🧶Inorganic Chemistry I Unit 6 – Acid-Base Concepts in Inorganic Chem

Key Concepts and Definitions

• Acids produce hydrogen ions ($H^+$) in aqueous solutions
• Bases produce hydroxide ions ($OH^-$) in aqueous solutions
• Amphoteric substances can act as either an acid or a base depending on the reaction conditions (water)
• Conjugate acid-base pairs consist of two species that differ by a proton ($H^+$) ($NH_4^+$ and $NH_3$)
  • The conjugate base is formed by removing a proton from the acid
  • The conjugate acid is formed by adding a proton to the base
• Neutralization reactions involve the combination of an acid and a base to form water and a salt
• Autoionization of water refers to the self-ionization of water molecules into $H^+$ and $OH^-$ ions
  • The ion product constant of water ($K_w$) is equal to $[H^+][OH^-] = 1.0 \times 10^{-14}$ at 25°C

Acid-Base Theories

• Arrhenius theory defines acids as substances that dissociate in water to produce $H^+$ ions and bases as substances that dissociate to produce $OH^-$ ions
• Brønsted-Lowry theory defines acids as proton donors and bases as proton acceptors
  • Focuses on the transfer of protons ($H^+$) between species
  • Allows for the identification of conjugate acid-base pairs
• Lewis theory defines acids as electron pair acceptors and bases as electron pair donors
  • Broadens the definition of acids and bases beyond proton transfer
  • Includes species that do not contain protons ($BF_3$, $AlCl_3$)
• Solvent system definition considers the role of the solvent in acid-base reactions
  • In non-aqueous solvents, the solvent can act as an acid or base (liquid ammonia, glacial acetic acid)
• Usanovich theory combines the Lewis and Brønsted-Lowry theories
  • Defines acids as species that accept negative species or donate positive species
  • Defines bases as species that accept positive species or donate negative species

Types of Acids and Bases

• Strong acids completely dissociate in aqueous solutions ($HCl$, $HNO_3$, $H_2SO_4$)
  • Have a large acid dissociation constant ($K_a$)
  • Produce a high concentration of $H^+$ ions in solution
• Weak acids partially dissociate in aqueous solutions ($CH_3COOH$, $HF$, $H_2CO_3$)
  • Have a small acid dissociation constant ($K_a$)
  • Establish an equilibrium between the acid and its dissociated ions
• Strong bases completely dissociate in aqueous solutions ($NaOH$, $KOH$)
  • Produce a high concentration of $OH^-$ ions in solution
• Weak bases partially dissociate in aqueous solutions ($NH_3$, $CH_3NH_2$)
  • Establish an equilibrium between the base and its dissociated ions
• Polyprotic acids can donate multiple protons ($H_3PO_4$, $H_2SO_4$)
  • Each successive proton dissociation has a smaller acid dissociation constant ($K_{a1} > K_{a2} > K_{a3}$)
• Organic acids contain a carboxyl group ($-COOH$) and can act as weak acids (benzoic acid, citric acid)

Acid-Base Reactions

• Neutralization reactions involve the combination of an acid and a base to form water and a salt
  • The net ionic equation for a strong acid-strong base reaction is: $H^+(aq) + OH^-(aq) \rightarrow H_2O(l)$
  • The pH at the equivalence point of a strong acid-strong base titration is 7
• Acid-base titrations are used to determine the concentration of an unknown acid or base solution
  • Involves the gradual addition of a titrant (acid or base of known concentration) to the analyte (unknown solution)
  • The equivalence point is reached when the moles of titrant equal the moles of analyte
• Hydrolysis reactions involve the reaction of a salt with water to produce an acidic or basic solution
  • Salts of weak acids and strong bases produce basic solutions (sodium acetate)
  • Salts of strong acids and weak bases produce acidic solutions (ammonium chloride)
• Acid-base indicators are weak acids or bases that change color depending on the pH of the solution
  • The color change occurs over a specific pH range, allowing for the visual determination of pH
  • Common indicators include phenolphthalein, methyl orange, and bromothymol blue
• Lewis acid-base reactions involve the formation of a coordinate covalent bond
  • The Lewis acid accepts an electron pair from the Lewis base ($BF_3$ + $NH_3$ $\rightarrow$ $F_3B-NH_3$)

pH and pOH

• pH is a logarithmic scale that measures the acidity of a solution
  • Defined as the negative logarithm of the hydrogen ion concentration: $pH = -\log[H^+]$
  • A pH less than 7 indicates an acidic solution, while a pH greater than 7 indicates a basic solution
• pOH is a logarithmic scale that measures the basicity of a solution
  • Defined as the negative logarithm of the hydroxide ion concentration: $pOH = -\log[OH^-]$
  • A pOH less than 7 indicates a basic solution, while a pOH greater than 7 indicates an acidic solution
• The relationship between pH and pOH is given by: $pH + pOH = 14$
• The pH of a solution can be calculated from the concentration of $H^+$ ions
  • For a strong acid: $pH = -\log[H^+] = -\log(C_a)$, where $C_a$ is the concentration of the acid
  • For a weak acid: $pH = -\log[H^+] = -\log(\sqrt{K_aC_a})$, where $K_a$ is the acid dissociation constant
• The pOH of a solution can be calculated from the concentration of $OH^-$ ions
  • For a strong base: $pOH = -\log[OH^-] = -\log(C_b)$, where $C_b$ is the concentration of the base
  • For a weak base: $pOH = -\log[OH^-] = -\log(\sqrt{K_bC_b})$, where $K_b$ is the base dissociation constant

Buffers and Buffer Solutions

• Buffers are solutions that resist changes in pH when small amounts of acid or base are added
  • Consist of a weak acid and its conjugate base, or a weak base and its conjugate acid
  • Maintain a relatively constant pH through the consumption of added $H^+$ or $OH^-$ ions
• The Henderson-Hasselbalch equation relates the pH of a buffer solution to the concentrations of the acid and its conjugate base
  • For an acidic buffer: $pH = pK_a + \log\frac{[A^-]}{[HA]}$, where $pK_a = -\log K_a$
  • For a basic buffer: $pOH = pK_b + \log\frac{[B]}{[BH^+]}$, where $pK_b = -\log K_b$
• The buffer capacity is a measure of the ability of a buffer to resist changes in pH
  • Depends on the concentrations of the acid and its conjugate base, as well as the ratio of these concentrations
  • A higher buffer capacity indicates a greater ability to maintain a stable pH
• Buffers are important in biological systems to maintain the pH of bodily fluids (blood, cytoplasm)
  • The bicarbonate buffer system ($H_2CO_3$/$HCO_3^-$) is a major buffer in blood
  • Phosphate buffers ($H_2PO_4^-$/$HPO_4^{2-}$) are important in intracellular fluids
• Buffers are also used in chemical processes to control the pH of reactions
  • Maintain optimal conditions for enzyme activity, product formation, or separation processes

Applications in Inorganic Chemistry

• Acid-base reactions are involved in the synthesis and purification of inorganic compounds
  • Precipitation reactions rely on the formation of insoluble salts from the reaction of acids and bases ($BaCl_2 + H_2SO_4 \rightarrow BaSO_4 \downarrow + 2HCl$)
  • Acid-base extraction is used to separate compounds based on their solubility in acidic or basic solutions
• Acid-base properties influence the solubility and reactivity of inorganic compounds
  • Amphoteric oxides ($Al_2O_3$, $ZnO$) can react with both acids and bases
  • Basic oxides ($Na_2O$, $CaO$) react with acids to form salts and water
• Acid-base reactions are involved in the corrosion of metals
  • Acidic environments promote the oxidation of metals and the formation of soluble metal ions
  • Basic environments can lead to the formation of protective oxide layers on metal surfaces
• Acid-base chemistry is important in environmental and industrial processes
  • The pH of soil and water affects the availability of nutrients and the survival of organisms
  • Acid rain, caused by the dissolution of atmospheric pollutants ($SO_2$, $NO_x$), can harm ecosystems
  • The pH control of industrial wastewater is necessary to meet environmental regulations and prevent equipment damage
• Acid-base reactions are used in analytical chemistry techniques
  • Acid-base titrations are used to determine the concentration of acids or bases in solutions
  • pH indicators are used to visually detect the endpoint of titrations or to estimate the pH of solutions

Problem-Solving and Calculations

• Calculating the pH or pOH of a solution from the concentration of $H^+$ or $OH^-$ ions
  • Example: Calculate the pH of a 0.025 M $HCl$ solution
    • $pH = -\log[H^+] = -\log(0.025) = 1.60$
• Calculating the concentration of $H^+$ or $OH^-$ ions from the pH or pOH of a solution
  • Example: Calculate the $[OH^-]$ in a solution with a pH of 9.5
    • $pOH = 14 - pH = 14 - 9.5 = 4.5$
    • $[OH^-] = 10^{-pOH} = 10^{-4.5} = 3.2 \times 10^{-5} M$
• Calculating the pH of a weak acid or weak base solution using the acid or base dissociation constant
  • Example: Calculate the pH of a 0.1 M $CH_3COOH$ solution ($K_a = 1.8 \times 10^{-5}$)
    • $pH = -\log[H^+] = -\log(\sqrt{K_aC_a}) = -\log(\sqrt{(1.8 \times 10^{-5})(0.1)}) = 2.87$
• Calculating the pH of a buffer solution using the Henderson-Hasselbalch equation
  • Example: Calculate the pH of a buffer containing 0.2 M $CH_3COOH$ and 0.1 M $CH_3COONa$ ($K_a = 1.8 \times 10^{-5}$)
    • $pH = pK_a + \log\frac{[A^-]}{[HA]} = -\log(1.8 \times 10^{-5}) + \log\frac{0.1}{0.2} = 4.74 + (-0.30) = 4.44$
• Solving acid-base titration problems to determine the concentration of an unknown solution
  • Example: A 25.0 mL sample of $HCl$ is titrated with 0.100 M $NaOH$. The equivalence point is reached after 20.0 mL of $NaOH$ is added. Calculate the concentration of the $HCl$ solution.
    • At the equivalence point: $n(HCl) = n(NaOH)$
    • $M(HCl) \times V(HCl) = M(NaOH) \times V(NaOH)$
    • $M(HCl) = \frac{M(NaOH) \times V(NaOH)}{V(HCl)} = \frac{(0.100 M)(20.0 mL)}{25.0 mL} = 0.0800 M$