74</div> <div> Problem:</div> <div> What is the pH of a 0.016 mol/dm3 solution of NaOH?</div> <div> Solution:</div> <div> NaOH is a strong base and dissociates almost 100% to form OH- ions, thus</div> <div> [OH-] = [NaOH]</div> <div> = 0.016 mol/dm3</div> <div> Therefore,</div> <div> pOH = - log [OH-]</div> <div> = - log 0.016</div> <div> = 1.8</div> <div> pH + pOH = 14</div> <div> pH = 14 - pOH</div> <div> = 14 – 1.8</div> <div> = 12.2</div> <div> _______________________________________________</div> <div> Measurement of pH</div> <div> The pH of a solution can be determined by two general methods:</div> <div> 1. Indicators 2. pH meters</div> <div> Indicators</div> <div> Indicators are chemical compounds which change colour with pH and thus give an</div> <div> approximate value of the pH of a given solution. There are a number of indicators. A special</div> <div> indicator is the universal indicator, which is a mixture of several indicators that has several</div> <div> colour changes over a wide range of pH.</div> <div> pH Meters</div> <div> A pH meter is an electrical instrument that employs two electrodes to measure the</div> <div> potential difference between the solution to be tested and a standard solution of known pH. This</div> <div> potential difference is related to the pH of the solution being tested. The pH is read directly from</div> <div> the dial of the instrument.</div> <div> Common Indicators:</div> <div> To determine the end-point in an acid-base titration or to determine the pH of a solution,</div> <div> various compounds are used, that are called indicators. Broadly speaking, “an indicator is any</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 75</div> <div> compound which changes colour with the change in pH”. An acid-base indicator is a weak acid</div> <div> (or base) whose colour is different from that of its conjugate base (or acid). Take the example of</div> <div> phenolphthalein (C20H13O4H):</div> <div> C20H13O4H (aq) C20H13O4</div> <div> - (aq) + H+ (aq)</div> <div> (colourless) (red) conjugate base</div> <div> Table: Some indicators and their colour changes</div> <div> Indicators Approximate pH range Corresponding coloru change</div> <div> Methyl orange</div> <div> Bromophenol blue</div> <div> Methyl red</div> <div> Phenolphthalein</div> <div> 3.1 – 4.4</div> <div> 3.8. - 4.6</div> <div> 4.4 – 6.2</div> <div> 8.0 – 10.0</div> <div> Orange to yellow</div> <div> Yellow to blue</div> <div> Red to yellow</div> <div> Colourless to red</div> <div> A paper strip impregnated with a universal indicator is often used. This is known as pH paper.</div> <div> When a part of this strip is dipped in a solution, the paper shows a colour from which the pH can</div> <div> be judged.</div> <div> Neutralization</div> <div> When an acidic solution is added to a basic solution in the right proportion, the properties of</div> <div> each are lost, and a salt and water are formed. This process is called neutralization.</div> <div> “The reaction between an acid and a base that forms</div> <div> a salt and water is called neutralization”.</div> <div> For Example,</div> <div> HCI (aq) + NaOH (aq) NaCI (aq) + H2O</div> <div> HNO3 (aq) + NaOH (aq) NaNO3 (aq) + H2O</div> <div> Buffers or Buffer Solutions</div> <div> Normally, solutions change their pH when an acid or a base is added to them. For example,</div> <div> pure water, when kept open, changes its pH from neutral to acidic due to the absorption of CO2</div> <div> from air forming carbonic acid (H2CO3).</div> <div> However, there are certain solutions, which tend to maintain their pH. These are known as</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 76</div> <div> buffers or buffer solutions.</div> <div> Definition:</div> <div> “A buffer solution is a solution which undergoes only a slight change of pH when an acid or a base</div> <div> is added to it”.</div> <div> A buffer offers resistance against any change in its pH. Most commonly:</div> <div> “A buffer solution is a solution of a weak acid and a salt of it, or a buffer solution is a</div> <div> solution of a weak base and a salt of it.”</div> <div> For example, a solution of acetic acid and sodium acetate forms a buffer: (CH3COOH +</div> <div> CH3COONa) solution and a solution of ammonium hydroxide and ammonium chloride forms a</div> <div> buffer: (NH4OH + NH4CI) solution</div> <div> Buffers maintain their pH upon addition of acids or bases by shifting equilibrium position.</div> <div> Buffers play a very important role in many chemical and biological processes. For</div> <div> example, the control of the pH is very important in the treatment of sewage in electroplating,</div> <div> and in the manufacturing of photographic materials. Maintaining a constant pH is essential for</div> <div> many metabolic processes; the pH of the blood has to be close to 7.4 and the pH of saliva is</div> <div> close to 6.8.</div> <div> Oxidation and Reduction</div> <div> Many reactions involve transfer of electrons from one atom to another atom. These reactions</div> <div> are known as oxidation-reduction reactions. The loss of electrons from an atom is known as</div> <div> oxidation, while the gain of electrons by an atom is called reduction. Oxidation and reduction</div> <div> always take place simultaneously and the extent of oxidation is equal to the extent of reduction.</div> <div> That is, the loss of electrons is equal to the gain of electrons. The substance that undergoes</div> <div> oxidation is known as the reducing agent, and the substance that undergoes reduction is known as</div> <div> the oxidizing agent.</div> <div> Oxidation Number</div> <div> A very important question is how to decide whether a given reaction is an oxidation- reduction</div> <div> reaction or not. To recognize such reactions, we can use the concept of oxidation number, also</div> <div> known as oxidation state:</div> <div> “The oxidation number of an atom is an apparent charge that is assigned to it on the basis of</div> <div> certain rules.”</div> <div> For compounds, the oxidation number may be defined as follows:</div> <div> “Oxidation number of an atom in a compound is the apparent charge on the atom if both electrons</div> <div> of each bond are arbitrarily assigned to the atom of higher electronegativity”.</div> <div> The oxidation number is not the real electrical charge. It is the charge that would be associated</div> <div> ©Copy right Virtual University of Pakistan</div> <div> 77</div> <div> with a particular atom if we assign both the electrons of each bond to the atom of higher</div> <div> electronegativity. The oxidation number is the measure of the extent or degree of oxidation</div> <div> of an element in its compounds.. ab</div> </div> <div> <br /></div> <div> <b><a href="http://www.genrica.com/vustuff/BIO102/BIO102_handouts_1_21.pdf" target="_blank"> Basic II-Chemistry (BIO102) book HANDOUTS in pdf</a></b></div>

Basic II-Chemistry (BIO102) book HANDOUTS in pdf

Basic II-Chemistry (BIO102) book HANDOUTS in pdf Handouts | Lectures | Contents | Books ©Copy right Virtual University of Pakistan...

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