Chapter 14: Acids and Bases

Acids and Bases

Properties of Acids

• Aqueous solutions of acids have a sour taste.

• Acids change the color of acid-base indicators.

• Some acids react with active metals and release hydrogen gas ($H_2$).

  • Example: Ba(s) + H2SO4(aq) —> BaSO4(s) + H2(g)

  • The reaction demonstrates the reactivity of barium with sulfuric acid, illustrating the release of hydrogen gas as a product. Additionally, acids can neutralize bases to form salts and water, exemplifying their role in various chemical reactions. Furthermore, this neutralization process can be represented by the general equation: {Acid} + {Base} —> {Salt} + {Water}, showcasing the importance of acids and bases in producing essential compounds.

• Acids react with bases to produce salts and water.

• Acids conduct electric current.

Acid Nomenclature

• Binary Acid: An acid containing only two different elements: hydrogen and a more electronegative element (e.g., HF, HCl, HBr, HI).

  • Binary Acid Nomenclature:

    1. Begins with the prefix hydro-

    2. Followed by the root name of the second element.

    3. Ends with the suffix -ic.

• Examples:

  • HF: hydrofluoric acid

  • HCl: hydrochloric acid

  • HBr: hydrobromic acid

  • HI: hydriodic acid

  • H₂S: hydrosulfuric acid

• Oxyacid: An acid that is a compound of hydrogen, oxygen, and a third element (usually a nonmetal) (e.g., $HNO<em>3$, $H</em>2SO_4$).

  • The names of oxyacids follow a pattern, and the names of their anions are based on the names of the acids.

• Examples:

  • $CH<em>3COOH$: acetic acid, anion: $CH</em>3COO^-$, acetate

  • $H<em>2CO</em>3$: carbonic acid, anion: $CO_3^{2-}$, carbonate

  • $HIO<em>3$: iodic acid, anion: $IO</em>3^{-}$, iodate

  • $HClO$: hypochlorous acid, anion: $ClO^{-}$, hypochlorite

  • $HClO<em>2$: chlorous acid, anion: $ClO</em>2^{-}$, chlorite

  • $HClO<em>3$: chloric acid, anion: $ClO</em>3^{-}$, chlorate

  • $HClO<em>4$: perchloric acid, anion: $ClO</em>4^{-}$, perchlorate

  • $HNO<em>2$: nitrous acid, anion: $NO</em>2^{-}$, nitrite

  • $HNO<em>3$: nitric acid, anion: $NO</em>3^{-}$, nitrate

  • $H<em>3PO</em>3$: phosphorous acid, anion: $PO_3^{3-}$, phosphite

  • $H<em>3PO</em>4$: phosphoric acid, anion: $PO_4^{3-}$, phosphate

  • $H<em>2SO</em>3$: sulfurous acid, anion: $SO_3^{2-}$, sulfite

  • $H<em>2SO</em>4$: sulfuric acid, anion: $SO_4^{2-}$, sulfate

Common Industrial Acids

• Sulfuric Acid ($H<em>2SO</em>4$): The most commonly produced industrial chemical worldwide.

• Nitric Acid ($HNO_3$)

• Phosphoric Acid ($H<em>3PO</em>4$)

• Hydrochloric Acid (HCl): Concentrated solutions commonly referred to as muriatic acid.

• Acetic Acid ($CH_3COOH$): Pure acetic acid is a clear, colorless, pungent-smelling liquid known as glacial acetic acid.

Properties of Bases

• Aqueous solutions of bases taste bitter.

• Bases change the color of acid-base indicators.

• Dilute aqueous solutions of bases feel slippery.

• Bases react with acids to produce salts and water.

• Bases conduct electric current.

Arrhenius Acids and Bases

• Arrhenius Acid: A compound that increases the concentration of hydrogen ions ($H^+$) in aqueous solution.

• Arrhenius Base: A substance that increases the concentration of hydroxide ions ($OH^−$) in aqueous solution.

• Arrhenius acids are molecular compounds with ionizable hydrogen atoms; their water solutions are known as aqueous acids which are all electrolytes.

Strength of Acids

• Strong Acid: Ionizes completely in aqueous solution; a strong electrolyte (e.g., $HClO<em>4$, HCl, $HNO</em>3$).

• Weak Acid: Releases few hydrogen ions in aqueous solution (e.g., HCN, organic acids such as acetic acid).

Aqueous Solutions of Bases

• Most bases are ionic compounds containing metal cations and the hydroxide anion ($OH^−$), which dissociate in water.

• Ammonia ($NH_3$) is molecular and produces hydroxide ions when it reacts with water molecules.

Strength of Bases

• The strength of a base depends on the extent to which the base dissociates; strong bases are strong electrolytes.

Relationship of Hydronium and Hydroxide Ion Concentrations

• Acidic solution: $[H_3O^+] > 10^{-7} M > [OH^-]$

• Neutral solution: $[H_3O^+] = 10^{-7} M = [OH^-]$

• Basic solution: $[H_3O^+] < 10^{-7} M < [OH^-]$

Brønsted-Lowry Acids and Bases

• Brønsted-Lowry Acid: A molecule or ion that is a proton donor.

  • Example: Hydrogen chloride (HCl) acts as a Brønsted-Lowry acid when it reacts with ammonia.

  • Water can also act as a Brønsted-Lowry acid.

• Brønsted-Lowry Base: A molecule or ion that is a proton acceptor.

  • Example: Ammonia accepts a proton from hydrochloric acid and acts as a Brønsted-Lowry base.

  • The $OH^−$ ion produced in solution by Arrhenius hydroxide bases (e.g., NaOH) is a Brønsted-Lowry base because it can accept a proton.

• In a Brønsted-Lowry acid-base reaction, protons are transferred from one reactant (the acid) to another (the base).

Monoprotic and Polyprotic Acids

• Monoprotic Acid: An acid that can donate only one proton (hydrogen ion) per molecule (e.g., $HClO<em>4$, HCl, $HNO</em>3$).

  • Involves only one ionization step.

• Polyprotic Acid: An acid that can donate more than one proton per molecule (e.g., $H<em>2SO</em>4$, $H<em>3PO</em>4$).

  • Involves multiple ionization steps.

  • Sulfuric acid solutions contain $H_3O^+$ ions.

• Diprotic Acid: A polyprotic acid that can donate two protons per molecule (e.g., $H<em>2SO</em>4$).

• Triprotic Acid: A polyprotic acid that can donate three protons per molecule (e.g., $H<em>3PO</em>4$).

Lewis Acids and Bases

• Lewis Acid: An atom, ion, or molecule that accepts an electron pair to form a covalent bond.

  • The Lewis definition is the broadest of the three acid definitions.

  • A bare proton (hydrogen ion) is a Lewis acid.

  • The formula for a Lewis acid need not include hydrogen (e.g., silver ion).

  • Any compound in which the central atom has three valence electrons and forms three covalent bonds can react as a Lewis acid.

• Lewis Base: An atom, ion, or molecule that donates an electron pair to form a covalent bond.

Acid Base Definitions Comparison

Conjugate Acids and Bases

• The species that remains after a Brønsted-Lowry acid has given up a proton is the conjugate base of that acid.

• Brønsted-Lowry acid-base reactions involve two acid-base pairs, known as conjugate acid-base pairs.

$acid<em>1 + base</em>2 \rightleftharpoons base<em>1 + acid</em>2$

Strength of Conjugate Acids and Bases

• The stronger an acid is, the weaker its conjugate base.

• The stronger a base is, the weaker its conjugate acid.

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

Relative Strengths of Acids and Bases

*Strong acids
† Strong bases

Amphoteric Compounds

• Any species that can react as either an acid or a base is described as amphoteric (e.g., water).

• Water can act as a base:

  • $acid<em>1 + base</em>2 \rightleftharpoons acid<em>2 + base</em>1$

• Water can act as an acid:

  • $base<em>1 + acid</em>2 \rightleftharpoons acid<em>1 + base</em>2$

Hydroxyl Group

• The covalently bonded —OH group in an acid is referred to as a hydroxyl group.

• Molecular compounds containing —OH groups can be acidic or amphoteric.

• The behavior of a compound is affected by the number of oxygen atoms bonded to the atom connected to the —OH group.

Oxyacids of Chlorine

Acidity increases as more oxygen atoms are present.

• Hypochlorous acid

• Chlorous acid

• Chloric acid

• Perchloric acid

Neutralization Reactions

• In aqueous solutions, neutralization is the reaction of hydronium ions and hydroxide ions to form water molecules.

• A salt is an ionic compound composed of a cation from a base and an anion from an acid.

Acid Rain

• $NO, NO<em>2, CO</em>2, SO<em>2$, and $SO</em>3$ gases from industrial processes can dissolve in atmospheric water to produce acidic solutions.

• Very acidic rain is known as acid rain.

• Acid rain can erode statues and affect ecosystems.