Acids and Bases: Properties, Nomenclature, and Theories

Chapter 14: Acids and Bases

Section 1: Properties of Acids and Bases

Acids and bases are two fundamental categories of compounds in chemistry, each with distinct properties and behaviors in aqueous solutions.

Properties of Acids

• Sour taste: Many acids, such as citric acid in lemons, have a characteristic sour taste.

• Indicator color change: Acids change the color of acid-base indicators (e.g., litmus turns red).

• Reaction with metals: Acids react with active metals to release hydrogen gas, H2. For example: $\mathrm{Ba(s) + H_2SO_4(aq) \rightarrow BaSO_4(s) + H_2(g)}$

• Neutralization: Acids react with bases to produce salts and water.

• Electrical conductivity: Acids conduct electric current in solution due to ionization.

Properties of Bases

• Bitter taste: Bases typically taste bitter.

• Indicator color change: Bases change the color of acid-base indicators (e.g., litmus turns blue).

• Slippery feel: Dilute solutions of bases feel slippery to the touch.

• Neutralization: Bases react with acids to produce salts and water.

• Electrical conductivity: Bases conduct electric current in solution due to dissociation into ions.

Acid Nomenclature

Acids are named based on their composition and the anion they contain. There are two main types: binary acids and oxyacids.

Binary Acids

• Contain hydrogen and one other, more electronegative element.

• Naming rules:

  1. Prefix "hydro-"

  2. Root of the second element's name

  3. Suffix "-ic"

Formula	Acid name
HF	hydrofluoric acid
HCl	hydrochloric acid
HBr	hydrobromic acid
HI	hydroiodic acid
H2S	hydrosulfuric acid

Oxyacids

• Contain hydrogen, oxygen, and a third element (usually a nonmetal).

• The name is based on the polyatomic anion present.

Formula	Acid name	Anion
CH3COOH	acetic acid	CH3COO−, acetate
H2CO3	carbonic acid	CO32−, carbonate
HNO3	nitric acid	NO3−, nitrate
H2SO4	sulfuric acid	SO42−, sulfate

Common Industrial Acids

• Sulfuric acid (H2SO4): Most produced industrial chemical.

• Nitric acid (HNO3): Used in fertilizers and explosives.

• Phosphoric acid (H3PO4): Used in soft drinks and detergents.

• Hydrochloric acid (HCl): Also known as muriatic acid, used in cleaning and processing steel.

• Acetic acid (CH3COOH): Pure form is called glacial acetic acid, used in vinegar.

Arrhenius Acids and Bases

The Arrhenius definition is one of the earliest and simplest ways to classify acids and bases based on their behavior in water.

• Arrhenius acid: Increases the concentration of hydrogen ions (H+) or hydronium ions (H3O+) in aqueous solution.

• Arrhenius base: Increases the concentration of hydroxide ions (OH−) in aqueous solution.

Strength of Acids and Bases

The strength of an acid or base depends on its degree of ionization or dissociation in water.

Strong and Weak Acids

• Strong acids: Ionize completely in aqueous solution (e.g., HCl, HNO3, HClO4).

• Weak acids: Ionize only partially, producing fewer hydrogen ions (e.g., CH3COOH, HCN).

Strong acids		Weak acids
HI + H2O → H3O+ + I−		HSO4− + H2O ⇌ H3O+ + SO42−

Strong and Weak Bases

• Strong bases: Dissociate completely in water (e.g., NaOH, KOH, Ba(OH)2).

• Weak bases: Only partially react with water to produce hydroxide ions (e.g., NH3).

Strong bases		Weak bases
Ca(OH)2 → Ca2+ + 2OH−		NH3 + H2O ⇌ NH4+ + OH−

Relationship of [H3O+] to [OH−]

The concentrations of hydronium and hydroxide ions determine whether a solution is acidic, neutral, or basic.

• Acidic solution: [H3O+] > [OH−]

• Neutral solution: [H3O+] = [OH−] = $1 \times 10^{-7}$ M

• Basic solution: [H3O+] < [OH−]

Section 2: Acid-Base Theories

There are three main theories for defining acids and bases: Arrhenius, Brønsted-Lowry, and Lewis.

Type	Acid	Base
Arrhenius	H+ or H3O+ producer	OH− producer
Brønsted-Lowry	proton (H+) donor	proton (H+) acceptor
Lewis	electron-pair acceptor	electron-pair donor

Brønsted-Lowry Acids and Bases

• Acid: Proton (H+) donor

• Base: Proton (H+) acceptor

• Example: $\mathrm{HCl + NH_3 \rightarrow NH_4^+ + Cl^-}$

Lewis Acids and Bases

• Lewis acid: Electron-pair acceptor (may not contain hydrogen)

• Lewis base: Electron-pair donor

• Example: $\mathrm{BF_3 + F^- \rightarrow BF_4^-}$

Monoprotic and Polyprotic Acids

Acids are classified by the number of protons (hydrogen ions) they can donate per molecule.

• Monoprotic acid: Donates one proton (e.g., HCl, HNO3).

• Diprotic acid: Donates two protons (e.g., H2SO4).

• Triprotic acid: Donates three protons (e.g., H3PO4).

Section 3: Acid-Base Reactions

Conjugate Acids and Bases

In Brønsted-Lowry theory, acids and bases exist in conjugate pairs. When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid.

• Conjugate acid: The species formed when a base gains a proton.

• Conjugate base: The species formed when an acid loses a proton.

• Example: $\mathrm{HF + H_2O \rightleftharpoons F^- + H_3O^+}$

Strength of Conjugate Acids and Bases

• The stronger an acid, the weaker its conjugate base, and vice versa.

• Proton transfer reactions favor the formation of the weaker acid and base.

Conjugate acid	Formula	Conjugate base	Formula
Hydriodic acid	HI	Iodide ion	I−
Perchloric acid	HClO4	Perchlorate ion	ClO4−

Amphoteric Compounds

Amphoteric substances can act as either acids or bases depending on the reaction. Water is a classic example.

• Hydroxyl group (–OH): The presence and bonding of –OH groups influence whether a compound is acidic, basic, or amphoteric.

Oxyacids of Chlorine

Oxyacids of chlorine demonstrate how increasing the number of oxygen atoms increases acidity.

Neutralization Reactions

Neutralization is the reaction between hydronium ions (from acids) and hydroxide ions (from bases) to form water. The other product is a salt, an ionic compound composed of a cation from the base and an anion from the acid.

• General equation: $\mathrm{HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)}$