Acids, Bases, and Equilibrium: Study Notes for General Chemistry

Acids and Bases: Fundamental Concepts

Introduction to Acids and Bases

Acids and bases are essential chemical species in aqueous chemistry, defined by their ability to donate or accept protons (H+). The behavior of acids and bases is crucial for understanding chemical reactions, biological systems, and industrial processes.

• Acidic Hydrogen: A hydrogen atom bonded to a strongly electronegative atom (such as oxygen) can be released as a proton (H+) if the resulting anion is stable.

• Example: In acetic acid (CH3COOH), the hydrogen attached to the oxygen is acidic.

Arrhenius Definition of Acids and Bases

The Arrhenius definition is one of the earliest ways to classify acids and bases in aqueous solutions.

• Acid: Contains a hydrogen atom and dissolves in water to form H+ ions. $ext{HCl}(g) ightarrow ext{H}^+(aq) + ext{Cl}^-(aq)$

• Base: Contains hydroxide and dissolves in water to form OH- ions. $ext{NaOH}(s) ightarrow ext{Na}^+(aq) + ext{OH}^-(aq)$

Brønsted–Lowry Definition of Acids and Bases

The Brønsted–Lowry definition expands the concept of acids and bases beyond aqueous solutions.

• Acid: Proton (H+) donor.

• Base: Proton (H+) acceptor.

• Example: $ext{HCl}(g) + ext{H}_2 ext{O}(l) ightarrow ext{H}_3 ext{O}^+(aq) + ext{Cl}^-(aq)$ HCl donates a proton (acid), H2O accepts a proton (base).

• Requirement: Brønsted–Lowry acids must contain a hydrogen atom; bases must have a lone pair of electrons to bond with H+.

Common Brønsted–Lowry Bases

• Neutral Compounds with Lone Pairs: NH3 (ammonia), H2O (water).

• Metal Hydroxides: NaOH (sodium hydroxide), KOH (potassium hydroxide), Mg(OH)2 (magnesium hydroxide), Ca(OH)2 (calcium hydroxide).

Conjugate Acid–Base Pairs

Definition and Identification

Conjugate acid–base pairs are species on opposite sides of a chemical reaction whose formulas differ by one H+.

• Example: $ext{HF}(g) + ext{OH}^-(aq) ightleftharpoons ext{F}^-(aq) + ext{H}_2 ext{O}(l)$ HF and F- are a conjugate acid–base pair; OH- and H2O are another.

• Example: $ext{CH}_3 ext{COOH} + ext{H}_2 ext{O} ightleftharpoons ext{CH}_3 ext{COO}^- + ext{H}_3 ext{O}^+$

Classification of Acids: Monoprotic, Diprotic, Triprotic

Types of Acids

Acids are classified by the number of protons (H+) they can donate.

• Monoprotic: One proton to donate (e.g., HCl, HNO3).

• Diprotic: Two protons to donate (e.g., H2SO4).

• Triprotic: Three protons to donate (e.g., H3PO4).

Naming Acids

Binary Acids

Binary acids contain hydrogen and one other nonmetal element.

• Anions ending in -ide: Add prefix hydro- and change -ide to -ic acid.

Formula	Anion	Acid Name
HF	F- (fluoride)	Hydrofluoric acid
HCl	Cl- (chloride)	Hydrochloric acid
HI	I- (iodide)	Hydroiodic acid

Polyatomic Acids

Polyatomic acids contain hydrogen, oxygen, and another element.

• Anions ending in -ate: Change -ate to -ic acid.

• Anions ending in -ite: Change -ite to -ous acid.

Formula	Anion	Acid Name
HNO3	NO3- (nitrate)	Nitric acid
H2SO4	SO42- (sulfate)	Sulfuric acid
H2SO3	SO32- (sulfite)	Sulfurous acid

Names of Common Acids and Their Anions

Acid	Name of Acid	Anion	Name of Anion
HCl	Hydrochloric acid	Cl-	Chloride
HBr	Hydrobromic acid	Br-	Bromide
HI	Hydroiodic acid	I-	Iodide
HCN	Hydrocyanic acid	CN-	Cyanide
HNO3	Nitric acid	NO3-	Nitrate
H2SO4	Sulfuric acid	SO42-	Sulfate
H2SO3	Sulfurous acid	SO32-	Sulfite
H3PO4	Phosphoric acid	PO43-	Phosphate
CH3COOH	Acetic acid	CH3COO-	Acetate

Acid and Base Strength

Strong Acids and Bases

Strong acids and bases dissociate completely in water, producing high concentrations of ions.

• Strong Acid Example: $ext{HI}(aq) + ext{H}_2 ext{O}(l) ightarrow ext{H}_3 ext{O}^+(aq) + ext{I}^-(aq)$

• Strong Base Example: $ext{NaOH}(s) + ext{H}_2 ext{O}(l) ightarrow ext{Na}^+(aq) + ext{OH}^-(aq)$

• List of Strong Acids: Hydroiodic acid (HI), Hydrobromic acid (HBr), Hydrochloric acid (HCl), Sulfuric acid (H2SO4), Nitric acid (HNO3), Perchloric acid (HClO4).

Weak Acids and Bases

Weak acids and bases only partially dissociate in water, resulting in equilibrium between reactants and products.

• Weak Acid Example: $ext{HF}(aq) + ext{H}_2 ext{O}(l) ightleftharpoons ext{H}_3 ext{O}^+(aq) + ext{F}^-(aq)$

• Weak Base Example: $ext{NH}_3(g) + ext{H}_2 ext{O}(l) ightleftharpoons ext{NH}_4^+(aq) + ext{OH}^-(aq)$

pH Scale and Calculations

Definition and Formula

The pH scale quantifies the acidity or basicity of a solution, ranging from 0 (most acidic) to 14 (most basic).

• Formula: $ext{pH} = - ext{log}[ ext{H}_3 ext{O}^+]$

• Neutral Solution: pH = 7

• Acidic Solution: pH < 7

• Basic Solution: pH > 7

The Self-Ionization of Water

Water can self-ionize, producing equal concentrations of H+ and OH- ions.

• Equation: $ext{H}_2 ext{O} + ext{H}_2 ext{O} ightleftharpoons ext{H}_3 ext{O}^+ + ext{OH}^-$

• Equilibrium Constant: $K_w = [ ext{H}_3 ext{O}^+][ ext{OH}^-] = 1.0 imes 10^{-14}$

• Relationship: $ext{pH} + ext{pOH} = 14$

Neutral, Acidic, and Basic Solutions

Type of Solution	[H3O+]	[OH-]	Kw (25°C)
Neutral	1.0 × 10-7 M	1.0 × 10-7 M	1.0 × 10-14
Acidic	2.5 × 10-5 M	4.0 × 10-10 M	1.0 × 10-14
Basic	5.0 × 10-10 M	2.0 × 10-5 M	1.0 × 10-14

[H3O+], [OH-], and pH Comparison

[H3O+]	pH	[OH-]
100	0	10-14
10-1	1	10-13
10-2	2	10-12
10-3	3	10-11
10-4	4	10-10
10-5	5	10-9
10-6	6	10-8
10-7	7	10-7
10-8	8	10-6
10-9	9	10-5
10-10	10	10-4
10-11	11	10-3
10-12	12	10-2
10-13	13	10-1
10-14	14	100

Acid–Base Equilibrium and Le Châtelier’s Principle

Equilibrium of Weak Acids

Weak acids establish an equilibrium in solution, with both forward and reverse reactions occurring simultaneously.

• General Reaction: $ext{HF}(aq) + ext{H}_2 ext{O}(l) ightleftharpoons ext{F}^-(aq) + ext{H}_3 ext{O}^+(aq)$

• At equilibrium, the rates of the forward and reverse reactions are equal, and concentrations remain constant.

Le Châtelier’s Principle

When a system at equilibrium is disturbed, it will shift to counteract the disturbance and reestablish equilibrium.

• Adding more reactant (HF): Increases the rate of the forward reaction.

• Removing reactant (HF): Increases the rate of the reverse reaction.

• Adding or removing product (F-): Shifts equilibrium accordingly.

Reactions of Acids

Acids and Metals

Acids react with metals to produce a salt and hydrogen gas.

• Example: $2 ext{K}(s) + 2 ext{HCl}(aq) ightarrow 2 ext{KCl}(aq) + ext{H}_2(g)$

• Example: $ext{Zn}(s) + 2 ext{HCl}(aq) ightarrow ext{ZnCl}_2(aq) + ext{H}_2(g)$

Acids and Hydroxides: Neutralization

Acids react with bases to produce a salt and water in a neutralization reaction.

• Example: $ext{HCl}(aq) + ext{NaOH}(aq) ightarrow ext{NaCl}(aq) + ext{H}_2 ext{O}(l)$

Acid–Base Titrations

Principles and Calculations

Titration is a laboratory technique used to determine the concentration (molarity) of an unknown acid or base by reacting it with a solution of known concentration.

• End Point: The stage in titration when the indicator changes color, signaling completion.

• Indicator: A substance that changes color at (or near) the equivalence point.

• Calculation Example: $ext{NaOH} + ext{HCl} ightarrow ext{NaCl} + ext{H}_2 ext{O}$ $0.02350~L imes rac{0.1~ ext{moles NaOH}}{1~L} imes rac{1~ ext{mol HCl}}{1~ ext{mol NaOH}} imes rac{1}{0.05~L~ ext{HCl}} = 0.047~M~ ext{HCl}$

Buffer Solutions

Definition and Function

Buffers are solutions that resist changes in pH when small amounts of acid or base are added. They typically consist of a weak acid and its conjugate base.

• Example: CH3COOH/CH3COO- buffer system.

• Buffers are essential in biological and chemical systems to maintain stable pH conditions.

Buffer in Action

• When acid is added, the conjugate base neutralizes it.

• When base is added, the weak acid neutralizes it.

Practice Problems and Applications

Sample Calculations

• Find the pH of a 0.15 M H+ solution: $ext{pH} = - ext{log}(0.15)$

• Find the pH of a 5.0 × 10-3 M H+ solution: $ext{pH} = - ext{log}(5.0 imes 10^{-3})$

• Find the pH of a 3.0 × 10-8 M OH- solution: First, calculate pOH: $ext{pOH} = - ext{log}(3.0 imes 10^{-8})$ Then, $ext{pH} = 14 - ext{pOH}$

• Titration Example: How many mL of 0.150 M NaOH are needed to neutralize 50.00 mL of a 0.120 M solution of H2SO4? Additional info: Use stoichiometry and molarity relationships.

• Buffer Solution Identification: A buffer is produced by combining a weak acid and its salt (conjugate base), e.g., H2CO3 and NaHCO3.

Summary Table: Key Acid–Base Concepts

Concept	Definition/Example
Arrhenius Acid	Produces H+ in water (HCl)
Arrhenius Base	Produces OH- in water (NaOH)
Brønsted–Lowry Acid	Proton donor (HCl)
Brønsted–Lowry Base	Proton acceptor (NH3)
Strong Acid	Complete dissociation (HNO3)
Weak Acid	Partial dissociation (HF)
Strong Base	Complete dissociation (NaOH)
Weak Base	Partial dissociation (NH3)
Buffer	Resists pH change (CH3COOH/CH3COO-)

Additional info: These notes cover the essential concepts of acids, bases, equilibrium, titration, and buffer solutions as presented in a General Chemistry course.