Chemistry is traditionally divided in compounds found in plants and animals, known as Organic Chemistry, and those found in non-life matter, known as Inorganic Chemistry. Then there is Physical Chemistry that deals with the physical interactions of atoms and molecules in both Organic and Inorganic Chemistry.
In chemistry our aim is to understand how atoms and molecules interact with each other. So we generally introduce the study of Chemistry with lessons on atoms and molecules. Some would consider this as Physical Chemistry while others would name it General Chemistry. So here we are continuing from where we left off in our introduction.
LESSON 1 : INTERMOLECULAR FORCES
There are basically three physical states for a covalent compound, namely; solid, liquid and gas. Ionic compounds generally exist as crystals. The forces in ionic compounds are the columbic forces, while for the covalent compounds it will be the London forces, the Van der Waals forces and hydrogen bond.
Ionic compounds are solids at standard atmospheric pressure and temperature (SATP), and for many they are solids at all temperatures. No one would think of melting a rock. For covalent compounds the intermolecular forces would determine whether the compound is to be a gas, a liquid or a solid at SATP.
Molecules have internal energies. The bonds are in various modes of motion and the molecules tend to roam about. The latter translation motion will depend on the size of the molecule and intermolecular attraction that exist. When the intermolecular interaction is greater than the energetic of the molecule the compound is a solid. At a particular temperature when the molecules break free from being "held tightly" to each other the compound melts to a liquid. If you continue putting energy into the system the energy of the molecules increase further the compound finally vapourises.
LESSON 2 : ENERGY of MOLECULES
At any temperature the molecules in the compound possess different amount of energies. There is a dynamic equilibrium between the molecules to ensure that the sum total energy for the sample is maintained. The Boltzmann's Distribution Function is a mathematical expression for the distribution of the molecules according to their energies, in a compound. The Maxwell's Distribution Function is a particular aspect of the Boltzmann's Distribution Function for the distribution of the molecules according to their translation energies (or speed), in the system.
LESSON 3 : GASEOUS STATE
Of all the physical states, the gaseous state is the easiest to study, as there are minimum intermolecular interactions. The parameters are limited only to temperature, pressure and the number of moles of the molecules. We can even assume that there is no intermolecular interaction between the molecules and referred to the gas as an ideal gas. The equation governing their behaviour is; pV = RnT, where R is the constant for 1 mole of gas in 22.4 dm� volume, at 760 torr of pressure and a temperature of 273.15 K, (that is at STP).
For a mixture of different molecules Dalton's Law of Partial Pressures states that the pressure measured is the sum of the pressures exerted by the individual gases if they have occupied the same volume individually. Again assuming that they are ideal gases.
Gases in the atmosphere qualify as ideal gases, however industrial gases are normally compressed under high pressure and so would not be ideal gases. They could not be defined by the ideal gas expression. As a matter of fact it is difficult to derive an equation to describe their behaviour. The best we have so far is the van der Waals Equation of; p = nRT / (V-nb) - a(n/V)� . "a and b" are experimentally determined for each and every gas.
LESSON 4 : KINETIC THEORY of GASES
Pressure is the force of bombardment of the gas molecules on the wall of the container. If the energy (or speed) is high and the number of molecules in the container is high, the impact on the wall will be more forceful and more often. This would result in a higher pressure on the container. The frequency of bombardment is of course dependent on the volume. If I have a container equipped with a piston, then the piston will move if the bombardment on both sides of the container is not equal.
According to Newton's Second Law of motion, pressure is the rate of change of momentum. Using this classical physics it is able to explain the ideal gas equation, and the transport properties - effusion, diffusion, thermal conductivity and viscosity - of ideal gases from a molecular level.
LESSON 5 : CHEMICAL KINETICS
Unlike gases, liquids are not free to move about freely, even though they are not held on to each other as rigidly as in the solid state. So their translational motion could not be as easily studied as in the gas state. The best science can do is to use the mathematics of chances known as the Monte Carlo probability, and the correlation between the mathematics and the practical results are not as simple as for the gas state. So let us move on to other topics.
Chemical reactions are better conducted in liquid state as they satisfied two important criteria. (1) It must be able to move about; and (2) the frequency of collision must be very high. The first criterion is not present in the solid state and the second critirion is not found in a gas state unless it is at very high pressure.
The first step to study a reaction is to conduct the reaction and determine all the products obtained - the compounds and their amounts. The next step is to understand how the reagents react to give the products. This is known as the reaction mechanism. It is made up of a proposed sequence of elementary reactions. A mathematical expression for the rate of the reaction can then be derived from the mechanism. This is known as the rate law. Chemist would then determine the rate of the reaction. When the derived expression agrees with the proposed rate law the chemist can then conclude that the reaction mechanism proposed is a possibility (not a certainty). Two different proposed mechanisms can give the same rate law.
The average number of collisions which took place and the minimum amount of impact energy needed before the species react are computed from the constant in the rate law by the use of the Arrhenius equation.
LESSON 6 : THERMODYNAMICS
We now move on to study the energetic of molecules with the help of thermodynamics. In chemical reaction we study the progress of the reaction. In thermodynamics we study the reaction when it has reached an equilibrium.
When the energy is transferred between a system and the surrounding because of a temperature difference, we say the energy is transformed in the form of heat. Thermodynamics studies the transformation of heat energy to work energy. The total amount of heat in a system is known as the internal energy (U) of the system. This would include the sum total of energies of all the particles (atoms, molecules, their interactions, their motions, etc). We are not able to measure U since it is so complex, so we have to make do with the study of the change of the internal energy of a system with respect to adding heat and work to a system. The correlation is known as the First Law of Thermodynamics. This law can be stated in different forms but the most used form is the mathematical representation - ΔU = q + w ; where q and w is the heat energy added to and work done on the system respectively.
LESSON 7 : WORK AND HEAT
U is a state function but work energy (w) and heat energy (q) are not state functions. We define another term H, known as enthalpy, as; H = U + pV. U, p and V are all state functions, so H must also be a state function. At constant pressure ∆U = ∆H - p∆V. This is the First Law of Thermodynamics with q = ∆H. So by expressing enthalpy the way we did, we have found a condition for q to be a state function. The condition that pressure be constant is not restrictive since a large portion of chemistry done in the laboratory is at one pressure; atmospheric pressure.
One parameter to measure how heat energy affect the internal energy of matter is the molar heat capacity. There are two molar heat capacities; one at constant pressure Cp = {∂H / ∂T}p,m and the other at constant volume Cv = {∂U / ∂T}v,m. The variation of the molar heat capacity with temperature is not great, except when you are dealing with very, very high temperatures, so for convenience we generally assume that they are constants for a particular matter.
There is no problem determining Cp,m in a laboratory, but Cv,m is a problem as it involves change in internal energy. However the relationship of the two is given by Cp - Cv = (α� TV) / kT, α is the expansion coefficient and kT the isothermal compressibility of the matter.
LESSON 8 : CHANGE in ENTHALPY
It is not possible to determine the absolute enthalpy value for any system, so we focus on the change in enthalpy (ΔH). ΔH of one mole of matter at 760 torr pressure is referred to as the standard enthalpy, ΔHΘ. We can define standard enthalpy for the different state transitions (fusion, vapourisation, etc) and chemical reactions. The Born-Haber cycle, states that the enthalpy of a physical process is the sum enthalpies of its composite processes, while the Hess's Law states that the enthalpy of a reaction is the sum enthalpies of its composite thermochemical reactions.
We find it useful to compile the standard enthalpy of formation of compounds. We can then compute the enthalpy for any reaction from the enthalpy of formation of the compounds involved in the reaction according to Hess's Law.
Kirchhoff's Law; ΔH(T) = ΔHΘ + ΔCp is used to compute the enthalpy at any temperature when the enthalpy at a certain temperature is known, from the Cp for the compounds involved.
Since the standard enthalpy of formation measures the change in bond energies we do the ultimate and compile data for the energy for each and every covalent bond, known as the enthalpy of bond dissociation. So we can now determine the energy of chemical bonds in molecules.
