CH 5: Principle of Chemcial Reactivity

Chemistry and Chemical Reactivity: Principles of Chemical Reactivity: Energy and Chemical Reactions

Overview

• Focus on energy changes associated with physical changes and chemical reactions.
• Key questions raised include:
  • How to measure and calculate energy changes?
  • Relationship between energy changes, heat, and work.
  • Determining product-favored vs reactant-favored chemical reactions at equilibrium (Chapter 18).
  • Identifying spontaneous chemical reactions or physical changes.

Types of Energy

• Energy: Capacity to do work.
  • Kinetic Energy (KE): Associated with motion.
  • Formula: KE = rac{1}{2}mv^2
  • Types: Thermal (motion of particles), mechanical (moving objects), electrical (electron movement).
  • Potential Energy: Associated with position.
  • Types: Gravitational, electrostatic, chemical energy associated with chemical bonds.

Law of Conservation of Energy

• Energy cannot be created or destroyed.
• Total energy of an isolated system remains constant. Mathematically presented as: $\Delta U = q + w$
  • Where:
  • $\Delta U$: Change in internal energy.
  • $q$: Heat exchanged.
  • $w$: Work done.

Molecular Level

• Thermal Energy: Kinetic energy of molecules; related to heat transfer between objects at different temperatures.
• Chemical Energy: Potential energy linked to the arrangement and bonds of atoms and molecules.

Heat

• Heat: Transfer of thermal energy from a hotter object to a colder one.
  • Heat flows from objects with higher average kinetic energy to those with lower.
  • Symbol for heat: $q$.

Systems and Surroundings

• System: Part of the universe under study (e.g., contents of a beaker).
  • Types:
  • Open System: Exchanges matter and energy with surroundings.
  • Closed System: Exchanges energy, but not matter.
  • Isolated System: Exchanges neither matter nor energy.
• Surroundings: Everything outside the system.

Energy Transfer

Energy Changes

• $\Delta U = q + w$
• Endothermic Process: System absorbs heat; q > 0.
• Exothermic Process: System releases heat; q < 0.

Energy Units

• Joules (J): Official SI unit for energy.
• Calories:
  • 1 Cal = 1000 cal (1 kcal).
  • 1 cal = 4.184 J (exact).

Concept Check

• Convert $3.75 \times 10^3$ calories into Joules:
  • $3.75 \times 10^3 \text{ cal} \times 4.184 \text{ J/cal} = 15665 \text{ J}$.

Heat Capacity

• Heat Capacity: Amount of heat required to raise the temperature of an object by 1 K (1 °C).
  • Formula: $q = C_{object} \times \Delta T$
• Where:
  • $q$: Energy gained or lost as heat.
  • $C_{object}$: Heat capacity of an object (J/K).
  • $\Delta T$: Change in temperature.
• Extensive Properties: Depend on material quantity (e.g., mass, volume).
• Intensive Properties: Independent of amount (e.g., density, specific heat capacity).

Specific Heat Capacity

• Quantity of heat required to raise 1 gram of substance by 1 K (or °C).
• Formula: $q = C \times m \times \Delta T$
  • Where:
  • $C$: Specific heat capacity (J/K/g).
  • $m$: Mass.

Example Calculation of Specific Heat Capacity

• For water: $C_{water} = 4.184 \, \text{J/g/K}$.
• Example: Heating 50 g of water from 25.3 °C to 33.7 °C.
  • $\Delta T = 33.7 - 25.3 = 8.4 \, K$
  • Calculate heat transfer: $q = 4.184 \, \text{J/g/K} \times 50 \, \text{g} \times 8.4 \, K = 1757.3 \, J$

Quantitative Aspects of Heat Transfer

• Energy transfers spontaneously from hotter to colder objects until thermal equilibrium is reached.
• Thermal Equilibrium: No net transfer of heat between objects.

Law of Conservation of Energy in Systems

• In an isolated system, the sum of energy changes is zero.

Calorimetry and Energy Transfer

Example Problem in Calorimetry

• A 145 g sample of copper at 100.0 °C placed in 250.0 g of water at 25.0 °C, reaching thermal equilibrium at 28.8 °C.
• Calculate the specific heat capacity of copper:
  • $q<em>{Cu} + q</em>{H2O} = 0$
  • Use:
    $m<em>{Cu} \times C</em>{Cu} \times (T<em>f - T</em>i) + m<em>{H2O} \times C</em>{H2O} \times (T<em>f - T</em>i) = 0$.

First Law of Thermodynamics

• Based on conservation of energy:
  • Internal energy change: $\Delta U = q + w$
• Implications for endothermic and exothermic processes:
  • Positive Heat Transfer: Energy absorbed; increases internal energy.
  • Negative Heat Transfer: Energy released; decreases internal energy.

Enthalpy

• Enthalpy (H): Defined as $H = U + PV$
• Change in enthalpy (ΔH): $\Delta H = H<em>{final} - H</em>{initial}$
• At constant pressure, $\Delta H = q_p$.

Thermochemical Equations

• Example reactions with enthalpy changes:
  • $C(s) + O<em>2(g) → CO</em>2(g) \, (\Delta H = -393.5 \, kJ/mol)$.

Example Problem for Hess’s Law

• Determine energy changes for reactions using known enthalpy values by manipulation of stoichiometric equations.
  • Energy change for combustion of glucose:
    $C<em>6H</em>{12}O<em>6(s) + 6 O</em>2(g) → 6 CO<em>2(g) + 6 H</em>2O(l) \, (answer = -2801 \, kJ)$.
• Standard enthalpy of formation (ΔfH°) defined at standard conditions,
  • Values indicate the stability of compounds under formation from elements.

Product vs Reactant Favored Reactions

• Product-favored reactions generally have negative ΔH°rxn.
• Reactant-favored reactions generally possess positive ΔH°rxn.

Note: This document is a fully detailed study guide based on the principles of chemical reactivity regarding energy changes, heat transfer, heat capacity, calorimetry, and thermochemistry.