8: Acid-Base Equilibria
- Page ID
- 502458
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- Blood, Balance, and Biochemistry: Acid–Base Equilibria in the Body
- 8.1: Brønsted-Lowry Acids and Bases
- A Brønsted–Lowry acid is a compound that donates a proton (a hydrogen ion), while a Brønsted–Lowry base is a compound that accepts a proton. When an acid loses a proton, it forms its conjugate base; when a base gains a proton, it forms its conjugate acid. Amphiprotic species, like water, can act as either an acid or a base depending on the situation.
- 8.2: pH and pOH
- The acidity or basicity of a solution can be expressed using the pH and pOH scales: pH=−log[\(\ce{H3O+}\)] and pOH=−log[\(\ce{OH−}\)]. At 25 °C, the product of these ion concentrations is constant, so pH and pOH are related by the equation pH + pOH = 14. In a neutral solution, both values are equal to 7. For an acidic solution, pH is less than 7, and for a basic solution, pH is greater than 7.
- 8.3: Relative Strengths of Acids and Bases
- Acid and base strength is determined by their ionization constants (Ka for acids, Kb for bases), indicating the extent of ionization in water. Strong acids and bases ionize completely, while weak ones do so partially. When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid. A stronger acid has a weaker conjugate base, and vice versa. In acid-base reactions, equilibrium favors the formation of the weaker acid and base.
- 8.4: pH of Strong Acids and Strong Bases
- Strong acids and bases fully ionize in water. Therefore, their initial concentrations directly determine the concentrations of \(\ce{H3O+}\) or \(\ce{OH−}\), allowing straightforward calculation of pH or pOH.
- 8.5: pH of Weak Acids and Bases
- Weak acids and weak bases ionize only slightly in water. To calculate the pH of their solutions, we analyze equilibrium concentrations. Typically, the extent of ionization is so small that it is negligible compared to the initial concentration, simplifying calculations. Polyprotic acids have more than one ionizable proton, and their successive Ka values differ by several orders of magnitude. Similarly, polyprotic bases can accept more than one proton.
- 8.6: Acid-Base Properties of Salt Solutions
- Salt solutions may be acidic, basic, or neutral depending on how their ions react with water. If the cation is a weak acid, it donates protons and makes the solution acidic. If the anion is a weak base, it accepts protons and makes the solution basic. Inert ions, such as those from strong acids or bases, do not affect pH. For salts that contain both acidic and basic ions (or amphiprotic ions), comparing the Ka and Kb values indicates whether the solution will be acidic, basic, or neutral.
- 8.7: Molecular Structure and Acid-Base Behavior
- Acid strength is closely related to molecular structure. A weaker H-E bond makes it easier for the acid to donate a proton and a more polar E–H bond also favors ionization. For binary acids (HE), acid strength increases as the atomic size of E increases (which weakens the E-H bond) and as the electronegativity of E increases (which makes the E-H bond more polar). For oxoacids, acid strength increases with both the number of oxygen atoms and the electronegativity of the central atom.
- 8.8: Lewis Acids and Bases
- A Lewis acid is a species that can accept an electron pair, whereas a Lewis base has an electron pair available for donation to a Lewis acid. The Lewis definition includes Brønsted-Lowry acid-base reactions but also other important reactions in chemistry.
- 8.E: Acid-Base Equilibrium (Exercises)
- These problems are essential for solidifying your understanding of acid–base equilibrium. They are designed to help you practise applying concepts, rather than memorizing the steps for a specific type of question. Aim to work through the problems and arrive at the final answer independently. This active approach builds understanding.