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7: Acid-Base Equilibria

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  • This chapter will illustrate the chemistry of acid-base reactions and equilibria, and provide you with tools for quantifying the concentrations of acids and bases in solutions.

    Unit 7 Objectives

    By the end of this unit, you will be able to:

    • Classify substances as acidic or basic based on their pH or chemical formula.
    • List the twelve common acids and bases, providing their name, formula, and whether they are weak or strong. 
    • Write equations for acid and base ionization reactions
    • Use the ion-product constant for water to calculate hydronium and hydroxide ion concentrations
    • Explain the characterization of aqueous solutions as acidic, basic, or neutral
    • Calculate pH.
    • Assess the relative strengths of acids and bases according to their ionization constants.
    • Write balanced chemical equations for neutralization reactions and determine if the resulting solution will be acidic, basic, or neutral. 
    • Determine if an acid is monoprotic, diprotic, or triprotic given the chemical formula.  
    • Describe the composition and function of acid–base buffers.

    • 7.1: Acids and Bases
      Compounds that donate a proton (a hydrogen ion) to another compound is called a Brønsted-Lowry acid. The compound that accepts the proton is called a Brønsted-Lowry base. The species remaining after a Brønsted-Lowry acid has lost a proton is the conjugate base of the acid. The species formed when a Brønsted-Lowry base gains a proton is the conjugate acid of the base. Amphiprotic species can act as both proton donors and proton acceptors. Water is the most important amphiprotic species.
    • 7.2: pH and pOH
      The concentration of hydronium ion in a solution of an acid in water is greater than \( 1.0 \times 10^{-7}\; M\) at 25 °C. The concentration of hydroxide ion in a solution of a base in water is greater than \( 1.0 \times 10^{-7}\; M\) at 25 °C. The concentration of H3O+ in a solution can be expressed as the pH of the solution; \(\ce{pH} = -\log \ce{H3O+}\). The concentration of OH− can be expressed as the pOH of the solution: \(\ce{pOH} = -\log[\ce{OH-}]\). In pure water, pH = 7 and pOH = 7.
    • 7.3: Relative Strengths of Acids and Bases
      The strengths of Brønsted-Lowry acids and bases in aqueous solutions can be determined by their acid or base ionization constants. Stronger acids form weaker conjugate bases, and weaker acids form stronger conjugate bases. Thus strong acids are completely ionized in aqueous solution because their conjugate bases are weaker bases than water. Weak acids are only partially ionized because their conjugate bases are compete successfully with water for possession of protons.
    • 7.4: Acid-Base Neutralization
      The characteristic properties of aqueous solutions of Brønsted-Lowry acids are due to the presence of hydronium ions; those of aqueous solutions of Brønsted-Lowry bases are due to the presence of hydroxide ions. The neutralization that occurs when aqueous solutions of acids and bases are combined results from the reaction of the hydronium and hydroxide ions to form water. Some salts formed in neutralization reactions may make the product solutions slightly acidic or slightly basic.
    • 7.5: Polyprotic Acids
      An acid that contains more than one ionizable proton is a polyprotic acid. The protons of these acids ionize in steps. The differences in the acid ionization constants for the successive ionizations of the protons in a polyprotic acid usually vary by roughly five orders of magnitude. As long as the difference between the successive values of Ka of the acid is greater than about a factor of 20, it is appropriate to break down the calculations of the concentrations sequentially.
    • 7.6: Buffers
      A solution containing a mixture of an acid and its conjugate base, or of a base and its conjugate acid, is called a buffer solution. Unlike in the case of an acid, base, or salt solution, the hydronium ion concentration of a buffer solution does not change greatly when a small amount of acid or base is added to the buffer solution. The base (or acid) in the buffer reacts with the added acid (or base).
    • 7.7: Unit 7 Practice Problems


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