1. Acids - Bases

An Acid is a substance that produces H3O+ (H+) when it is dissolved in water. It is a proton donor and an electron pair acceptor or a species that donates protons. For example: HCl, NH4, AlCl3.
A Base is a substance that produces an OH- when it is dissolved in water (Arrhenius). A proton acceptor (Brønsted), or a electron donor. For example: NaOH, KOH, CH3NH2.
Acids and bases were first identified as specific types of compounds because of their behavior in aqueous solutions.
Acids and bases relate to each other in Conjugate Pairs, somewhat like husbands and wives. For every acid there is a conjugate base; and for every base there is a conjugate acid. Just like every husband has a wife and vice versa. The two members of the conjugate pair are related by the donating and accepting of a single proton.
The equation below, Equation (1), demonstrates a power struggle going on between the two couples and within them. There is a competition for which base, H2O, (keep in mind that H2O can act as an acid or a base because it auto-ionizes itself, meaning it gives protons back and forth within itself, thus acting as both an acid and a base;) See Equation (2). Then A- will get the proton. The winner is the stronger base which has a greater affinity for H+ and everything will go its way. This base will determine whether the equation goes to the right or the left at equilibrium
HA(aq) + H2O(l) <==> H3O+(aq) + A-(aq)



To determine the strength of an acid or base can be difficult within conjugate pairs. Strong acids have weak conjugate bases so the equilibrium lies far to the right. A weak base has a lower affinity for protons that water. So water wins the H+ ion as in the reaction in Equation (1) above.
Of course, a weak acid has a strong conjugate base so the equilibrium will go to the left, and the acid will not dissociate that much.
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Acids

In order to determine how much the H+ will get, or the degree to which a weak monoprotic acid will dissociate, we use:
Ka, the acid dissociation constant: Ka = [H3O+][A-] / [HA]
The weaker the acid, the smaller its Ka and the less it will dissociate. Strong acids completely dissociate into their component ions in aqueous solution.
Note that [H2O is omitted from the Ka expression because the concentration of H2O is so high in an aqueous solution and changes so little, it is basically treated as a constant.
Below are examples of strong and weak acids:



Strong acids: HCl
H2SO4
HNO3
HClO4

Weak acids: NH4+
NCN
HF
HNO2


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Bases
We can also take a look at this from the base's point of view using the base dissociation constant, Kb. KB is a measure of the degree to which a base will dissociate:
B(aq) + H2O(l) <==> BH+ + OH-(aq)
The weaker the base, the smaller its Kb.
Kb always refers to the reaction of a base with water to from the conjugate acid and the hydroxide ion.
A compound that increases the concentration of hydronium ions (H3O+) in aqueous solution is an acid. A hydronium ion is a hydrated hydrogen ion, but can be written in a couple of ways; H+, H+(aq). Don't forget that even though it is written in both of those ways, a hydrogen ion is always associated with at least one water molecule in an aqueous solution. Hydrogen chloride, HCl, and sulfuric acid, H2SO4, are both acids; since they have hydrogen that can be released as protons, they are called protonic acids. The protonic acids help increase the concentration of the hydronium ion of water, by the following reactions when they really dissolve:


EQUATION (1)


HCl(g) + H2O(l) ==> H3O+(aq) + Cl-(aq)

EQUATION (2)


H2SO4(l) + H2O(l) ==> H3O+(aq) + HSO4-(aq)

Don't forget that some compounds don't contain any hydrogen but increase the hydronium ion concentration in solution by reacting with water and giving protonic acids as the result. The protonic acids that are given off start to act like electrolytes giving hydrogen ions to the solution; an example would be Sulfur trioxide. One molecule of SO3 reacts with one molecule of water to give a mole of sulfuric acid, H2SO4:

EQUATION (3)
SO3(g) + H2O(l) ==> H2SO4(l)


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If there is an excess of water present, the hydronium ion concentration will be increased by the reaction of H2SO4 with water, just as indicated in Equation (2) above. The net equation below, Equation (4), can be written as the sum of Equation (3) plus Equation (2):

EQUATION (4)
SO3(g) + 2H2O(l) ==> H3O+(aq) + HSO4-(aq)

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When a covalent molecule separates into ions it is said to ionize. Hydrogen chloride dissolves in water and ionizes in water; then ionizes into hydronium ions and chloride ions; nitric acid ionizes into hydronium ions and nitrate ions. A few protonic acids, such as hydrogen chloride, sulfuric acid, nitric acid, and perchloric acid, ionize completely. The preceding acids are called strong acids. Most acids however are weak acids. Only a small amount of the molecules ionize when dissolved in some water. Some examples of weak acids include hydrogen fluoride (HF), acetic acid (CH3CO2H), boric acid (H3BO3), and hydrogen cyanide (HCN). Although the hydronium ion concentration in an aqueous solution of a weak acid is greater that that of pure water, the hydronium ion concentration is a smaller amount than, if the weak acid were to be completely ionized.

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A compound that increases the concentration of hydroxide ion (OH-) in a solution is a base. Sodium Hydroxide, NaOH, and calcium hydroxide, Ca(OH)2, are examples of bases; since they and hydroxide ions, they are also called hydroxide bases. They also increase the hydroxide ion concentration by the following reactions when they dissolve in water:
EQUATION (5)
H2O (l)

NaOH(s) ==> Na+(aq) + OH-
EQUATION (6)


H2O(l)
Ca(OH)2(aq) ==> Ca2+(aq) + 2OH-(aq)

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Many compounds that contain a metal and oxygen react with water to form hydroxides. For example, one mole of sodium oxide, Na2O, reacts with one mole of water to give two moles of sodium hydroxide:

EQUATION (7)
Na2O(s) + H2O(l) ==> 2NaOH(s)

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If more that one mole of water is present, the excess basically acts as a solvent; and the sodium hydroxide dissolves with the formation of ions as described by Equation (8) below:

EQUATION (8)
Na2O(s) + H2O(l) ==> 2Na+(aq) + 2OH-(aq)

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Ammonia, NH3, is a base because it reacts with water to a limited extent to form ammonium ions, NH4+, and hydroxide ions:
EQUATION (9)
NH3(g) + H2O(l) <==> NH4+(aq) + OH-(aq)
Many hydroxide bases are similar to salts, as these bases are ionic compounds containing a cation and the hydroxide anion, OH-. These bases ionize completely in water and give solution that have the metal cations and hydroxide anions; the preceding are called strong bases. The hydroxides of the alkali metals (Group IA) and of calcium, strontium, barium, and radium (Group IIA, alkaline earth metals) are examples of strong bases. Weak bases are bases, such as beryllium hydroxide, Be(OH)2, which ionize only slightly in water, or bases such as ammonia, which also react with water to some extent. Don't forget that solutions of weak bases contain only a small amount of hydroxide ion and a large amount of undissociated or unreacted molecules of the base. Other weak bases include aluminum hydroxide, Al(OH)3, and pyriding, C5H5N.
During the formation of coordinate covalent bonds, one of the two atoms that are involved in forming the bond provides both electrons of the electron pair that bonds the atoms together. Both of these atoms, respectively, a base and an acid, according to the Lewis Theory of acids and bases. A Lewis Base is an ion or a molecule that helps in providing an electron pair. A Lewis Acid however, is a molecule or ion that accepts the pair of electrons to form the bond.

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