Acids, Bases and Salts | Jamb Chemistry

Jamb chemistry key points on General characteristics, properties and uses of acids, bases and salts etc

General Characteristics, Properties, and Uses of Acids

1. Definition of Acids: Substances that release hydrogen ions $(H^+)$ or hydronium ions $(H_3O^+)$ in aqueous solutions.
2. Taste: Acids typically have a sour taste.
3. pH Range: Acids have a pH less than 7.
4. Electrical Conductivity: Acids conduct electricity in aqueous solutions due to ionization.
5. Reaction with Metals: Acids react with reactive metals to produce hydrogen gas and a salt.
   • Example: $2HCl + Zn \rightarrow ZnCl_2 + H_2$.
6. Reaction with Carbonates: Acids react with carbonates to release carbon dioxide gas.
   • Example: $HCl + CaCO_3 \rightarrow CaCl_2 + CO_2 + H_2O$.
7. Neutralization: Acids react with bases to form salt and water.
8. Common Acids: Hydrochloric acid $(HCl)$, sulfuric acid $(H_2SO_4)$, nitric acid $(HNO_3)$.
9. Organic Acids: Naturally occurring acids include ethanoic acid (vinegar), citric acid (citrus fruits), and tartaric acid (grapes).
10. Uses of Acids:
    • Hydrochloric acid: Cleaning metals.
    • Sulfuric acid: Battery production.
    • Citric acid: Food preservative.
11. Acidic Oxides: Nonmetal oxides like $CO_2$ and $SO_2$ dissolve in water to form acids.
12. Corrosiveness: Strong acids like $HCl$ and $H_2SO_4$ are highly corrosive.
13. Volatility: Acids like $HCl$ are volatile, while $H_2SO_4$ is non-volatile.
14. Indicator Reaction: Acids turn blue litmus paper red.
15. Environmental Role: Acid rain, formed by $SO_2$ and $NO_x$, impacts ecosystems.
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General Characteristics, Properties, and Uses of Bases

16. Definition of Bases: Substances that release hydroxide ions $(OH^-)$ in aqueous solutions or accept protons.
17. Taste: Bases have a bitter taste.
18. Texture: Bases feel slippery or soapy to touch.
19. pH Range: Bases have a pH greater than 7.
20. Electrical Conductivity: Bases conduct electricity in aqueous solutions.
21. Neutralization: Bases react with acids to form salt and water.
    • Example: $NaOH + HCl \rightarrow NaCl + H_2O$.
22. Reaction with Ammonium Salts: Bases release ammonia gas upon reaction.
    • Example: $NaOH + NH_4Cl \rightarrow NaCl + NH_3 + H_2O$.
23. Common Bases: Sodium hydroxide $(NaOH)$, calcium hydroxide $(Ca(OH)_2)$, ammonia $(NH_3)$.
24. Alkaline Oxides: Metal oxides like $Na_2O$ and $CaO$ react with water to form bases.
25. Uses of Bases:
    • Sodium hydroxide: Soap production.
    • Calcium hydroxide: Neutralizing acidic soil.
    • Ammonia: Fertilizer production.
26. Indicator Reaction: Bases turn red litmus paper blue.
27. Cleaning Properties: Bases like $NaOH$ are used in cleaning agents.
28. Industrial Uses: Bases neutralize acidic waste in chemical processes.
29. Storage: Strong bases like $NaOH$ are corrosive and require careful storage.
30. Biological Role: Bicarbonates in the human body maintain pH balance.
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General Characteristics, Properties, and Uses of Salts

31. Definition of Salts: Ionic compounds formed from the reaction of an acid and a base.
32. Composition: Made of cations from bases and anions from acids.
33. Neutralization Reaction:
    • $NaOH + HCl \rightarrow NaCl + H_2O$.
34. Solubility: Most salts are soluble in water, though exceptions like $BaSO_4$ exist.
35. Electrical Conductivity: Salts conduct electricity when dissolved or molten.
36. Examples of Salts:
    • Table salt $(NaCl)$.
    • Epsom salt $(MgSO_4 \cdot 7H_2O)$.
    • Alum $(KAl(SO_4)_2 \cdot 12H_2O)$.
37. Double Salts: Salts like alum contain two different cations.
38. Uses of Salts:
    • $NaCl$: Preservative and seasoning.
    • $NH_4NO_3$: Fertilizer.
    • $CuSO_4$: Pesticide and fungicide.
39. Industrial Role: Salts are raw materials in chemical manufacturing.
40. Water Softening: Salts like washing soda $(Na_2CO_3)$ remove water hardness.
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Acids/Base Indicators and Basicity of Acids

41. Indicators: Substances that change color to indicate the pH of a solution.
42. Natural Indicators: Litmus, red cabbage extract, turmeric.
43. Synthetic Indicators: Phenolphthalein (pink in base, colorless in acid) and methyl orange (red in acid, yellow in base).
44. Universal Indicators: Provide a full pH range with distinct colors.
45. Basicity of Acids: Refers to the number of hydrogen ions $(H^+)$ an acid can donate.
46. Monobasic Acid: Releases one $H^+$ ion per molecule, e.g., $HCl$.
47. Dibasic Acid: Releases two $H^+$ ions per molecule, e.g., $H_2SO_4$.
48. Tribasic Acid: Releases three $H^+$ ions per molecule, e.g., $H_3PO_4$.
49. Strength of Acids: Strong acids ionize completely in water $(HCl)$, while weak acids ionize partially $(CH_3COOH)$.
50. Indicator Selection: Depends on the acid-base reaction and pH range of the endpoint.
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Preparation and Classification of Salts

Types of Salts

51. Normal Salts: Formed when all hydrogen ions in an acid are replaced by metal ions, e.g., $NaCl$.
52. Acid Salts: Contain replaceable hydrogen atoms, e.g., $NaHSO_4$.
53. Basic Salts: Contain hydroxide ions, e.g., $Bi(OH)_2NO_3$.
54. Double Salts: Contain two different cations, e.g., alum $(KAl(SO_4)_2 \cdot 12H_2O)$.
55. Complex Salts: Contain complex ions, e.g., $[Cu(NH_3)_4]SO_4$.
    
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Preparation Methods

56. Neutralization: Acid + base → salt + water.
57. Precipitation: Two solutions react to form an insoluble salt.
    • Example: $AgNO_3 + NaCl \rightarrow AgCl \downarrow + NaNO_3$.
58. Direct Combination: Reaction of elements, e.g., $Fe + Cl_2 \rightarrow FeCl_3$.
59. Action of Acid on Metal:
    • Example: $Zn + 2HCl \rightarrow ZnCl_2 + H_2$.
60. Action of Acid on Metal Carbonate:
    • Example: $HCl + CaCO_3 \rightarrow CaCl_2 + CO_2 + H_2O$.
      
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Oxides and Trioxocarbonate

61. Oxides: Compounds of oxygen with another element.
62. Types of Oxides:
    • Acidic Oxides: $CO_2, SO_2$.
    • Basic Oxides: $Na_2O, CaO$.
    • Amphoteric Oxides: $Al_2O_3, ZnO$.
    • Neutral Oxides: $CO, N_2O$.
63. Trioxocarbonate (IV) Salts: Contain $CO_3^{2-}$, e.g., $Na_2CO_3, CaCO_3$.
64. Thermal Decomposition:
    • Example: $CaCO_3 \rightarrow CaO + CO_2$.
65. Industrial Use:
    • Limestone $(CaCO_3)$: Cement production.
    • Washing soda $(Na_2CO_3)$: Detergent.
      
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Acid/Base Types and Identification

81. Strong Acids: Fully dissociate, e.g., $HCl, H_2SO_4$.
82. Weak Acids: Partially dissociate, e.g., $CH_3COOH$.
83. Strong Bases: Fully dissociate, e.g., $NaOH, KOH$.
84. Weak Bases: Partially dissociate, e.g., $NH_3$.
85. Organic Acids: Contain carbon, e.g., citric acid.
86. Inorganic Acids: Do not contain carbon, e.g., $HCl$.
87. Industrial Acids: Used in manufacturing, e.g., sulfuric acid.
88. Neutral Bases: Exhibit neither acidic nor basic properties, e.g., $NH_4Cl$.
89. Reaction Tests:
    • Acids react with metals to produce hydrogen gas.
    • Bases neutralize acids.
90. Litmus Test: Identifies acids and bases quickly.
91. Titration: Measures concentration of acids and bases.
92. Buffer Solutions: Resist changes in pH, crucial in biological systems.
93. Amphoteric Substances: Behave as both acid and base, e.g., $Al_2O_3$.
94. Uses of Buffers: Maintain pH in blood and industrial processes.
95. Eco-Friendly Bases: Lime reduces soil acidity.
96. Corrosive Acids: Require safe handling.
97. Natural Indicators: Easily available and environmentally friendly.
98. Identification by Reaction: Testing reactivity helps identify substances.
99. Properties Differentiation: Acids corrode metals, while bases feel slippery.
100. Safety Measures: Proper storage of acids and bases prevents accidents.
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Jamb chemistry Key points on Qualitative comparison of the conductance of molar solutions of strong and weak acids and bases

Conductance of Molar Solutions of Strong and Weak Acids and Bases

1. Conductance measures a solution’s ability to conduct electricity.
2. Strong acids (e.g., HCl, H₂SO₄, HNO₃) exhibit high conductance due to complete dissociation.
3. Weak acids (e.g., CH₃COOH, HF) show lower conductance because they partially dissociate.
4. Strong bases (e.g., NaOH, KOH) also exhibit high conductance, dissociating fully into ions.
5. Weak bases (e.g., NH₃, CH₃NH₂) have lower conductance due to limited ionization.
6. Conductance of strong acids and bases increases linearly with concentration at low molarity.
7. Weak acids and bases show a nonlinear increase in conductance with concentration.
8. Dilution enhances the dissociation of weak acids and bases, increasing their conductance.
9. For equimolar solutions, strong acids/bases always have higher conductance than weak ones.
10. Conductance decreases for all solutions at very high concentrations due to ion-pairing effects.
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Relationship Between Degree of Dissociation and Strength of Acids and Bases

11. The degree of dissociation is the fraction of molecules dissociated into ions.
12. Strong acids and bases have a degree of dissociation close to 1 (nearly complete).
13. Weak acids and bases have a degree of dissociation less than 1 (partial dissociation).
14. The strength of an acid or base is proportional to its degree of dissociation.
15. The degree of dissociation increases as the solution is diluted (Ostwald’s dilution law).
16. Strong acids maintain nearly constant conductance even at higher concentrations.
17. Weak acids/bases show significant increases in conductance when diluted.
18. For weak acids, the equilibrium constant (Ka) correlates with the degree of dissociation.
19. Strong acids have high Ka values, whereas weak acids have low Ka values.
20. Similar relationships exist for strong and weak bases with their Kb values.
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Relationship Between Conductance and Amount of Ions Present

21. Conductance depends on the number of free ions in the solution.
22. Strong acids/bases generate more ions per unit molarity than weak acids/bases.
23. Weak acids like acetic acid dissociate partially, contributing fewer ions to conductance.
24. Polyprotic acids (e.g., H₂SO₄) dissociate in steps, influencing their conductance.
25. Conductance increases with ion mobility; H⁺ and OH⁻ ions have high mobility.
26. Larger ions or ions with greater hydration shells reduce conductance.
27. The presence of impurities or added salts can alter conductance by increasing ion concentration.
28. Solutions with only neutral molecules (non-electrolytes) exhibit negligible conductance.
29. Ionic strength affects conductance by altering the interaction between ions.
30. Highly dissociated solutions of strong acids and bases generate greater ionic current.
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Relationship Between Degree of Dissociation and Conductance

31. The degree of dissociation is directly proportional to the conductance of a solution.
32. At infinite dilution, conductance represents the maximum dissociation of the acid/base.
33. Weak acids have low initial conductance but approach stronger conductance upon dilution.
34. Conductance for strong acids/bases changes less with dilution due to full dissociation.
35. Weak acids exhibit a steep increase in conductance as they dissociate more in dilute solutions.
36. Strong acids maintain consistent dissociation and thus conductance at different dilutions.
37. The molar conductivity at infinite dilution (Λ°) depends on the degree of dissociation.
38. Weak acids and bases have Λ° values that increase significantly with dilution.
39. The relationship between conductance and dissociation is quantitatively described by Kohlrausch’s Law.
40. For weak electrolytes, conductance approaches the value of strong electrolytes only at high dilution.
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Comparative Trends and Practical Implications

41. The conductance of weak acids is sensitive to changes in temperature and pH.
42. Strong acids have a relatively consistent conductance across different conditions.
43. For a given acid or base, the molar conductivity is a function of both concentration and dissociation.
44. Ion pairing at high concentrations reduces effective conductance for all electrolytes.
45. Weak acids/bases are useful for buffering due to their partial dissociation and moderate conductance.
46. Strong electrolytes serve as effective conductors in applications requiring high ion availability.
47. Conductance measurements can help identify unknown acids or bases by comparing dissociation behaviors.
48. The conductivity apparatus provides qualitative and quantitative insights into ion dissociation.
49. Weak acids like CH₃COOH exhibit distinct conductance curves, helping visualize dissociation trends.
50. Understanding conductance and dissociation is vital for applications in electrochemistry, medicine, and environmental science.
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Jamb chemistry Key points on pH and pOH scale Acid/base titrations, Hydrolysis of salts

pH and pOH Scale

1. The pH scale measures the hydrogen ion concentration $[H^+]$ in a solution.
2. The pOH scale measures the hydroxide ion concentration $[OH^-]$.
3. pH is calculated as $pH = -\log[H^+]$.
4. pOH is calculated as $pOH = -\log[OH^-]$.
5. The relationship between pH and pOH is $pH + pOH = 14$ (at 25°C).
6. A pH < 7 indicates an acidic solution.
7. A pH > 7 indicates a basic solution.
8. A pH of 7 indicates a neutral solution.
9. In strong acids, $[H^+]$ is equal to the molarity of the acid.
10. For strong bases, $[OH^-]$ is equal to the molarity of the base.
11. For weak acids, $pH$ is calculated using $Ka$ and $[HA]$:
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    $pH = -\log\sqrt{Ka}[HA]$
12. For weak bases, $pOH$ is calculated using $Kb$ and $[B]$:
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    $pOH = -\log\sqrt{Kb}[B]$
13. The concentration of $[H^+]$ and $[OH^-]$ can be deduced from the pH or pOH.
14. Strong acid-base neutralization produces a pH = 7 (neutral solution).
15. pH calculations can determine the acidity of unknown solutions in experiments.
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Acid/Base Titrations

16. Acid-base titrations are used to determine the concentration of an unknown solution.
17. The equivalence point is where moles of acid = moles of base.
18. For strong acid-strong base titrations, the equivalence point pH = 7.
19. For strong acid-weak base titrations, the equivalence point pH < 7 (acidic).
20. For strong base-weak acid titrations, the equivalence point pH > 7 (basic).
21. The endpoint is the observed color change in the indicator used.
22. Indicators must be chosen based on the pH range of the equivalence point (e.g., phenolphthalein for basic endpoints).
23. The titration curve shows the relationship between pH and volume of titrant added.
24. The initial pH depends on the nature of the analyte (acidic or basic).
25. The steepness of the curve at the equivalence point is greater for strong acid/base systems.
26. Half-equivalence points help calculate the $Ka$ or $Kb$ of weak acids or bases.
27. Molarity of the unknown solution can be determined using the formula:
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    ${M}_1{V}_1 = {M}_2{V}_2$
28. Titrations can confirm the purity of acids or bases.
29. Acid-base titration is fundamental in pharmaceutical and industrial applications.
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Calculations Based on the Mole Concept

30. The mole is the SI unit representing $6.022 \times 10^{23}$ entities of a substance.
31. Molar mass (g/mol) is the mass of one mole of a substance.
32. The number of moles is calculated as:
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    ${Moles} = \frac{Mass (g)}{Molar Mass (g/mol)}$
33. For solutions, the number of moles is:
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    ${Moles} = {Molarity} \times {Volume (L)}$
34. Stoichiometry relates the moles of reactants and products in a chemical equation.
35. Limiting reactant is the one that is completely consumed in a reaction.
36. Excess reactant remains after the limiting reactant is used up.
37. Molar ratios from balanced equations are used to determine product quantities.
38. Theoretical yield is the maximum product amount calculated from stoichiometry.
39. Percentage yield compares actual yield to theoretical yield:
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    $Percentage Yield = \frac{Actual Yield}{Theoretical Yield} \times 100$
40. Avogadro’s number helps calculate the number of particles in a sample.
41. Gas volumes can be calculated using:
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    ${Moles} = \frac{Volume}{Molar Volume (22.4 L at STP)}$
42. Mole concept is central to determining empirical and molecular formulas.
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Balancing Equations for Hydrolysis of Salts

43. Salt hydrolysis involves reaction of a salt with water to produce acidic or basic solutions.
44. Hydrolysis depends on the salt's parent acid and base (strong or weak).
45. Strong acid + strong base salts (e.g., NaCl) do not hydrolyze, resulting in neutral solutions.
46. Strong acid + weak base salts (e.g., NH₄Cl) hydrolyze to produce acidic solutions.
47. Weak acid + strong base salts (e.g., CH₃COONa) hydrolyze to produce basic solutions.
48. Weak acid + weak base salts (e.g., NH₄CH₃COO) may produce neutral, acidic, or basic solutions based on ${Ka}$ and ${Kb}$.
49. The hydrolysis of $NH₄Cl$ is:
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    $NH₄^+ + H₂O \leftrightarrow NH₃ + H₃O^+$
50. The hydrolysis of $CH₃COONa$ is:
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    $CH₃COO^- + H₂O \leftrightarrow CH₃COOH + OH^-$
51. Balanced equations involve the formation of $H^+$ or $OH^-$ ions from salt hydrolysis.
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Properties of Resultant Solutions

52. Neutral salts result in ${pH} = 7$ (e.g., NaCl).
53. Acidic salts result in ${pH} < 7$ (e.g., $NH₄Cl$).
54. Basic salts result in $pH > 7$ (e.g., $Na₂CO₃$).
55. The $pH$ of the solution depends on the strength of the acid and base forming the salt.
56. Hydrolysis alters the ionization equilibrium of water, shifting $[H^+]$ or $[OH^-]$.
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Hydrolysis of Salts

57. Salt hydrolysis occurs when ions from the salt react with water.
58. Cations from weak bases (e.g., $NH₄^+$) act as acids, donating $H^+$.
59. Anions from weak acids (e.g., $CH₃COO^-$) act as bases, accepting $H^+$.
60. Strong acid + strong base salts do not hydrolyze (no $H^+$ or $OH^-$ ions formed).
61. Strong acid + weak base salts yield $H^+$, making the solution acidic.
62. Weak acid + strong base salts yield $OH^-$, making the solution basic.
63. Weak acid + weak base salts depend on $Ka$ and $Kb$:
    • $Ka > Kb$: Acidic solution.
    • $Ka < Kb$: Basic solution.
    • $Ka = Kb$: Neutral solution.
64. The extent of hydrolysis is quantified by the hydrolysis constant ($K_h$).
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Miscellaneous

65. Hydrolysis affects the buffer capacity of solutions.
66. Acidic salts can accelerate corrosion due to $H^+$ presence.
67. Basic salts can alter enzyme activity due to $OH^-$ presence.
68. pH changes during hydrolysis impact biological systems, such as blood pH.
69. Hydrolysis is a key reaction in food preservation and fermentation.
70. Conductance measurements help verify hydrolysis effects.
71. Titration curves of hydrolyzed salts provide insight into their acid/base nature.
72. Calculations involving hydrolysis constants relate $pH$ to $Ka/Kb$.

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