Hydronium ion: The H3O+ ion, formed when an acid reacts with water.

An acid is a substance that produces hydrogen ions, H+, when dissolved in water.

A base is a substance that produces hydroxide ions, OH-, when dissolved in water.

The neutralization reaction of an acid with a base yields water plus a salt, an ionic compound composed of the cation from the base and the anion from the acid.

The Arrhenius definition of acids and bases applies only to processes that take place in an aqueous solution.

A Bronsted–Lowry acid is any substance that is able to give a hydrogen ion, H+, to another molecule or ion. A hydrogen atom consists of a proton and an electron, so a hydrogen ion, H+, is simply a proton.

Acids with one proton to donate, such as HCl or HNO3, are called monoprotic acids; H2SO4 is a diprotic acid because it has two protons to donate, and H3PO4 is a triprotic acid because it has three protons to donate. Notice that the acidic H atoms (i.e., the H atoms that are donated as protons) are bond.

Bronsted–Lowry base is any substance that can accept H+ ions from an acid.

Putting the acid and base definitions together, an acid-base reaction is one in which a proton is transferred.

Conjugate acid-base pair : Two substances whose formulas differ by only a hydrogen ion, H+

Conjugate base: The substance formed by loss of H+ from an acid.

Conjugate acid: The substance formed by addition of H+ to a base.

Strong acid is an acid that gives up H+ easily and completely dissociates in water.

Dissociation is the splitting apart of an acid in water to give H+ and an anion.

Weak acid is an acid that gives up H+ with difficulty and does not completely dissociate in water.

The stronger the acid, the weaker its conjugate base; the weaker the acid, the stronger its conjugate base.

An acid-base proton transfer equilibrium always favours reaction of the stronger acid with the stronger base and formation of the weaker acid and base. That is, the proton always leaves the stronger acid (whose weaker conjugate base cannot hold the proton) and always ends up in the weaker acid (whose stronger conjugate base holds the proton tightly).

Acid dissociation constant( Ka): The equilibrium constant for the dissociation of an acid (HA), equal to [H+][A-] / [HA].

In the Bronsted–Lowry sense, water can act as both an acid and a base.

• When in contact with a base, water reacts as a Bronsted–Lowry acid and donates a proton to the base.

• When in contact with an acid, water reacts as a Bronsted–Lowry base and accepts H+ from the acid.

• Substances like water, which can react as either an acid or a base depending on the circumstances, are said to be amphoteric in nature.

As a pure substance the concentration of water is essentially constant. We can therefore put the water concentrations [H2O] together to make a new equilibrium constant called the ion-product constant for water (Kw) , which is simply the H3O+ concentration times the OH- concentration. At 25 °C (298 K), Kw = 1.00 * 10^-14

The pH of a solution is the negative common logarithm of the H3O+ concentration.

Acid-base indicator is a dye that changes colour depending on the pH of a solution.

The well-known dye litmus is red below pH 4.8 but blue above pH 7.8 and the indicator phenolphthalein (fee-nol-THAY-lean) is colourless below pH 8.2 but red above pH 10. To make pH determination particularly easy, test kits are available that contain a mixture of indicators known as universal indicator to give approximate pH measurements in the range 2–10.

1 equivalent of acid contains 1 mol of H+ ions, and 1 equivalent of base contains 1 mol of OH- ions.

The normality (N) of an acid or base solution is defined as the number of equivalents (or milliequivalents) of acid or base per litre of solution.

• Normality of acid = (Molarity of acid2)* (Number of H+ ions produced per formula unit)

• Normality of base = (Molarity of base) * (Number of OH- ions produced per formula unit)

Salt of Weak Acid + Strong Base → Basic Solution.

Salt of Strong Acid + Strong Base → Neutral Solution.

Salt of Weak Acid + Weak Base :

Both cation and anion in this type of salt react with water, so we cannot predict whether the resulting solution will be acidic or basic without quantitative information.

Buffers are combinations of substances that act together to prevent a drastic change in pH.

When 0.010 mol of strong acid 1H3O+2 or 0.00 mol of strong base (OH-) are added to 1.0 L of pure water with an initial pH of 7.00, the pH of the solution varies between 12.00 (basic) and 2.000 (acidic). When the same amounts of strong acid or base is added to a 0.10 M acetic acid buffer solution having an initial pH of 4.74, the pH of the solution varies only between 4.85 and 4.68.

In general, the most effective buffers meet the following conditions:

• The pKa for the weak acid should be close to the desired pH of the buffer solution.

• The ratio of 3HA4 to 3A-4 should be close to 1, so that neither additional acid nor additional base changes the pH of the solution dramatically.

• The molar amounts of HA and A- in the buffer should be approximately 10 times greater than the molar amounts of either acid or base you expect to add so that the ratio [A-] /[HA] does not undergo a large change.

Titration is a procedure for determining the total acid or base concentration of a solution.