Comprehensive Acid-Base Chemistry: Theories, Properties, and Calculations

Acid

A substance which releases hydrogen ions (H+) in solution.

Base

A substance which releases hydroxide ions (OH-) in solution.

Neutralization

A reaction between an acid and a base that forms a salt and water.

Common Characteristics of Acids

1) Acids taste sour; 2) Acids turn litmus red; 3) Acids react with active metals to form hydrogen gas; 4) Acids react with bases to form salts and water.

Common Characteristics of Bases

1) Bases taste bitter; 2) Bases turn litmus blue; 3) Bases feel slippery; 4) Bases react with acids to form salts and water.

Arrhenius Theory

A theory that defines acids as substances that release H+ ions and bases as substances that release OH- ions in solution.

Bronsted-Lowry Theory

A theory that defines acids as proton donors and bases as proton acceptors.

Lewis Theory

A theory that expands the definition of acids and bases to include substances that are electron pair acceptors (acids) and electron pair donors (bases).

Hydrochloric Acid

Commonly known as stomach acid, its chemical formula is HCl.

Acetic Acid

Commonly known as vinegar, its chemical formula is HC2H3O2.

Carbonic Acid

Found in soda water, its chemical formula is H2CO3.

Sodium Hydroxide

Commonly known as Draino, its chemical formula is NaOH.

Ammonia Water

Commonly used as a cleaning agent, its chemical formula is NH4OH.

Aluminum Hydroxide

Commonly found in Rolaids, its chemical formula is Al(OH)3.

Magnesium Hydroxide

Commonly found in Tums, its chemical formula is Mg(OH)2.

Hydronium Ion

Formed when water acts as a base and accepts a proton, represented as H3O+.

Conjugate Acid

The species formed when a base accepts a proton.

Conjugate Base

The species formed when an acid donates a proton.

Amphiprotic

A term describing a substance that can act as either an acid or a base depending on the situation.

Proton Transfer

The process in which a proton is transferred from an acid to a base during an acid-base reaction.

Electron Pair Acceptor

A substance that accepts an electron pair, typically described as an acid in Lewis theory.

Electron Pair Donor

A substance that donates an electron pair, typically described as a base in Lewis theory.

Ammonia

An electron pair donor that does not contain hydroxide ions.

Acid Anhydrides

Nonmetal oxides that form acids when added to water.

Carbon Dioxide Reaction

CO2(g) + H2O(l) 🡪 H2CO3(aq) (Carbonic Acid) forms a mildly acidic solution.

Sulfur Dioxide Reaction

SO2(g) + H2O(l) 🡪 H2SO3(aq) (Sulfurous Acid) forms an acid.

Basic Anhydrides

Metallic oxides that form bases when added to water.

Calcium Hydroxide Reaction

CaO(s) + H2O(l) 🡪 Ca(OH)2(s) (Calcium Hydroxide) is a base.

Sodium Hydroxide Reaction

Na2O(s) + H2O(l) 🡪 2 NaOH(s) (Sodium Hydroxide) is a base.

Dissociation

The process where dissolved substances separate into free mobile ions.

Acid Strength

Depends on the ability to dissociate in solution.

Concentration

Refers to the molarity of the solution.

Strong Acids

Acids that dissociate well (~ 100%).

Weak Acids

All other acids that dissociate poorly.

Perchloric Acid

HCLO4 is a strong acid.

Hydroiodic Acid

HI is a strong acid.

Hydrobromic Acid

HBr is a strong acid.

Nitric Acid

HNO3 is a strong acid.

Sulfuric Acid

H2SO4 is a strong acid.

Strong Bases

Bases that dissociate well (~ 100%).

Lithium Hydroxide

LiOH is a strong base.

Potassium Hydroxide

KOH is a strong base.

pH

A system for measuring the acidity of a solution.

pH Definition

Defined as the negative logarithm of the hydrogen ion concentration in a solution.

Logarithm

A power of 10; if a number is written as 10^x, then its log is x.

pH Calculation Example

For [H+] = 0.000001 M, pH = -LOG(10^-6) = 6.00.

pH Scale

Ranges from 0 to 14, with lower values indicating higher acidity.

Pure Water pH

pH = 7.00 with [H+] = 1.00 x 10^-7 M.

pOH

Defined as -LOG[OH-], where brackets indicate concentration of OH-.

pOH Calculation Example

For [OH-] = 0.00001 M, calculate pOH.

Hydrogen Ion

An aqueous proton formed when a hydrogen atom loses its electron.

Acid-Base Neutralization

H+(aq) + OH-(aq) 🡪 H2O(l)

Binary Acids

Consist of two elements, one of which is hydrogen.

Ternary Acids

Consist of three elements, one of which is hydrogen and another oxygen, also called oxyacids.

Naming Binary Acids

Begin with 'HYDRO' and end in 'IC'; e.g., H2S is 'Hydrosulfuric Acid'.

Naming Ternary Acids

If the anion ends in 'ATE', the acid name ends in 'IC'; e.g., H2SO4 is 'Sulfuric Acid'.

Sulfite Ion

SO3-2, associated with the acid name 'Sulfurous Acid'.

Nitrate Ion

NO3-, associated with the acid name 'Nitric Acid'.

Nitrite Ion

NO2-, associated with the acid name 'Nitrous Acid'.

Degree of Dissociation

The extent to which an acid dissociates into ions in solution.

Conjugate Base Strength

Inversely related to the strength of its corresponding acid.

Kb Values

Measure the strength of bases in solution.

HClO4

Strong acid that dissociates into H+ and ClO4-.

HCl

Strong acid that dissociates into H+ and Cl-.

HF

Weak acid that dissociates into H+ and F-.

HCOOH

Weak acid that dissociates into H+ and HCOO-.

HC2H3O2

Weak acid that dissociates into H+ and C2H3O2-.

NH4+

Weak acid that dissociates into H+ and NH3.

Kw

The equilibrium constant for the dissociation of water, always equal to 1.00 x 10-14.

Ka

The acid dissociation constant, a measure of the strength of an acid in solution.

Strong Acid

An acid that completely dissociates in solution, indicated by an undefined Ka value.

Hydrofluoric Acid

A weak acid with a Ka of 6.7 x 10-4.

Formic Acid

A weak acid with a Ka of 1.8 x 10-4.

Benzoic Acid

A weak acid with a Ka of 6.0 x 10-5.

pH of Strong Acid

For a 2.0 liters solution of nitric acid containing 15.75 grams, pH = 0.903.

Molar Mass of HNO3

The molar mass of nitric acid is 63.0 grams.

Hydrogen Ion Concentration

In pure water, [H+] = [OH-] = 1.00 x 10-7 M.

Kw Calculation

Kw = [H+] x [OH-]; for pure water, (1.00 x 10-7)(1.00 x 10-7) = 1.00 x 10-14.

Acidic Solution

Solutions with [H+] > 1.00 x 10-7 M and [OH-] < 1.00 x 10-7 M are acidic.

Basic Solution

Solutions with [H+] < 1.00 x 10-7 M and [OH-] > 1.00 x 10-7 M are basic.

Weak Acid Dissociation

For a weak acid HX, Ka = [H+] x [X-] / [HX].

pH of Acetic Acid

For a 0.10 M solution of acetic acid, pH = 2.87.

pH of Ammonia Solution

The pH of a 0.10 M ammonia solution is 11.37.

Kb for NH3

The base dissociation constant for ammonia, Kb = 1.8 x 10-5.

Titration

The addition of an acid and base in measured quantities, often to an endpoint where moles of acid and base are equal.

Endpoint of Titration

The point in titration where the moles of added acid and base are equal, not always at pH = 7.00.

Strong Acid with Strong Base Titration

Results in a neutral solution at the endpoint (pH = 7.00).

Strong Acid with Weak Base Titration

Results in an acidic solution at the endpoint (pH < 7.00).

Weak Acids and Bases Titration

The resulting acidity of the solution is determined by the relative strengths of the acids and bases involved.

Strong Base

A base that completely dissociates in solution, such as KOH.

End Point

The point in a titration at which the reaction is complete.

Neutral

A solution with a pH of 7.00.

Acidic

A solution with a pH less than 7.00.

Basic

A solution with a pH greater than 7.00.

Weak Acid

An acid that partially dissociates in solution, such as HF.

Weak Base

A base that partially dissociates in solution, such as NH3.

Normality

A system of measuring the concentration of solutions, often used in titrations.

Normality of an Acid

Equal to the moles of hydrogen ions available per liter of solution.

Normality of a Base

Equal to the moles of hydroxide ions available per liter of solution.

Equivalent

A mole of H+ ions or OH- ions.