BASIC CONCEPTS FROM GENERAL CHEMISTRY
Atomic Weight – refer to the relative weights of the atoms as compared with carbon. Carbon has
atomic weight standard of 12.
Gram atomic weight – refers to a quantity of the element in grams corresponding to the atomic
Gram molecular weight – GMW or MW refers to the molecular weight in grams of any particular
compound (referred also as mole).
Molar solution – consists of 1-gram molecular weight dissolved in enough water to make 1 liter
of solution. Molar concentrations are used for equilibrium calculations.
Molal solution – consists of 1-gram molecular weight dissolved in 1 liter of water, resulting
solutions having a volume slightly in excess of 1 liter. This is used when physical properties of
solutions such as vapor pressure, freezing point, boiling point, are involved.
Equivalent weight: EW = ----------
Where Z = absolute value of the ion charge
= the number of H+
ions a species can react with or yield in an acid-base
= the absolute value of the change in valence occurring in an oxidation-
1. Determine equivalent weight of Ca2+
a. EW = MW / Z = 40 g per mole/ 2 = 20 g per equivalent
b. MW of CaCO3 = 40 x 1 + 12 x 1 + 16 x 3 = 100 g per mole
Z = 2 E = MW / Z = 100 g per mole / 2 = 50 g per equivalent
2. What concentration is 40 mg/l of Ca2+
when expressed in CaCO3?
One equivalent of an ion or molecule is “chemically” equivalent to one equivalent of a
different ion or molecule. Thus is concentrations are expressed in terms of equivalents per
little (eq/l) they can be added, subtracted or converted easily.
= 40 mg/l x 1eq/20 g x 1 g/1000mg = 0.002 eq./l of Ca2+
- mg/L as CaCO3 = 0.002 eq/l x 50 g/eq x 1000mg/g = 100 mg/l as CaCO3
Ionic bond – formed by the transfer of electrons from one atom to another
Covalent bond – when electrons are not transferred by are shared between atoms e..gCl2, N2, O2
Valency or oxidation number – determined by the number of electrons that it can take on, give up
or share with other atoms.
If electrons are lost, the atom becomes positively charged ion, and if electrons are gained, the
atom becomes negatively charged ion
- Metal or metal-like element loses electrons to gain or approach a stable condition
with no electrons on its outer ring.
- The nonmetal steals electrons from the metal to complete its outer ring to eight
electrons, a stable configuration.
Chemical reactions become equations only when they are balanced
Mass must be conserved – the total number of each kind of atom must be the same on both
sides of the equation
The sum of the charge on one side of the equation must be equal that on the other
Example: NaOH + HCl -----> NaCl + H2O
(23 + 16 + 1) (1 + 35.5) (23 + 35.5) (2x1 + 16)
40 36.5 58.5 18
40 g of NaOH combined with 36.5 g HCl will yield 58.5 g NaCl and 18 g H2O
Example 2: 2 HCl + Na2CO3 -----------> H2O + 2NaCl + CO2
73 106 18 117 44
1. Boyle’s Law: The volume of a gas varies inversely with its pressure at a constant
2. Charles Law: The volume of gas at constant pressure varies in direct proportion to the
3. Generalized Gas Law: For a given quantity of gas, Boyle’s law and Charles’ law can be
P V = n R T
Where n = to the number of moles of gas in a particular sample
R = universal constant for all gases
= depends on the units chosen for the measurement of P, V and T
= 1 mole of gas at 1 atm pressure occupies a volume of 22. 414 liters at 273 K.
= 0.082 liter atmosphere per mole per Kelvin (0.082 l-atm/mol-K)
Example: What tank volume is required to hold 10,000 kg of methane gas
(CH4) at 25o
C and 2 atm pressure?
Molecular weight of CH4 gas is 12 + 4x1 = 16 g
Number of moles in 10,000 kg is 10,000,000 / 16 g = 625,000 mol
General gas law: V = nRT/P = 625,000 (0.082)(273+ 25)/ 2 = 7.64 x 106
Volume of tank = 7.64 x 106
4. Dalton’s law of Partial Pressure: In a mixture of gas, such as air, each gas exerts pressure
independently of the others. The partial pressure of each gas is proportional to the amount
(percent by volume) of the gas in the mixture, or in other words, it is equal to the pressure
that gas would exert if it were the sole occupant of the volume available to the mixture.
5. Henry’s law: The weight of any gas that will dissolve in a given volume of a liquid, at
constant temperature, is directly proportional to the pressure that the gas exerts above the
C equi = KH Pgas
Where C equi = concentration of gas dissolved in the liquid at equilibrium
Pgas = is the partial pressure of the gas above the liquid
KH = is the Henry’s law constant for the gas at the given temperature.
KH for oxygen in water at 20o
C is 43.8 mg/1-atm. Since air contains 21 percent by volume of
oxygen, the partial pressure in air according to Dalton’s law would be 0.21 atm when the total
air pressure is 1 atm. Therefore the equilibrium concentration of oxygen in water at 20o
in the presence of 1 atm of air would be 43.8 x 0.21 = 9.2 mg/l
Application of Henry’s Law:
1. Aeration of water supply for the removal of gases
2. Aeration of domestic sewage for the application of oxygen
3. Removal of gases from industrial wastewater
6. Graham’s Law – the rates of diffusion of gases are inversely proportional to the square root
of their density.
Example: Using atomic weights of hydrogen, oxygen, chlorine, bromine: 1, 16, 36, 80
respectively. Oxygen diffuses about one-fourth, chlorine about 1/6, and bromine about 1/9 as
fast as hydrogen.
7. Gay-Lussac’s Law of Combining Volume
The volumes of all gases that react and that are produced during the course of reaction are
related numerically to one another as a group of small, whole numbers.
C + CO2 -----------> CO2
Solid 1 vol 1 vol
CH4 + 2O2+ ----------> CO2 + 2 H2O
1 vol 2 vol 1 vol 2 vol (>100o
0 vol (< 100o
Vapor Pressure – the presence of a non-volatile solute in a liquid always lowers the vapor
pressure of the solution.
Example: When sugar, sodium chloride, or similar substance is dissolved in water, the vapor
pressure is decreased.
Raoult’s Law: The extent of the physical blocking effect or depression of the vapor pressure is
directly proportional to the concentration of the particles in solution. Raoult’s law is applicable
only to dilute solutions.
For solutes that do not ionize – the effect is proportional to the molal concentration
For solutes that do ionize – the effect is proportional to the molal concentration times the
number of ions formed per molecule of solute modified by the degree of ionization.
Chemical kinetics is concerned with the speed or velocity of reactions.
If a reaction is considered in which A, B, and C are possible reactants, then the rate of equations
that express the concentration dependence of the reaction may take one of the following forms:
Rate = kCao
= k 0 order
Rate = kCa 1st
Rate = kCaCb = kCa
Rate = kCaCbCc = kCa
Cb = kCa
Where Ca, Cb, C c represent the concentrations of reactants A, B and C
The order of reaction is the sum of the exponents of the concentration terms in the reaction
Applications of chemical kinetics:
1. Biotransformation reactions
2. Microbial growth and decay
4. Radioactive decay
6. Chemical hydrolysis
7. Oxidation – reduction
The rate of reaction is independent of concentration.
Plot of concentration vs. time is a straight line
C = Co + kt
Where C = concentration of reactant at time t
Co = initial concentration of the reactant
1. Biologically induced reactions particularly growth on simple, soluble substrates
2. Oxidation of ammonia to nitrite
3. Oxidation of glucose by aerobic bacteria
The rate of reaction is directly proportional to the amount undecayed material
C = Co e –kt
or C = Co 10 –k’t
where k’ = k / 2.303
Application to radioactive substances:
Decomposition is expressed as half-life
Half-life is the time required for the amount of substance to decrease to half its initial
value ( C = ½ Co) at time t1/2
. t1/2 = 0.693 / k