Physical Chemistry Notes - PDF

Physical Chemistry By: Dr. Heba Salah Mousa I. Matter: Properties and measurements. States of matter Table (1): Different characteristics of matter Matter properties and measurements 1 Physical Chemistry...

Physical Chemistry By: Dr. Heba Salah Mousa I. Matter: Properties and measurements. States of matter Table (1): Different characteristics of matter Matter properties and measurements 1 Physical Chemistry By: Dr. Heba Salah Mousa A. Physical properties Examples of Physical properties: 1. Color (intensive). 2. Density (intensive). 3. Volume (extensive). 4. Mass (extensive). 5. Boiling point (intensive): temperature at which substance boils. 6. Melting point (intensive): temperature at which substance melts B. Chemical properties Examples of chemical properties: 1. Heat of combustion (∆Hc): Energy released upon complete combustion (burning) of compound with oxygen. 2. Stability: refers to reactions the alter chemical structures of compounds such as oxidation (reaction with oxygen), hydrolysis (reaction with water) and photosensitivity (decomposition by light). 3. Flammability: ability of compound to burn when exposed to flame. Commonly high temperature in presence of oxygen. 2 Physical Chemistry By: Dr. Heba Salah Mousa 4. Oxidation-Reduction: oxidation refers to loss of electrons while reduction is gain of electrons. 5. Chemical change or chemical reaction: process that cause a substance to change into a new substance with a new chemical formula. A+B (reactants) C (product) Units of Measurement All systems of weights and measures, metric and non-metric, are linked through a network of international agreements supporting the International System of Units 3 Physical Chemistry By: Dr. Heba Salah Mousa All SI units can be expressed in terms of standard multiple or fractional quantities, as well as directly. Multiple and fractional SI units are defined by prefix multipliers according to the powers of 10 ranging from 10 -24 to 10 24. There are seven basic units in SI system Table (2): Different basic units of SI system: 4 Physical Chemistry By: Dr. Heba Salah Mousa Table (3): Standard prefixes for SI units of measure: 1) Units of Mass ✓ The kilogram (Symbol, kg) is the SI unit of mass. ✓ It is defined as: The mass of a particular international prototype made of platinum-iridium and kept at the International Bureau of Weights and Measures. ✓ Variants on the base unit are formed by 1000 times increases or divisions of the gram. ✓ Care must be exercised in carrying out these calculations as errors may not be immediately obvious. 5 Physical Chemistry By: Dr. Heba Salah Mousa Notes In order to Avoid confusion between the abbreviations for microgram and nanogram, It is recommended that, the units of the weight are usually written in full, whereas possible. For example: 3 nanogram rather than 3 ng. micrograms rather than 5.8 µg. This is because the first letter of the abbreviation (the m or n) when written by hand may be mistaken for the letter ‘m’. This could then result in dosing errors of one-thousand times plus. Quantities should be written as whole numbers For example: 30 mg not 0.03 g 100 microgram not 0.1 mg The decimal point should always be preceded by a zero For example: 0.6 g not.6 g. 0.2 Kg not.2 Kg. 2) Units of Amounts of Substance ✓ The mole (abbreviation, mol) is the SI unit of the quantity of the material (amount of substance). ✓ One mole is the number of atoms in 0.012 kilogram of the most common isotope of elemental carbon (C-12). ✓ Amount of substance conversions are easy as they are all based around 1000 times multiplications or divisions of the base unit. 6 Physical Chemistry By: Dr. Heba Salah Mousa 3) Units of Length ✓ The meter (abbreviation, m) is the SI unit of length. ✓ One meter is the distance traveled by a ray of electromagnetic (EM) energy through a vacuum in 1/299,792,458 (3.33564095 x 10 -9) second. 4) Units of Capacity (Volume) ✓ Liquids are universally measured in litres. ✓ Although not part of the SI system of units, the litre is commonly used. ✓ A litre is defined as one cubic decimeter (0.001 m3), i.e. the volume that a cube with equal sides of 10 cm would occupy. 7 Physical Chemistry By: Dr. Heba Salah Mousa ✓ volume conversions are easy as they are all based around 1000 times multiplications or divisions of the base unit. 5) Units of Concentration The concentrations of electrolyte solutions for parenteral use are expressed as the amount of substance concentration (mol per L). The concentrations of substances normally found in blood, electrolytes, glucose, ceratinine, urea, etc. are expressed as the amount of substance concentration (mol per L). An exception: the amount of haemoglobin is expressed as g per 100 mL. 8 Physical Chemistry By: Dr. Heba Salah Mousa 9 Physical Chemistry By: Dr. Heba Salah Mousa ❖ In order to convert the moles of a substance to grams, you will need to multiply the mole value of the substance by its molar mass. No. of moles= no. of grams/molar mass Molar mass= x grams/1mole 1. How many grams are in 0.572 moles of glucose, C6H12O6? [ C=12, O= 16 and H=1] Answer Molecular weight= 12×6+ 12×1+ 16×6= 180 g/mol No. of grams= 0.572×180/1mol= 102.96 g C6H12O6 2. How many grams are in 3.79 moles of calcium bromide, CaBr2? (M.wt= 199.88) Answer 3.79×199.88/1 mole= 758 grams CaBr2 10 Physical Chemistry By: Dr. Heba Salah Mousa Relationship between density and volume ✓ The volume of substance related to quantity of substance at definite pressure and temperature. ✓ The volume of substance can be measured in volumetric or graduated measurements. ✓ Density is the amount of substance contained in a definite volume ✓ Density is used to define any substance. ▪ ρ (rho)= m/v where m is mass v is volume ✓ The density of a material varies with temperature and pressure ✓ This variation is typically small for solids and liquids but much greater for gases ✓ Increasing the pressure on an object decreases the volume of the object and thus increases its density. ✓ Increasing the temperature of a substances decreases its density by increasing its volume. ✓ Increasing mass will increase density except water is unusual case where its solidification (ice) becomes lighter than liquid state (water) and floats. ✓ Specific gravity is ratio between density of substance to density of reference material (water) at constant volume. Table (4): Different densities of water at different degrees. 11 Physical Chemistry By: Dr. Heba Salah Mousa Temperature measurement ✓ Measurement of temperature should be against accepted standards. ✓ Temperature can be measured in several scales as Celsius, Kelvin and Fahrenheit ✓ Temperature is a measure of cold or heat and measured by Thermometer. Comparison of temperature scale Table (5): Calculation of temperature by various scales. Dimensional analysis (unit conversion) A method of converting one unit to another. List of some commonly derived units. 12 Physical Chemistry By: Dr. Heba Salah Mousa II. Atomic structure History of atomic structure ❑ The word atom comes from the ancient Greek adjective atomos, meaning "indivisible”. ❑ The idea of an indivisible particle was explained by a number of philosophers and scientists such as Galileo, Dalton, Newton, Boyle and Lavoisier. ❑ John Dalton is the first chemist that postulates the modern atomic theory based on his experiments on gases. ❑ Atom is smallest constituent of matter and consists of nucleus and surrounding electrons. ❑ Every atom is composed of a nucleus and one or more electrons bound to the nucleus. ❑ The nucleus is made of one or more protons and typically a similar number of neutrons. ❑ Protons and neutrons are called nucleons. More than 99.94% of an atom's mass is in the nucleus. ❑ The protons have a positive electric charge, the electrons have a negative electric charge, and the neutrons have no electric charge. ❑ If the number of protons and electrons are equal, that atom is electrically neutral. ❑ If an atom has more or fewer electrons than protons, then it has an overall negative or positive charge, respectively, and it is called an ion. 13 Physical Chemistry By: Dr. Heba Salah Mousa Dalton’s atomic theory ✓ These ideas based on his studies on gases: A. All atoms of a given element are identical. B. Atoms of different elements vary in size and mass. C. Atoms are indivisible and the chemical reaction leads to rearrangement of atoms not to their creation or destruction. Dalton outlined law of multiple proportions that describes how reactants combine in set ratio e.g. Dalton knew that the element carbon forms two oxides by combining with oxygen in different proportions carbon monoxide (CO) and carbon dioxide (CO2). Periodic table 14 Physical Chemistry By: Dr. Heba Salah Mousa Figure (1): List of all metals with their atomic numbers in periodic table. Isotopes 15 Physical Chemistry By: Dr. Heba Salah Mousa ✓ Despite having different numbers of neutrons, isotopes of the same element have very similar physical properties. ✓ Some isotopes are unstable and will undergo radioactive decay to be other element. ✓ The predictable half-lives of isotopes allows scientist to date materials based on its isotopic composition such as dating with C-14 (The longest-lived radioisotope with a half-life of 5,700 years). ✓ Half-life of isotope is the time it takes for half of original concentration of an isotope to decay back to its more stable form. ✓ Radioactive carbon determines age of an object by comparing ratio of C-14 amount found in it to amount in atmosphere. Molecules 16 Physical Chemistry By: Dr. Heba Salah Mousa Figure (2): 3D of caffeine molecule showing its atoms and bonds. Isomers 17 Physical Chemistry By: Dr. Heba Salah Mousa 18 Physical Chemistry By: Dr. Heba Salah Mousa III. Introduction to bonds ✓ Chemical bonds join atoms together to form more complex structures (like molecules or crystals). ✓ Bonds can form between atoms of the same element, or between atoms of different elements. ✓ There are several types of chemical bonds which have different properties and give rise to different structure. ✓ These types include ionic, covalent, hydrogen, metallic and co-ordination bonds. 1. Ionic bond: ✓ It is formed between positive ions (cations) and negative ions (anions). ✓ In an ionic solid, the ions arrange themselves into a rigid crystal lattice. NaCl (common salt) is an example of an ionic substance. ✓ When ionic bonds form, there is an attractive force established between the positive cation and the negative anion. This attraction between oppositely- charged ions is the ionic bond ✓ Generally, when metals react with non-metals, electrons are transferred from the metals to the non-metals. The metals form positively-charged ions and the non-metals form negatively-charged ions. ❑ Characteristics of ionic bonds are: 1. High melting point (solid at room temperature). 2. Hard but brittle. 3. Many dissolve in water. 4. Conductors of electricity when dissolved or melted. Figure (3): Lewis structure of the ionic bond between sodium and chlorine. 19 Physical Chemistry By: Dr. Heba Salah Mousa How ionic bond formed? ✓ Ionic bonds form when metals and non-metals chemically react. By definition, a metal is relatively stable if it loses electrons to form a complete valence shell and becomes positively charged. ✓ Likewise, a non-metal becomes stable by gaining electrons to complete its valence shell and becomes negatively charged. When metals and non-metals react, the metals lose electrons by transferring them to the non-metals, which gain them. Consequently, ions are formed, which instantly attract each other—ionic bonding. ✓ For instance, in the reaction of Na (sodium) and Cl (chlorine), each Cl atom takes one electron from Na atom. Therefore each Na becomes an Na+ cation and each Cl atom becomes a Cl- anion. ✓ It should also be noted that some atoms can form more than one ion. This usually happens with the transition metals. ✓ For example Fe (iron) can become Fe2+ (called iron (II) or ferrous). Fe can also become Fe3+ (called iron (III) or ferric). 2. Covalent bond: ✓ It is a common type of bonding, in which two or more atoms share valence electrons more or less equally ✓ The simplest and most common type is a single bond in which two atoms share two electrons. Figure (4): Diagram of a covalent bond between hydrogen atoms. 20 Physical Chemistry By: Dr. Heba Salah Mousa ✓ Other types include the double bond, the triple bond. Formation of covalent bond: ✓ Covalent bonds form between two atoms which have incomplete octets that is, their outermost shells have fewer than eight electrons. They can share their electrons in a covalent bond. ✓ The simplest example is water (H2O). Oxygen has six valence electrons (and needs eight) and the hydrogens have one electron each (and need two). The oxygen shares two of its electrons with the hydrogens, and the hydrogens share their electrons with the oxygen. The result is a covalent bond between the oxygen and each hydrogen. The oxygen has a complete octet and the hydrogens have the two electrons they each need. Double and Triple Bonds (covalent bonds) ✓ Covalent bonds can also form between other non-metals, for example chlorine. A chlorine atom has 7 electrons in its valence shell—it needs 8 to complete it. Two chlorine atoms can share 1 electron each to form a single covalent bond. They become a Cl2 molecule. 21 Physical Chemistry By: Dr. Heba Salah Mousa ✓ Oxygen can also form covalent bonds, however, it needs a further 2 electrons to complete its valence shell (it has 6). Two oxygen atoms must share 2 electrons each to complete each other’s shells, making a total of 4 shared electrons. ✓ Because twice as many electrons are shared, this is called a double covalent bond. ✓ Double bonds are much stronger than single bonds, so the bond length is shorter and the bond energy is higher. ✓ Furthermore, nitrogen has 5 valence electrons (it needs a further 3). Two nitrogen atoms can share 3 electrons each to make a N2 molecule joined by a triple covalent bond. ✓ Triple bonds are stronger than double bonds. They have the shortest bond lengths and highest bond energies. Triple bonds more strong than double than single and so triple bonds require more energy than other to be destroyed. 22 Physical Chemistry By: Dr. Heba Salah Mousa 3. Metallic bond: It occurs among metal atoms. Metallic bonding joins a bulk of metal atoms (same or different metals). Formation of metallic bond: ✓ When metallic bonds form, the s and p electrons delocalize. Instead of orbiting their atoms, they form a "sea of electrons" surrounding the positive metal ions. The electrons are free to move throughout the resulting network e.g. Hg2+2 (mercurous ions). ✓ Metallic bonds can occur between different elements. A mixture of two or more metals is called an alloy. The delocalized nature of the electrons explains a number of unique characteristics of metals: 1. Metals are good conductors of electricity 2. Metals have very high melting and boiling points (Metallic bonding is very strong). ✓ Metallic bonds can occur between different elements. A mixture of two or more metals is called an alloy. ✓ Depending on the size of the atoms being mixed, there are two different kinds of alloys that can form: 23 Physical Chemistry By: Dr. Heba Salah Mousa Substitutional alloy Interstitial alloy The resulting mixture will have a combination of the properties of both metals involved. 4. Coordinate bond: ✓ It is a sharing of lone pair of electrons from one atom called donor (Lewis base) to another atom called acceptor (Lewis acid). ✓ Lewis acid: electron pair acceptor e.g. H+, AlCl3, FeBr3, BF3. ✓ Lewis base: electron pair donor e.g. compounds containing heteroatoms (O, S, N) e.g. NH3, H2O 5. Hydrogen bond: ✓ It will occur when a hydrogen atom is attached to an oxygen, nitrogen, or fluorine atom. ✓ This is because there is a large electronegativity difference between hydrogen and fluorine, oxygen, and nitrogen. Thus, molecules such as HF, H2O, NH3 are extremely polar molecules with very strong dipole-dipole forces. 24 Physical Chemistry By: Dr. Heba Salah Mousa ✓ As a result of the high electronegativities of fluorine, oxygen, and nitrogen, these elements will pull the electrons almost completely away from the hydrogen. ✓ The hydrogen becomes a bare proton sticking out from the molecule, and it will be strongly attracted to the negative side of any other polar molecules. ✓ Hydrogen bonding is an extreme type of dipole-dipole bonding. These forces are weaker than intramolecular bonds, but are much stronger than other intermolecular forces, causing these compounds to have high boiling points. 25 Physical Chemistry By: Dr. Heba Salah Mousa Determine which compound can form inter or intramolecular hydrogen bonding 26 Physical Chemistry By: Dr. Heba Salah Mousa IV. Chemical formula 27 Physical Chemistry By: Dr. Heba Salah Mousa For above equation, the following ratios can be observed: 1 mol CH4: 1 mol CO2 1 mol CH4: 2 mol H2O 1 mol CH4: 2 mol O2 2 mol O2: 1 mol CO2 2 mol O2: 2 mol H2O 28 Physical Chemistry By: Dr. Heba Salah Mousa V. Electromagnetic spectrum (EMS) 29 Physical Chemistry By: Dr. Heba Salah Mousa ❑ Properties of light (waves or electromagnetic radiation): 1. Wave motion arises when a periodic disturbance of some kind propagated in elastic medium as air. 2. There are four measurable properties of wave motion such as wavelength (λ), frequency (υ, nu), amplitude and wavenumber (ύ, nu par). Wavenumber is reciprocal of wavelength in cm. (ύ=1/λ) 30 Physical Chemistry By: Dr. Heba Salah Mousa Q) Calculate wavenumber of light with λ=300 nm? Answer ύ=1/λ= 1/300×10-7= 33333 cm-1 Velocity of light (C)= λ×υ= 2.8×108 m/s. λ= C/υ Energy of photons (units of light), E= hυ=hC/λ where h is Planck’s constant. As a result, E is inversely proportional to λ i.e. increase energy of light, decrease of its wavelength. Gamma rays have highest energy and so, lower λ. Radio waves have lowest energy and so, higher λ. Most of EMS are used in the science of spectroscopy (interaction of EMS with matter). Interaction between matter and light in the form of electronic excitation, vibration of bonds and rotation of nuclei. 31 Physical Chemistry By: Dr. Heba Salah Mousa VI. Reactions in aqueous solutions All salts are strong electrolyte such as NaCl, KCl, CH3COONa. HClO4 is perchloric acid, CH3COOH is acetic acid, H3PO4 is phosphoric acid. 32 Physical Chemistry By: Dr. Heba Salah Mousa HCN is hydrocyanic acid, H3BO3 is boric acid, H2CO3 is carbonic acid (very weak acids). Water solvent ability: ❑ In summary, water can dissolve many substances (universal solvent) by two mechanism dipole-dipole interaction and hydrogen bond formation. Figure (5): Dissociation of NaCl in water aided by dipole-dipole interaction.. 33 Physical Chemistry By: Dr. Heba Salah Mousa The concentration: 1. Physical units 2. Chemical units 34 Physical Chemistry By: Dr. Heba Salah Mousa ❖ e.g. the weights required for preparing 1 liter of a molar (1 M) solution of: 1. Sodium hydroxide (NaOH), = 23 + 16 + 1 = 40 g in one liter of solution. 2. Sulphuric acid (H2SO4), = (2x1) + 32 + (4x16) = 98 g in one liter of solution. 3. Copper sulphate pentahydrate (CuSO4.5H2O) = 63.6 + 32 + 64 + (5x18) = 249.6 g in one liter of solution. 35 Physical Chemistry By: Dr. Heba Salah Mousa 36 Physical Chemistry By: Dr. Heba Salah Mousa 1) Calculate the number of moles represented by each of the following: (a) 20 g NaOH (b) 74 g K2CrO4 (c) 148.2 g Ca(OH)2 (d) 50 g CuSO4.5 H2O (e) 57 g Na3PO4. I2 H2O The Answer of 1.(a) Na= 23, O= 16, H=1 Number of moles = weight (g) / molecular weight (M.wt) = 20 / ( 23 + 16 + 1) =20/40 = 0.5 mole Ans. (a) 0.5; (b) 0.38; (c) 2; (d) 0.2; (e) 0.15 2) How many grams of SnCl2.2H2O are needed to prepare 75 ml of 0.25 M solution, Sn= 23, O= 16, H=1, Cl=35.5? The Answer: SnCl2. 2 H2O molecular weight= 118.7 + 2 (35.5) + 2 (2+16) = 225.7 Molarity (M)= weight (g)/M.wt× volume (L) Weight= molarity× M.wt× volume (L)= 0.25× 225.7×75/1000=4.3 g 37 Physical Chemistry By: Dr. Heba Salah Mousa 3) What is the normality of a solution containing 35 g of MnCl2.4 H2O in 300 ml of solution Mn= 55, Cl=35.5, H=1, O=1? The Answer: Molecular weight of MnCl2.4 H2O = 55 + 71 + 72 = 198 Normality= weight/equivalent weight×volume (L) Normality= 35×2/198×0.3= 1.18 N. 4) Ferrous salts react with oxidizing agents according to the equation: Fe2+ → Fe3+ + e How many grams of FeSO4.7H2O are required to prepare 500 ml of a 0.2N solution of the reducing agent Fe= 55.8, O= 16, H=1? The Answer: M.wt of FeSO4.7 H2O = 55.8 + 32 + 64 + 126 = 277.8 Weight= normality× equivalent weight× volume (L) Weight (g)= 0.2×277.8×0.5= 27.78 g. 5) In acid solution the permanganate ion reacts with reducing agents according to the equation MnO4- + 8 H+ + 5 e → Mn2+ + 4 H2O Calculate the weight of KMnO4 necessary to prepare 1 liter of a 1 normal solution. The Answer: Weight= normality× equivalent weight× volume (L)= 1× Mno4-/5×1= Mno4-/5 6) What volume in ml of 2 N barium chloride solution can be prepared from 8 g of BaCl2.2 H2O, Ba= 137.3, Cl=35.5, H=1, O=16 ? 38 Physical Chemistry By: Dr. Heba Salah Mousa Chemical reactions between ions ✓ Combination of ions occurs through the formation of any of the following: 39 Physical Chemistry By: Dr. Heba Salah Mousa 40 Physical Chemistry By: Dr. Heba Salah Mousa 41 Physical Chemistry By: Dr. Heba Salah Mousa Kinetics of chemical reactions ✓ All chemical reactions require time for their completion; some reactions proceed rapidly whereas others proceed slowly. ✓ The rate of chemical change is expressed in terms of the quantities of substances, which are transformed in a given interval of time; this is called the reaction rate. ✓ A study of many reactions reveals that some substances inherently have a greater tendency to react than others (e.g. magnesium reacts more rapidly with dilute acids than iron under the same conditions). ✓ Reactions between ions take place very rapidly, provided no change in valence occurs. For example, the following reactions take place almost instantaneously: H+ + OH- → H2O Ag+ + Cl- → AgCl ↓ ✓ On the other hand, if changes in valence occur, reactions in general proceed more slowly: 2 MnO4- + 5 C2O42- + 16 H+ → 2 Mn2+ + 10 CO2 + 8 H2O ✓ Molecular reactions usually proceed more slowly than ionic reactions: 2 SO2 + O2 → 2 SO3 N2 + 3 H2 → 2 NH3 42 Physical Chemistry By: Dr. Heba Salah Mousa 43 Physical Chemistry By: Dr. Heba Salah Mousa Law of concentration effect ✓ Reaction rates depend on the number of collisions of the reacting particles, and this in turn depends on the number of particles of each of the reactants per unit of volume. ✓ In other words, the reaction rate is related to the gram-ionic or gram- molecular (molar) concentration of each of the reactants. This relationship is expressed as follows: ✓ The rate of a chemical reaction is directly proportional to the molar concentration of each of the reactants raised to a power equal to the number of molecules or ions of that substance appearing in the balanced equation for that reaction. ✓ This principle was known originally as the law of mass action, but it is now more properly called the law of concentration effect. According to the law of concentration effect, in a reaction of the type: A→X+... Rate α [A] [A], is the concentration of that substance in moles per litre. If one quantity is proportional to another, the first quantity is equal to the second multiplied by a constant, K, or: Rate = K  [A] In a reaction of the type: A + B → X +... The reaction rate is proportional to the molar concentration of substance A (law of concentration effect), or Rate α [A]. The rate is also proportional to the molar concentration of substance B, or Rate α [B]. If one quantity is proportional to each of two other quantities, it is proportional also to their product, or Rate α [A]  [B]: Rate = K  [A]  [B] 44 Physical Chemistry By: Dr. Heba Salah Mousa Again, if the reaction is of the type: A + 2 B → X +... In this reaction, two particles of one kind must collide with one particle of another kind to produce a reaction, the equation for the reaction rate is: Rate = K  [A]  [B]2 The concentration term for substance B is squared in this equation, since the reaction involves collisions of two similar molecules. In general: aA + bB + cC +... → X +... Rate = K  [A]a  [B]b  [C]c ... Reversible reactions 45 Physical Chemistry By: Dr. Heba Salah Mousa Equilibrium constant (K) ✓ All reversible reactions are, in fact, two reactions proceeding in opposite directions. Let us consider a reversible reaction of the type: A+B  C+D ✓ If pure substance A and substance B are mixed, the reaction at the instant of mixing proceeds only to the right, since no products C and D are initially present. The rate of the forward reaction (to the right in the above equation) is proportional to the product of the molar concentrations of reactants A and B; the reaction rate is expressed as: Rate(F) = KF  [A]  [B] As A and B react, C and D are formed; they in turn react with one another to re- form A and B. The rate of the latter reaction is proportional to the product of the molar concentrations of C and D; the reaction rate is expressed as: Rate(l) = Kl  [C]  [D] At equilibrium there is no further change in the concentrations of the reactants, since for every particle of A and of B that react to form C and D, another particle of each is formed by the interaction of C and D. At equilibrium, KR and KL (in the above equations) are equal, and consequently: Kr  [A]  [B] = Kl  [C]  [D] 46 Physical Chemistry By: Dr. Heba Salah Mousa Characteristics of equilibrium constant (K): 1. At constant temperature, it has definite numerical value, which is independent of the original concentration of the reactants. 2. It has different values for different reactions. 3. For given reaction, it has different values for different temperatures. 4. The same state of equilibrium is attained, and consequently the value for the equilibrium constant is the same, regardless of which set of reactants is used. Thus, in the reaction: A + B→C + D The same state of equilibrium will be attained whether A and B or C and D are initially mixed. 5. The greater the numerical value of the equilibrium constant the more complete the reaction. 6. The value of the equilibrium constant is not changed by the presence of a catalyst. 47 Physical Chemistry By: Dr. Heba Salah Mousa Calculation of the equilibrium constant ❑ Assume that 1 mole of SO2 and 1 mole of NO2 are introduced into a vessel having a capacity of 1 litre, and the mixture is then allowed to stand until equilibrium is established. SO2 + NO2  SO3 + NO ❑ By experiment the concentration of SO3 is then found to be 0.6 mole per litre. ❑ From the equation, it is noticed that 1 mole of SO2 reacts with 1 mole of NO2 to form 1 mole of SO3 and 1 mole of NO. ❑ Therefore, if 0.6 mole of SO3 is formed, there will also be 0.6 mole of NO. Further, 0.6 mole each of SO2 and NO2 are required to form SO3 and NO and consequently (1 - 0.6) or 0.4 mole of SO2 and also of NO2 remain in the equilibrium mixture ❑ It should be noted that the reactions do not stop at equilibrium, but since forward and reverse reaction rates are equal at equilibrium there is no net change in the concentrations. For the above reaction we write the equation for the equilibrium constant as follows 48 Physical Chemistry By: Dr. Heba Salah Mousa Effect of changing concentration on the reaction ❑ After equilibrium is established, no further change in the concentrations of any of the reactants in the system occurs. ❑ Now let us consider the effect of changing the concentration of one of the reactants in the system, which has come to equilibrium: A + B  C + D ❑ If more A is added, the forward reaction rate is accordingly increased, since the forward reaction rate is proportional to the molar concentration of A. The reverse reaction, however, is not immediately influenced by a change in the concentration of A. ❑ Therefore, C and D are produced for a time at a rate faster than that at which they combine. ❑ As more of C and D are formed, however, the rate at which they react becomes greater, and eventually the two reaction rates again become equal. ❑ When equilibrium is once again established, the concentrations of the reactants will not be the same as in the original mixture, nor will the two reaction rates be the same as before the original equilibrium was disturbed. 49 Physical Chemistry By: Dr. Heba Salah Mousa Le-Chatelier principal 50 Physical Chemistry By: Dr. Heba Salah Mousa No shift occurs with the following system reaction since no change in the total volume accompanies the reaction. H2 (1 volume) + I2 (1 volume)  2 HI (2 volumes) A system in equilibrium is not changed by the addition of a catalyst since under equilibrium conditions a catalyst has the same effect on both forward and reverse reaction. 51