63d ago

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 (H2H_2$$H_2$$).

    • 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., $$HNO3,, $$, $$H2SO_4$$).

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

  • Examples:

    • $$CH3COOH:aceticacid,anion:: acetic acid, anion: $$: acetic acid, anion: $$ CH3COO^-$$, acetate

    • $$H2CO3:carbonicacid,anion:: carbonic acid, anion: $$: carbonic acid, anion: $$CO_3^{2-}$$, carbonate

    • $$HIO3:iodicacid,anion:: iodic acid, anion: $$: iodic acid, anion: $$IO3^{-}$$, iodate

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

    • $$HClO2:chlorousacid,anion:: chlorous acid, anion: $$: chlorous acid, anion: $$ClO2^{-}$$, chlorite

    • $$HClO3:chloricacid,anion:: chloric acid, anion: $$: chloric acid, anion: $$ClO3^{-}$$, chlorate

    • $$HClO4:perchloricacid,anion:: perchloric acid, anion: $$: perchloric acid, anion: $$ClO4^{-}$$, perchlorate

    • $$HNO2:nitrousacid,anion:: nitrous acid, anion: $$: nitrous acid, anion: $$NO2^{-}$$, nitrite

    • $$HNO3:nitricacid,anion:: nitric acid, anion: $$: nitric acid, anion: $$NO3^{-}$$, nitrate

    • $$H3PO3:phosphorousacid,anion:: phosphorous acid, anion: $$: phosphorous acid, anion: $$PO_3^{3-}$$, phosphite

    • $$H3PO4:phosphoricacid,anion:: phosphoric acid, anion: $$: phosphoric acid, anion: $$PO_4^{3-}$$, phosphate

    • $$H2SO3:sulfurousacid,anion:: sulfurous acid, anion: $$: sulfurous acid, anion: $$SO_3^{2-}$$, sulfite

    • $$H2SO4:sulfuricacid,anion:: sulfuric acid, anion: $$: sulfuric acid, anion: $$SO_4^{2-}$$, sulfate

Common Industrial Acids

  • Sulfuric Acid ($$H2SO4$$): The most commonly produced industrial chemical worldwide.

  • Nitric Acid (HNO3HNO_3$$HNO_3$$)

  • Phosphoric Acid ($$H3PO4$$)

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

  • Acetic Acid (CH3COOHCH_3COOH$$CH_3COOH$$): 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+H^+$$H^+$$) in aqueous solution.

  • Arrhenius Base: A substance that increases the concentration of hydroxide ions (OHOH^−$$OH^−$$) 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., $$HClO4,HCl,, HCl, $$, HCl, $$HNO3$$).

  • 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 (OHOH^−$$OH^−$$), which dissociate in water.

  • Ammonia (NH3NH_3$$NH_3$$) 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: [H3O+]>107M>[OH][H_3O^+] > 10^{-7} M > [OH^-]$$[H_3O^+] > 10^{-7} M > [OH^-]$$

  • Neutral solution: [H3O+]=107M=[OH][H_3O^+] = 10^{-7} M = [OH^-]$$[H_3O^+] = 10^{-7} M = [OH^-]$$

  • Basic solution: [H3O+]<107M<[OH][H_3O^+] < 10^{-7} M < [OH^-]$$[H_3O^+] < 10^{-7} M < [OH^-]$$

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 OHOH^−$$OH^−$$ 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., $$HClO4,HCl,, HCl, $$, HCl, $$HNO3$$).

    • Involves only one ionization step.

  • Polyprotic Acid: An acid that can donate more than one proton per molecule (e.g., $$H2SO4,, $$, $$H3PO4$$).

    • Involves multiple ionization steps.

    • Sulfuric acid solutions contain H3O+H_3O^+$$H_3O^+$$ ions.

  • Diprotic Acid: A polyprotic acid that can donate two protons per molecule (e.g., $$H2SO4$$).

  • Triprotic Acid: A polyprotic acid that can donate three protons per molecule (e.g., $$H3PO4$$).

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

Type

Acid

Base

Arrhenius

H+H^+$$H^+$$ or H3O+H_3O^+$$H_3O^+$$ producer

OHOH^-$$OH^-$$ producer

Brønsted-Lowry

Proton (H^$$H^$$) donor

Proton (H+H^+$$H^+$$) acceptor

Lewis

Electron-pair acceptor

Electron-pair donor

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.

$$acid1 + base2 \rightleftharpoons base1 + acid2$$

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

Conjugate acid

Formula

Conjugate base

Formula

hydriodic acid*

HI

iodide ion

I-

perchloric acid*

HClO4HClO_4$$HClO_4$$

perchlorate ion

ClO4ClO_4$$ClO_4$$

hydrobromic acid*

HBr

bromide ion

Br-

hydrochloric acid*

HCl

chloride ion

Cl-

sulfuric acid*

$$H2SO4$$

hydrogen sulfate ion

HSO4HSO_4$$HSO_4$$

chloric acid*

HClO3HClO_3$$HClO_3$$

chlorate ion

ClO3ClO_3$$ClO_3$$

nitric acid*

HNO3HNO_3$$HNO_3$$

nitrate ion

NO3NO_3$$NO_3$$

hydronium ion

H3O+H_3O^+$$H_3O^+$$

water

H2OH_2O$$H_2O$$

chlorous acid

HClO2HClO_2$$HClO_2$$

chlorite ion

ClO2ClO_2$$ClO_2$$

hydrogen sulfate ion

HSO4HSO_4$$HSO_4$$

sulfate ion

SO4SO_4$$SO_4$$

phosphoric acid

$$H3PO4$$

dihydrogen phosphate ion

$$H2PO4$$

hydrofluoric acid

HF

fluoride ion

F-

acetic acid

CH3COOHCH_3COOH$$CH_3COOH$$

acetate ion

CH3COOCH_3COO^-$$CH_3COO^-$$

carbonic acid

$$H2CO3$$

hydrogen carbonate ion

HCO3HCO_3$$HCO_3$$

hydrosulfuric acid

H2SH_2S$$H_2S$$

hydrosulfide ion

HSHS$$HS$$

dihydrogen phosphate ion

$$H2PO4$$

hydrogen phosphate ion

HPO4HPO_4$$HPO_4$$

hypochlorous acid

HClOHClO$$HClO$$

hypochlorite ion

ClOClO^-$$ClO^-$$

ammonium ion

NH4+NH_4^+$$NH_4^+$$

ammonia

NH3NH_3$$NH_3$$

hydrogen carbonate ion

HCO3HCO_3$$HCO_3$$

carbonate ion

CO3CO_3$$CO_3$$

hydrogen phosphate ion

HPO4HPO_4$$HPO_4$$

phosphate ion

PO4PO_4$$PO_4$$

water

H2OH_2O$$H_2O$$

hydroxide ion

OHOH^-$$OH^-$$

ammonia

NH3NH_3$$NH_3$$

amide ion*

NH2NH_2$$NH_2$$

hydrogen

H2H_2$$H_2$$

hydride ion*

HH^-$$H^-$$

*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:

    • $$acid1 + base2 \rightleftharpoons acid2 + base1$$

  • Water can act as an acid:

    • $$base1 + acid2 \rightleftharpoons acid1 + base2$$

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

  • $$NO, NO2, CO2, SO2,and, and $$, and $$SO3$$ 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.


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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 (H2H_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., HNO3HNO3, H2SO4H2SO_4).

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

  • Examples:

    • CH3COOHCH3COOH: acetic acid, anion: CH3COOCH3COO^-, acetate

    • H2CO3H2CO3: carbonic acid, anion: CO32CO_3^{2-}, carbonate

    • HIO3HIO3: iodic acid, anion: IO3IO3^{-}, iodate

    • HClOHClO: hypochlorous acid, anion: ClOClO^{-}, hypochlorite

    • HClO2HClO2: chlorous acid, anion: ClO2ClO2^{-}, chlorite

    • HClO3HClO3: chloric acid, anion: ClO3ClO3^{-}, chlorate

    • HClO4HClO4: perchloric acid, anion: ClO4ClO4^{-}, perchlorate

    • HNO2HNO2: nitrous acid, anion: NO2NO2^{-}, nitrite

    • HNO3HNO3: nitric acid, anion: NO3NO3^{-}, nitrate

    • H3PO3H3PO3: phosphorous acid, anion: PO33PO_3^{3-}, phosphite

    • H3PO4H3PO4: phosphoric acid, anion: PO43PO_4^{3-}, phosphate

    • H2SO3H2SO3: sulfurous acid, anion: SO32SO_3^{2-}, sulfite

    • H2SO4H2SO4: sulfuric acid, anion: SO42SO_4^{2-}, sulfate

Common Industrial Acids

  • Sulfuric Acid (H2SO4H2SO4): The most commonly produced industrial chemical worldwide.

  • Nitric Acid (HNO3HNO_3)

  • Phosphoric Acid (H3PO4H3PO4)

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

  • Acetic Acid (CH3COOHCH_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+H^+) in aqueous solution.

  • Arrhenius Base: A substance that increases the concentration of hydroxide ions (OHOH^−) 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., HClO4HClO4, HCl, HNO3HNO3).

  • 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 (OHOH^−), which dissociate in water.

  • Ammonia (NH3NH_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: [H3O+]>107M>[OH][H_3O^+] > 10^{-7} M > [OH^-]

  • Neutral solution: [H3O+]=107M=[OH][H_3O^+] = 10^{-7} M = [OH^-]

  • Basic solution: [H3O+]<107M<[OH][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 OHOH^− 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., HClO4HClO4, HCl, HNO3HNO3).

    • Involves only one ionization step.

  • Polyprotic Acid: An acid that can donate more than one proton per molecule (e.g., H2SO4H2SO4, H3PO4H3PO4).

    • Involves multiple ionization steps.

    • Sulfuric acid solutions contain H3O+H_3O^+ ions.

  • Diprotic Acid: A polyprotic acid that can donate two protons per molecule (e.g., H2SO4H2SO4).

  • Triprotic Acid: A polyprotic acid that can donate three protons per molecule (e.g., H3PO4H3PO4).

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

Type

Acid

Base

Arrhenius

H+H^+ or H3O+H_3O^+ producer

OHOH^- producer

Brønsted-Lowry

Proton (H^) donor

Proton (H+H^+) acceptor

Lewis

Electron-pair acceptor

Electron-pair donor

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.

acid1+base2base1+acid2acid1 + base2 \rightleftharpoons base1 + acid2

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

Conjugate acid

Formula

Conjugate base

Formula

hydriodic acid*

HI

iodide ion

I-

perchloric acid*

HClO4HClO_4

perchlorate ion

ClO4ClO_4

hydrobromic acid*

HBr

bromide ion

Br-

hydrochloric acid*

HCl

chloride ion

Cl-

sulfuric acid*

H2SO4H2SO4

hydrogen sulfate ion

HSO4HSO_4

chloric acid*

HClO3HClO_3

chlorate ion

ClO3ClO_3

nitric acid*

HNO3HNO_3

nitrate ion

NO3NO_3

hydronium ion

H3O+H_3O^+

water

H2OH_2O

chlorous acid

HClO2HClO_2

chlorite ion

ClO2ClO_2

hydrogen sulfate ion

HSO4HSO_4

sulfate ion

SO4SO_4

phosphoric acid

H3PO4H3PO4

dihydrogen phosphate ion

H2PO4H2PO4

hydrofluoric acid

HF

fluoride ion

F-

acetic acid

CH3COOHCH_3COOH

acetate ion

CH3COOCH_3COO^-

carbonic acid

H2CO3H2CO3

hydrogen carbonate ion

HCO3HCO_3

hydrosulfuric acid

H2SH_2S

hydrosulfide ion

HSHS

dihydrogen phosphate ion

H2PO4H2PO4

hydrogen phosphate ion

HPO4HPO_4

hypochlorous acid

HClOHClO

hypochlorite ion

ClOClO^-

ammonium ion

NH4+NH_4^+

ammonia

NH3NH_3

hydrogen carbonate ion

HCO3HCO_3

carbonate ion

CO3CO_3

hydrogen phosphate ion

HPO4HPO_4

phosphate ion

PO4PO_4

water

H2OH_2O

hydroxide ion

OHOH^-

ammonia

NH3NH_3

amide ion*

NH2NH_2

hydrogen

H2H_2

hydride ion*

HH^-

*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:

    • acid1+base2acid2+base1acid1 + base2 \rightleftharpoons acid2 + base1

  • Water can act as an acid:

    • base1+acid2acid1+base2base1 + acid2 \rightleftharpoons acid1 + base2

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,NO2,CO2,SO2NO, NO2, CO2, SO2, and SO3SO3 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.