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

Acid-base concepts form the foundation of inorganic chemistry, explaining how substances interact and react. These principles cover everything from simple proton transfers to complex buffer systems, providing a framework for understanding chemical behavior in various environments. From Arrhenius to Lewis theories, acid-base concepts have evolved to encompass a wide range of chemical interactions. Understanding pH, neutralization reactions, and buffer solutions is crucial for predicting and controlling chemical processes in both laboratory and real-world applications.

Study Guides for Unit 6

6.1

Brønsted-Lowry and Lewis Acid-Base Theories

4 min read

6.2

Hard-Soft Acid-Base (HSAB) Theory

4 min read

6.3

Acid-Base Properties of Oxides and Aqueous Solutions

3 min read

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 ( $H^+$ $H^{+}$ ) ( $NH_4^+$ $N H_{4}^{+}$ and $NH_3$ $N H_{3}$ )
- 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 $H^+$ $H^{+}$ and $OH^-$ $O H^{-}$ 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 ( $HCl$ $H Cl$ , $HNO_3$ $H N O_{3}$ , $H_2SO_4$ $H_{2} S O_{4}$ )
- Have a large acid dissociation constant ( $K_a$ )
- Produce a high concentration of $H^+$ ions in solution
Weak acids partially dissociate in aqueous solutions ( $CH_3COOH$ $C H_{3} COO H$ , $HF$ $H F$ , $H_2CO_3$ $H_{2} C O_{3}$ )
- 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 ( $NaOH$ $N a O H$ , $KOH$ $K O H$ )
- Produce a high concentration of $OH^-$ ions in solution
Weak bases partially dissociate in aqueous solutions ( $NH_3$ $N H_{3}$ , $CH_3NH_2$ $C H_{3} N H_{2}$ )
- Establish an equilibrium between the base and its dissociated ions
Polyprotic acids can donate multiple protons ( $H_3PO_4$ $H_{3} P O_{4}$ , $H_2SO_4$ $H_{2} S O_{4}$ )
- 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 $H^+$ $H^{+}$ 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 $OH^-$ $O H^{-}$ 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 $H^+$ $H^{+}$ or $OH^-$ $O H^{-}$ ions
- Example: Calculate the pH of a 0.025 M $HCl$ $H Cl$ solution
  - $pH = -\log[H^+] = -\log(0.025) = 1.60$
Calculating the concentration of $H^+$ $H^{+}$ or $OH^-$ $O H^{-}$ ions from the pH or pOH of a solution
- Example: Calculate the $[OH^-]$ $[O H^{-}]$ 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 $CH_3COOH$ $C H_{3} COO H$ solution ( $K_a = 1.8 \times 10^{-5}$ $K_{a} = 1.8 \times 1 0^{- 5}$ )
  - $pH = -\log[H^+] = -\log(\sqrt{K_aC_a}) = -\log(\sqrt{(1.8 \times 10^{-5})(0.1)}) = 2.87$
Calculating the pH of a buffer solution using the Henderson-Hasselbalch equation
- Example: Calculate the pH of a buffer containing 0.2 M $CH_3COOH$ $C H_{3} COO H$ and 0.1 M $CH_3COONa$ $C H_{3} COON a$ ( $K_a = 1.8 \times 10^{-5}$ $K_{a} = 1.8 \times 1 0^{- 5}$ )
  - $pH = pK_a + \log\frac{[A^-]}{[HA]} = -\log(1.8 \times 10^{-5}) + \log\frac{0.1}{0.2} = 4.74 + (-0.30) = 4.44$
Solving acid-base titration problems to determine the concentration of an unknown solution
- Example: A 25.0 mL sample of $HCl$ $H Cl$ is titrated with 0.100 M $NaOH$ $N a O H$ . The equivalence point is reached after 20.0 mL of $NaOH$ $N a O H$ is added. Calculate the concentration of the $HCl$ $H Cl$ solution.
  - At the equivalence point: $n(HCl) = n(NaOH)$
  - $M(HCl) \times V(HCl) = M(NaOH) \times V(NaOH)$
  - $M(HCl) = \frac{M(NaOH) \times V(NaOH)}{V(HCl)} = \frac{(0.100 M)(20.0 mL)}{25.0 mL} = 0.0800 M$