Comprehensive Guide to Acids, Bases, Salts, and pH Concepts in Chemistry

Arrhenius acid

Hydrogen-containing compound that produces H+ ions in solution

Arrhenius base

Hydroxide-containing compound that produces OH- ions in solution

Ionization

The process in which individual positive and negative ions are produced from a molecular compound that is dissolved in solution

Dissociation

The process in which individual positive and negative ions are released from an ionic compound that is dissolved in solution

Bronsted-Lowry acid

Substance that can donate a proton (H+ ion) to some other substance

Brønsted-Lowry base

Substance that can accept a proton (H+ ion) from some other substance

Acid

A substance that donates a hydrogen ion to another substance.

Base

A substance that accepts a hydroxide ion from another substance.

Acid in Water

A substance that produces hydrogen ions in water.

Base in Water

A substance that produces hydroxide ions in water.

Conjugate Acid-Base Pair

Two members that differ from one another by one H+.

Amphiprotic Substance

A substance that can either lose or accept a proton and thus can function as either a Brønsted-Lowry acid or a Brønsted-Lowry base.

Example of Amphiprotic Substance

H2O, H3O+, OH-.

Monoprotic Acid

An acid that supplies one proton (H+ ion) per molecule during an acid-base reaction.

Examples of Monoprotic Acid

HCl, HNO3, HBr.

Diprotic Acid

An acid that supplies two protons (H+ ions) per molecule during an acid-base reaction.

Examples of Diprotic Acid

H2CO3, H2SO4.

Triprotic Acid

An acid that supplies three protons (H+ ions) per molecule during an acid-base reaction.

Example of Triprotic Acid

H3PO4.

Polyprotic Acid

An acid that supplies two or more protons (H+ ions) during an acid-base reaction.

Difference between Monoprotic, Diprotic, and Triprotic Acid

A monoprotic acid supplies one proton per molecule; a diprotic acid supplies two protons per molecule; and a triprotic acid supplies three protons per molecule.

Acidic Hydrogen Atoms

Acidic hydrogen atoms are routinely listed at the beginning of the chemical formula.

Example of Acidic Hydrogen Atom

Acetic acid's formula is HC2H3O2.

Strong Acid

Transfers ~100% of its protons to water in an aqueous solution.

Weak Acid

Transfers only a small percent of its protons to water in an aqueous solution.

Equilibrium position of Strong Acid

Lies far to the right.

Equilibrium position of Weak Acid

Lies far to the left.

Strong Acids

Strong electrolytes.

Weak Acids

Weak electrolytes.

Strong Bases

Hydroxides of Groups IA and IIA.

Acid Ionization Constant (Ka)

An equilibrium constant for the reaction of a weak acid with water.

Acid Strength

Increases along with an increase in percent ionization.

Percent Ionization

Increases with an increase in the magnitude of Ka.

Base Ionization Constant (Kb)

The equilibrium constant for the reaction of a weak base with water.

Ionic Compounds

Contain a metal or polyatomic ion as the positive ion and a nonmetal or polyatomic ion (except hydroxide) as the negative ion.

Soluble Salts

All common soluble salts are completely dissociated into ions in solution.

Definition of a Salt

Ionic compound containing a metal or polyatomic ion as the positive ion and a nonmetal or polyatomic ion (except hydroxide) as the negative ion.

Neutralization Reaction

The chemical reaction between an acid and a hydroxide base in which a salt and water are the products.

Self-Ionization

An extremely small percentage of water molecules in pure water interact with one another to form ions.

Ion Product Constant for Water

At 24°C: Kw = [H3O+][OH-] = 1.00 × 10-14.

Three Possible Situations

[H3O+] = [OH-]; neutral solution, [H3O+] > [OH-]; acidic solution, [H3O+] < [OH-]; basic solution.

Acidic Solution Definition

An aqueous solution in which the concentration of H3O+ ion is higher than that of OH- ion.

pH Formula

pH = -log[H3O+].

pH Range

pH range between 0 and 14 in aqueous solutions at 24°C.

Net Effect of Self-Ionization

The formation of equal amounts of hydronium and hydroxide ions.

Acidic Solution Characteristics

An aqueous solution in which the concentration of H3O+ ion is higher than that of OH- ion.

Basic Solution Characteristics

An aqueous solution in which the concentration of OH- ion is higher than that of H3O+ ion.

Neutral Solution Characteristics

An aqueous solution in which the concentration of H3O+ ion is equal to that of OH- ion.

pH

A measure of the acidity or basicity of a solution, calculated as pH = -log [H3O+].

Neutral pH

A pH of 7.00, indicating a neutral solution.

Basic pH

A pH greater than 7.00, indicating a basic solution.

Acidic pH

A pH less than 7.00, indicating an acidic solution.

pKa

A method of expressing the strength of acids, calculated as pKa = -log Ka.

Ka

The acid dissociation constant, used to measure the strength of an acid.

Hydrolysis

The reaction of a salt with water to produce hydronium ions and/or hydroxide ions.

Salt of strong acid and strong base

Does not hydrolyze, resulting in a neutral solution (e.g., NaCl, KBr).

Salt of strong acid and weak base

Hydrolyzes to produce an acidic solution (e.g., NH4Cl).

Salt of weak acid and strong base

Hydrolyzes to produce a basic solution (e.g., NaC2H3O2).

Salt of weak acid and weak base

Hydrolyzes to produce a slightly acidic, neutral, or slightly basic solution, depending on the relative strengths.

pH + pOH

The relationship that states pH + pOH = 14.00.

Calculation of [H3O+]

To calculate the concentration of H3O+, use the formula [H3O+] = 10^-pH.

Calculation of Ka

To calculate the Ka value of an acid, use the equation Ka = 10^-pKa.

Example of pKa calculation

Calculate the pKa for HF given that the Ka for this acid is 6.80 × 10^-4.

Neutral Solution

When the salts of a strong acid and strong base are placed in water, the resulting solution is neutral.

Buffer

An aqueous solution containing substances that prevent major changes in solution pH when small amounts of acid or base are added to it.

Buffer Components

Typically, a buffer system is composed of a weak acid and its conjugate base.

Active Chemical Species in Buffers

A substance to react with and remove added base and a substance to react with and remove added acid.

Addition of Base to Buffer

The added OH- ion reacts with H3O+ ion, producing water (neutralization), shifting the equilibrium to produce more H3O+ ion.

Addition of Acid to Buffer

The added H3O+ ion increases the overall amount of H3O+ ion present, shifting the equilibrium to consume excess H3O+ ion.

Electrolyte

Substance whose aqueous solution conducts electricity.

Nonelectrolytes

They do not conduct electricity because there are no ions when dissolved. Example - Table sugar (sucrose) and glucose.

Strong Electrolytes

They completely ionize/dissociate into ions, resulting in high conductivity in solution. Example - Strong acids, bases, and soluble salts.

Weak Electrolytes

They incompletely ionize/dissociate into ions, resulting in low conductivity in solution. Example - Weak acids and bases.

Henderson-Hasselbalch Equation

A formula used to calculate the pH of a buffer solution.

Hydronium Ion

An ion formed when an acid donates a proton to water.

Hydroxide Ion

An ion formed when a base accepts a proton from water.

Conjugate Base

The species that remains after an acid donates a proton.

Conjugate Acid

The species that is formed when a base accepts a proton.

Acidic Solution

A solution with a pH less than 7.

Basic Solution

A solution with a pH greater than 7.

Equilibrium Shift

The change in the position of equilibrium in a chemical reaction in response to a change in concentration, temperature, or pressure.

Stress on Buffer System

The addition of acid or base that disrupts the equilibrium of the buffer solution.