Chemical Reactions and Thermochemistry ✏ SAT Chemistry

Rucete ✏ SAT Chemistry In a Nutshell

8. Chemical Reactions and Thermochemistry

This chapter explains the four basic types of chemical reactions and introduces key thermochemistry concepts such as enthalpy, entropy, and Hess’s Law. It also teaches how to predict whether reactions occur and how to calculate energy changes during chemical processes.

Four Types of Chemical Reactions

1. Combination (Synthesis)

• Two or more elements or compounds combine to form one compound.
  • Example: Zn + S → ZnS (ΔH = −202.7 kJ)
  • If the heat of formation is large and negative, the reaction is likely spontaneous and exothermic.

2. Decomposition (Analysis)

• One compound breaks down into simpler substances.
  • Example: 2HgO → 2Hg + O₂ (ΔH = +90.8 kJ)
  • A low negative or positive ΔH means the compound can decompose more easily.

3. Single Replacement

• One element replaces another in a compound.
  • Example: Zn + 2HCl → ZnCl₂ + H₂
  • Can be predicted using ΔH values or the activity series:
    • Elements higher in the series replace lower ones.

4. Double Replacement

• Two ionic compounds exchange partners.
  • Occurs if:
    1. An insoluble precipitate forms
    2. A nonionizing substance (e.g., water) forms
    3. A gas is released
  • Example:
    AgNO₃ + KCl → AgCl (s) + KNO₃
    HCl + NaOH → H₂O + NaCl
    CaCO₃ + HCl → CaCl₂ + H₂O + CO₂ (gas)

Hydrolysis Reactions

• Hydrolysis is the reverse of neutralization. A salt reacts with water to form an acid and/or a base.
• The pH of the resulting solution depends on the strengths of the parent acid and base:
  • Strong acid + strong base → neutral solution (e.g., NaCl + H₂O)
  • Strong base + weak acid → basic solution (e.g., Na₂CO₃ + H₂O)
  • Strong acid + weak base → acidic solution (e.g., ZnCl₂ + H₂O)
  • Weak acid + weak base → neutral or slightly shifted, depending on relative strength (e.g., (NH₄)₂CO₃ + H₂O)

Entropy (ΔS)

• Entropy is the measure of disorder or randomness in a system.
• Systems tend to move toward higher entropy (more disorder).
• Gases have higher entropy than liquids or solids.
• Reactions that increase the number of gas molecules or produce a mixture often increase entropy.
• Entropy is a key driving force of spontaneous reactions, alongside enthalpy.

Enthalpy (ΔH)

• Enthalpy = heat content of a system at constant pressure.
• ΔH < 0 → exothermic (releases heat)
  ΔH > 0 → endothermic (absorbs heat)
• Standard conditions: 25°C and 1 atm → values given as ΔH⁰
• Molar heat of formation (ΔH⁰f) = enthalpy change when 1 mole of a compound is formed from its elements in standard states.
• Example:
  H₂(g) + ½O₂(g) → H₂O(g)  ΔH⁰ = −241.8 kJ/mol

Heat of Combustion

• Heat of combustion (ΔHc) = heat released when 1 mole of substance is burned in oxygen.
• Differs from heat of formation, which refers to product formation.
• Example: Combustion of CH₄ produces large amounts of heat and is exothermic.

Hess’s Law

• Hess’s Law of Heat Summation:
  If a reaction is the sum of two or more other reactions, the ΔH of the overall reaction is the sum of the ΔH values of those reactions.
• Based on the First Law of Thermodynamics (energy cannot be created or destroyed).
• Steps:
  • Equations can be reversed (change the sign of ΔH).
  • Equations can be multiplied or divided (adjust ΔH accordingly).
  • Add all modified reactions to get the target reaction.
• Example:
  Want to find ΔH for H₂O(l) → H₂O(g)
  Use:
  H₂(g) + ½O₂(g) → H₂O(l)  ΔH = −285.8 kJ
  H₂(g) + ½O₂(g) → H₂O(g)  ΔH = −241.8 kJ
  Subtract (reverse the first equation and add):
  H₂O(l) → H₂O(g)  ΔH = +44.0 kJ

Alternate Method: Using ΔH⁰f Values

• You can also calculate the enthalpy change of a reaction using this formula:
  ΔHreaction = ΣΔHf(products) − ΣΔHf(reactants)
• Example:
  4NH₃(g) + 5O₂(g) → 6H₂O(g) + 4NO(g)
  Given:
  ΔHf(H₂O) = −241.8 kJ/mol
  ΔHf(NO) = +90.3 kJ/mol
  ΔHf(NH₃) = −46.2 kJ/mol
  ΔHf(O₂) = 0 (standard element)
  
  Products: (6 × −241.8) + (4 × 90.3) = −1089.6 kJ
  Reactants: (4 × −46.2) + 0 = −184.8 kJ
  ΔHreaction = −1089.6 − (−184.8) = −904.8 kJ

In a Nutshell

This chapter explains how to classify and predict chemical reactions—synthesis, decomposition, single and double replacement—and introduces the energy changes involved. You learned how to identify reaction spontaneity using ΔH and entropy, calculate enthalpy changes using ΔHf values, and apply Hess’s Law to multi-step reactions. These skills are essential for understanding reaction behavior and solving thermochemical problems on the SAT Chemistry test.