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.
Binary Acid: An acid containing only two different elements: hydrogen and a more electronegative element (e.g., HF, HCl, HBr, HI).
Binary Acid Nomenclature:
Begins with the prefix hydro-
Followed by the root name of the second element.
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: acetic acid, anion: CH3COO^-, acetate
H2CO3: carbonic acid, anion: CO_3^{2-}, carbonate
HIO3: iodic acid, anion: IO3^{-}, iodate
HClO: hypochlorous acid, anion: ClO^{-}, hypochlorite
HClO2: chlorous acid, anion: ClO2^{-}, chlorite
HClO3: chloric acid, anion: ClO3^{-}, chlorate
HClO4: perchloric acid, anion: ClO4^{-}, perchlorate
HNO2: nitrous acid, anion: NO2^{-}, nitrite
HNO3: nitric acid, anion: NO3^{-}, nitrate
H3PO3: phosphorous acid, anion: PO_3^{3-}, phosphite
H3PO4: phosphoric acid, anion: PO_4^{3-}, phosphate
H2SO3: sulfurous acid, anion: SO_3^{2-}, sulfite
H2SO4: sulfuric acid, anion: SO_4^{2-}, sulfate
Sulfuric Acid (H2SO4): The most commonly produced industrial chemical worldwide.
Nitric Acid (HNO_3)
Phosphoric Acid (H3PO4)
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.
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 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.
Strong Acid: Ionizes completely in aqueous solution; a strong electrolyte (e.g., HClO4, HCl, HNO3).
Weak Acid: Releases few hydrogen ions in aqueous solution (e.g., HCN, organic acids such as acetic acid).
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.
The strength of a base depends on the extent to which the base dissociates; strong bases are strong electrolytes.
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 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 Acid: An acid that can donate only one proton (hydrogen ion) per molecule (e.g., HClO4, 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 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 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.
Type | Acid | Base |
---|---|---|
Arrhenius | H^+ or H_3O^+ producer | OH^- producer |
Brønsted-Lowry | Proton (H^) donor | Proton (H^+) acceptor |
Lewis | Electron-pair acceptor | Electron-pair donor |
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
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.
Conjugate acid | Formula | Conjugate base | Formula |
---|---|---|---|
hydriodic acid* | HI | iodide ion | I- |
perchloric acid* | HClO_4 | perchlorate ion | ClO_4 |
hydrobromic acid* | HBr | bromide ion | Br- |
hydrochloric acid* | HCl | chloride ion | Cl- |
sulfuric acid* | H2SO4 | hydrogen sulfate ion | HSO_4 |
chloric acid* | HClO_3 | chlorate ion | ClO_3 |
nitric acid* | HNO_3 | nitrate ion | NO_3 |
hydronium ion | H_3O^+ | water | H_2O |
chlorous acid | HClO_2 | chlorite ion | ClO_2 |
hydrogen sulfate ion | HSO_4 | sulfate ion | SO_4 |
phosphoric acid | H3PO4 | dihydrogen phosphate ion | H2PO4 |
hydrofluoric acid | HF | fluoride ion | F- |
acetic acid | CH_3COOH | acetate ion | CH_3COO^- |
carbonic acid | H2CO3 | hydrogen carbonate ion | HCO_3 |
hydrosulfuric acid | H_2S | hydrosulfide ion | HS |
dihydrogen phosphate ion | H2PO4 | hydrogen phosphate ion | HPO_4 |
hypochlorous acid | HClO | hypochlorite ion | ClO^- |
ammonium ion | NH_4^+ | ammonia | NH_3 |
hydrogen carbonate ion | HCO_3 | carbonate ion | CO_3 |
hydrogen phosphate ion | HPO_4 | phosphate ion | PO_4 |
water | H_2O | hydroxide ion | OH^- |
ammonia | NH_3 | amide ion* | NH_2 |
hydrogen | H_2 | hydride ion* | H^- |
*Strong acids
† Strong bases
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
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.
Acidity increases as more oxygen atoms are present.
Hypochlorous acid
Chlorous acid
Chloric acid
Perchloric acid
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.
NO, NO2, CO2, SO2, 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.