CHEMICAL ENERGETICS - DETAILED STUDY NOTES

CHEMICAL ENERGETICS - DETAILED STUDY NOTES

6.1 ENTHALPY

1. THE COLLISION THEORY

  • Chemical energy of any substance includes:

    • Kinetic energy: Energy due to motion.

    • Potential energy: Energy stored in bonds.

  • The Collision Theory posits:

    1. A chemical reaction occurs only when particles collide.

    2. Particles must collide with sufficient energy, termed activation energy (Ea), to react.

  • Activation Energy (Ea):

    • Defined as the minimum energy required for reactants to collide successfully and form products.

    • Successful collision allows the reactants to overcome the energy barrier and transform into products.

  • Energy Change During Chemical Reactions:

    1. Bonds in the reactant molecules break, which requires energy.

    2. Bonds in the product molecules form, which releases energy.

    3. The predominant energy transfer during a reaction is thermal energy (heat).

    4. Heat transfer occurs between the system (reactants & products) and the surroundings (external environment such as a container, solvent, air, etc.).

    5. Enthalpy change (ΔH) is defined as the amount of heat energy transferred between the system and surroundings during a chemical reaction.

2. ENDOTHERMIC & EXOTHERMIC REACTIONS

  • Endothermic Reactions:

    • Overview: Reactions that gain thermal energy from the surroundings.

    • Characteristics:

    1. Energy of the surrounding decreases while the system energy increases.

    2. The energy gained for breaking bonds (Ea) exceeds the energy released from forming bonds.

    3. Enthalpy change (ΔH) is always positive.

    4. The energy of reactants is less than the energy of products.

    • Examples:

    • Dehydration reaction: CuSO4.5H2O + Heat → CuSO4 + 5H₂O

    • Thermal Decomposition: CuCO3 + Heat → CuO + CO₂

    • Process of melting, boiling, evaporation, and sublimation.

  • Exothermic Reactions:

    • Overview: Reactions that release thermal energy to the surroundings.

    • Characteristics:

    1. The surrounding temperature increases while the system energy decreases.

    2. The energy gained for breaking bonds (Ea) is less than the energy released from forming product bonds.

    3. Enthalpy change (ΔH) is always negative.

    4. The energy of reactants is greater than that of products.

    • Examples:

    • Combustion: CH4 + 2O2 → CO2 + 2H2O + Heat

    • Neutralization reaction: HCl + NaOH → NaCl + H2O + Heat

    • Processes like condensation and freezing.

6.2 STANDARD ENTHALPY CHANGE (ΔH)

1. THE STANDARD CONDITIONS

  • Standard conditions include:

    1. Pressure: 100 kPa (or 1.0 atm)

    2. Temperature: 298 K (or 25 °C)

    3. Concentration for aqueous solutions: 1.0 mol/dm³.

  • Each reactant and product must be in its standard physical state at these conditions (e.g., H₂O (l), O₂ (g)).

  • The standard enthalpy change (ΔH) is measured under these conditions and is expressed in kJ/mol.

Examples of Standard Enthalpy Changes

  1. Enthalpy Change of Reaction (ΔHᵣ):

    • Measured for the complete reaction under standard conditions for a balanced equation.

    • Example reaction:
      N2(g) + 3H2(g) → 2NH3(g)

  2. Enthalpy Change of Combustion (ΔHᶜ):

    • Measured when 1 mole of a substance is burned in excess oxygen under standard conditions.

    • Example:
      CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ΔHᶜ (CH4) = -890 kJ/mol

  3. Enthalpy Change of Neutralization (ΔHₙ):

    • Defined as the heat change when 1 mole of water is produced from neutralization of acid with base.

    • Enthalpy changes for strong acid-strong base reactions typically range around -57 to -58 kJ/mol.

    • Example:
      NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l) ΔHₙ = -57.1 kJ/mol

  4. Enthalpy Change of Formation (ΔH꜀):

    • Change measured when 1 mole of a compound is formed from its constituent elements in their standard states.

    • Example:
      2C(graphite) + 5H2(g) + 1/2O2(g) → C2H5OH(l) ΔH꜀ = -277 kJ/mol

  5. Enthalpy Change of Atomization (ΔHₐ):

    • Change measured when 1 mole of gaseous atoms is formed from an element in its standard state.

    • Example:
      C(s) → C(g)
      Also applicable to halogens and other elements transitioning from solid to gas phase.

2. DETERMINING STANDARD ENTHALPY CHANGE OF COMBUSTION (ΔHᶜ)

  • ΔHᶜ is typically negative, indicating that combustion is an exothermic process.

  • To compute in lab, calorimetry techniques can be employed:

  • Major Apparatus Needed:

    1. Spirit Burner: Contains fuel for combustion.

    2. Copper Can: Conducts heat efficiently to water used in the experiment.

    3. Thermometer: Measures temperature changes without touching container walls.

    4. Wind Shield: Reduces heat loss and improves accuracy.

Procedure Summary:

  1. Measure mass of the copper can (with water) - m_water.

  2. Measure mass of spirit burner (with fuel) - m₁.

  3. Record initial water temperature - Ti.

  4. Ignite the spirit burner, observe temperature increases.

  5. Turn off burner, record final temperature - Tf.

  6. Measure spirit burner mass again to find remaining fuel - m₂.

Calculation of ΔH:

  • Key Assumptions:

    1. All heat goes to water with no loss.

    2. No evaporation of fuel during combustion.

    3. Complete combustion occurs.

  • Density of Water: Equivalent to 1.0 g/cm³ for calculation purposes.

  • Specific Heat Capacity: For solutions considered to be the same as water, c = 4.18 J/g·K.

Heat Transfer Calculation:

  • The equation for heat transfer to water is: Q=mimescimesriangleTQ = m imes c imes riangle T Where:

    • $Q$ = Heat transfer to water (in J)

    • $m$ = Mass of water (in g)

    • $c$ = Specific heat capacity (4.18 J/g·K)

    • $T$ = Change in temperature (Tf - Ti)

Final Calculation for Enthalpy Change:

  • extNo.ofmolesoffuel=racextmassoffuelusedextmolarmassoffuelext{No. of moles of fuel} = rac{ ext{mass of fuel used}}{ ext{molar mass of fuel}}

  • Enthalpy change ($ ext{ΔH}$) is computed by:
    extΔH=−racQextNo.ofmolesoffuelext{ΔH} = - rac{Q}{ ext{No. of moles of fuel}}

Calculations Samples:

  1. For Hexane (C6H14):

    • Volume of water heated = 75.0 cm³

    • Initial mass of hexane = 250.0 g

    • Final mass of hexane = 65.50 g

    • Initial temperature = 23.7 °C

    • Final temperature = 52.2 °C

      • Result: extΔH=−4.16extkJ/molext{ΔH} = -4.16 ext{ kJ/mol}.

  2. For Propanol (C3H7OH):

    • Volume of water heated = 120.0 cm³, initial mass = 135.0 g, final mass = 78.0 g, Ti = 23.7 °C, Tf = 44.2 °C.

      • Result: extΔH=−0.058extkJ/molext{ΔH} = -0.058 ext{ kJ/mol}.

3. THE STANDARD ENTHALPY CHANGE OF NEUTRALIZATION (ΔHₙ)

  • ΔHₙ represents the heat change when an acid neutralizes a base to produce water under standard conditions.

  • Neutralization reactions for strong acids and bases yield a consistent ΔHₙ value around -57 to -58 kJ/mol.

  • Example of Reaction:

    • HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

  • Laboratory Method:

    1. Use a polystyrene cup as an insulating reaction container.

    2. Record initial temperatures for the acid and the base.

    3. Combine acid and base, monitor temperature rise; record maximum.

Calculation for ΔHₙ:

  • Key Assumptions:

    1. No heat loss to surroundings; all heat remains in solution.

    2. Density of aqueous solution approximated as water (1.0 g/cm³).

    3. Specific heat capacity equal to that of water.

Calculations Samples:

  1. For HCl + NaOH Reaction:

    • Initial acid volume = 25 cm³, initial alkali volume = 25 cm³, initial temperature = 25.4 °C, final temperature = 32.5 °C.

      • Result: extΔH=−57.7extkJ/molext{ΔH} = -57.7 ext{ kJ/mol}.

  2. For H2SO4 + KOH Reaction:

    • Volumes are 30.0 cm³ of H2SO4 and 20.0 cm³ of KOH; corresponding temperatures noted down.

      • Result: extΔH=−57.2extkJ/molext{ΔH} = -57.2 ext{ kJ/mol}.

6.3 HESS'S LAW

1. THE STANDARD ENTHALPY OF FORMATION (ΔHᶠ)

  • Defined under standard conditions when 1 mole of compound forms from elements in their standard states.

  • Values of ΔHᶠ are primarily negative due to bond formations, though exceptions may exist.

  • Examples of Standard Enthalpy of Formation:

    • Formation of ethanol: 2C(g) + 3H2(g) + 1/2O2(g) → C2H5OH(l). ΔHᶠ = -277 kJ/mol.

2. HESS'S LAW

  • Definition: The total enthalpy change of a reaction is independent of the pathway taken, which allows calculation of enthalpy changes using indirect routes through related reactions.

  • Practical Application: Allows determination of ΔH for difficult to measure reactions using the enthalpy changes of known reactions (using Hess's cycle or enthalpy cycles).

  • Methods for Calculation:

  1. Rule Method:

    • extΔH(reaction)=extΣΔH(products)−extΣΔH(reactants)ext{ΔH(reaction)} = ext{Σ ΔH(products)} - ext{Σ ΔH(reactants)}

  2. Energy (Hess's) Cycle Method:

    • Sequence of necessary reactions leading from reactants to products.

Calculation Examples:

  1. For Reaction

    • C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g) gives:

    • −2221extkJ/mol-2221 ext{ kJ/mol} using the summation of formation enthalpies.

  2. Using Hess's Cycle:

    • Construct cycles from the enthalpy formations of individual components to calculate overall changes.

3. CALCULATING AHF FROM ΔHᶜ

  • Method 1: Flip & Subtract

  • Method 2: Energy Cycle Method

  • Calculation Samples:

  1. Standard Enthalpy of Formation of Ethane, C2H6:

    • From combustion enthalpies:

    • −84.7extkJ/mol-84.7 ext{ kJ/mol} calculated using both methods presented above.

  2. Ammonium Ethanoate Reaction:

    • Uses enthalpy calculations to determine values using Hess's cycles as described previously.

6.4 BOND ENTHALPY

  • The bond enthalpy is the measure of strength of a bond, defined as the energy required to break 1 mole of a bond in gaseous molecules.

  • Examples:

  1. O=O bond in O2.

  2. C=O bond and C-H bonds in hydrocarbons to assess their stability and reactivity in various reactions.

Summary Points:

  • Stronger the bond, higher the bond enthalpy.

  • Reactants and products' bond enthalpies can be utilized to calculate overall enthalpy changes for reactions through the strategies outlined (Hess's law and bond enthalpy).

Notes for Further Study:

  • Explore exam questions and perform practical experiments to enhance comprehension of these topics.