Skip to main content
Chemistry LibreTexts

8.S: Acids, Bases and pH (Summary)

  • Page ID
  • \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

    \( \newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\)

    ( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\)

    \( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

    \( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\)

    \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

    \( \newcommand{\Span}{\mathrm{span}}\)

    \( \newcommand{\id}{\mathrm{id}}\)

    \( \newcommand{\Span}{\mathrm{span}}\)

    \( \newcommand{\kernel}{\mathrm{null}\,}\)

    \( \newcommand{\range}{\mathrm{range}\,}\)

    \( \newcommand{\RealPart}{\mathrm{Re}}\)

    \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)

    \( \newcommand{\Argument}{\mathrm{Arg}}\)

    \( \newcommand{\norm}[1]{\| #1 \|}\)

    \( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)

    \( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\AA}{\unicode[.8,0]{x212B}}\)

    \( \newcommand{\vectorA}[1]{\vec{#1}}      % arrow\)

    \( \newcommand{\vectorAt}[1]{\vec{\text{#1}}}      % arrow\)

    \( \newcommand{\vectorB}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vectorC}[1]{\textbf{#1}} \)

    \( \newcommand{\vectorD}[1]{\overrightarrow{#1}} \)

    \( \newcommand{\vectorDt}[1]{\overrightarrow{\text{#1}}} \)

    \( \newcommand{\vectE}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash{\mathbf {#1}}}} \)

    \( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)

    \( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)

    • A bond that is formed from a hydrogen atom, which is part of a polar covalent bond (such as the O—H bond) to another, more electronegative atom (that has at least one unshared pair of electrons in its valence shell) is called a hydrogen bond. Hydrogen bonds are weak, partially covalent bonds. The bond dissociation energy of the O—H covalent bond 464 kJ/mole; the bond dissociation energy an O—H••••O hydrogen bond is about 21 kJ/mole.
    • Even though hydrogen bonds are relatively weak, the vast network of hydrogen bonds in water makes the energy significant, and hydrogen bonding is generally used to explain the high boiling point of water (100 ˚C), relative to molecules of similar mass that cannot hydrogen bond. The extra energy represents the energy required to break down the hydrogen bonding network.
    • Polar molecules, such as acids, strongly hydrogen bond to water. This hydrogen bonding not only stabilizes the molecular dipoles, but also weakens the H—A covalent bond (A represents the acid molecule). As a result of this weakening, the H—A bond in these acids stretches (the bond length increases) and then fully breaks. The hydrogen that was hydrogen-bonded to the water molecule now becomes fully bonded to the oxygen, forming the species H3O+ (the hydronium ion) and the acid now exists as an anion (A); this is the process of acid dissociation.
    • The hydronium ion and the acid anion that are formed in an acid dissociation can react to re-form the original acid. This represents a set of forward- and back-reactions that occur together on a very fast time-scale; this type of a set of reactions is called an equilibrium and a double arrow is used in the chemical reaction to show this. This type of reaction is referred to as an acid dissociation equilibrium. HA (aq) + H2O ⇄ H3O+ (aq) + A (aq)
    • For any equilibrium, an equilibrium constant can be written that describes whether the products or the reactants will be the predominant species in solution. For the dissociation of the simple acid, HA, the equilibrium constant, Ka, is simply given by the ratio of the concentrations of the products and the reactants, remembering that the molarity of the solutes have been used to approximate their activity, and that solvents, such as water, have an activity of 1. Thus for the ionization of HA; \[K_{a}=\frac{[H_{3}O^{+}]{[A^{-}]}}{[HA]} \nonumber \]
    • According to the Brønsted Acid-Base Theory, any substance that ionizes in water to form hydronium ions (a proton donor) is called an acid; any substance that accepts a proton from a hydronium ion is a base. In an acid-base equilibrium, the conjugate acid is defined as the species that donates a hydrogen (a proton) in the forward reaction, and the conjugate base is the species that accepts a hydrogen (a proton) the reverse reaction. Thus for the ionization of HCl, HCl is the conjugate acid and Cl is the conjugate base. HCl (aq) + H2O ⇄ H3O+ (aq) + Cl (aq)
    • Metal hydroxides, such as NaOH, dissolve in water to form metal cations and hydroxide anion. Hydroxide anion is a strong Brønsted base and, therefore, hydroxide anion accepts a proton from the hydronium ion to form two moles of water. The reaction of a Brønsted acid with a Brønsted base to form water is the process of neutralization.
    • Just like water can promote the ionization of acids, water can also promote the ionization of itself. This equilibrium process occurs very rapidly in pure water and any sample of pure water will always contain a small concentration of hydronium and hydroxide ions. In pure water, at 25 oC, the concentration of hydronium ions ([H3O+]) and hydroxide ions ([HO]) will both be equal to exactly 1 × 10-7 M. This is referred to as the autoprotolysis of water.
    • The equilibrium for the autoprotolysis of water is defined as Kw, according to the equation shown below:

    \[K_{W}=[H_{3}O^{+}]{[HO^{-}]} \nonumber \]

    and at neutrality, [H3O+] and [HO] are both 1 × 10-7 M, making the value of Kw

    \[K_{W}=[1\times 10^{-7}]{[1\times 10^{-7}]}=1\times 10^{-14} \nonumber \]

    • Based on the autoprotolysis equilibrium, acidic, basic and neutral solutions can be defined as:
      • A solution is acidic if [H3O+] > 1 × 10-7 M.
      • A solution is basic if [H3O+] < 1 × 10-7 M.
      • A solution is neutral if [H3O+] = 1 × 10-7 M.
    • A pH value is simply the negative of the logarithm of the hydronium ion concentration (-log[H3O+]).
    • Remember that a logarithm consists of two sets of numbers; the digits to the left of the decimal point (the characteristic) reflect the integral power of 10, and are not included when you count significant figures. The numbers after the decimal (the mantissa) have the same significance as your experimental number, thus a logarithm of 4.15482 represents five significant figures.
    • In an acid-base titration a solution of acid or base with a known concentration is slowly add to an acid or base solution with an unknown concentration, using a volumetric burette, to until neutrality has been achieved. Typically an indicator or a pH meter is used to signify neutrality.
    • At neutrality, the volume and concentration of the reactant you have added is known, which means that you can calculate the number of moles that you added (remember, concentration × volume = moles). Based on the stoichiometry of your neutralization reaction, you then know how many moles of acid or base were in the unknown sample.

    This page titled 8.S: Acids, Bases and pH (Summary) is shared under a CC BY-SA 4.0 license and was authored, remixed, and/or curated by Paul R. Young ( via source content that was edited to the style and standards of the LibreTexts platform; a detailed edit history is available upon request.