Thermochemistry Lecture Review

Introduction to Thermochemistry

Thermochemistry is defined as the study of heat change in chemical reactions.

• The System: This refers to the specific part of the universe that is of interest in a particular study.

• The Universe: Comprises both the system and the surroundings.

• Exchange Types:     * Open System: Exchanges both mass and energy with surroundings.      * Closed System: Exchanges energy, but nothing (mass) else.      * Isolated System: Exchanges nothing (neither mass nor energy).

Interdisciplinary Applications:

• Food Science: Measuring calories and energy content in various foods.

• Biology: Studying metabolism in living organisms.

• Materials Science: Investigating heat transfer in different materials.

Energy Definitions and Classifications

Energy is broadly defined as the capacity to supply heat or do work. While chemistry focuses primarily on the supply of heat, physics focuses on the ability to do work.

Forms of Energy:

• Radiant Energy: Derived from the sun; it is Earth's primary energy source.

• Thermal Energy: Energy associated with the random motion of atoms and molecules.

• Chemical Energy: Energy stored within the chemical bonds of substances.

• Nuclear Energy: Energy stored within the collection of neutrons and protons in an atom.

• Potential Energy: Energy available by virtue of an object's position.

• Kinetic Energy (KE): The energy of motion.

Kinetic Energy Formula and Units: Kinetic energy is measured in Joules ($J$).

$KE = \frac{1}{2}mv^2$

• $m$ = mass (measured in $kg$)

• $v$ = velocity (measured in $m/s$)

• $1\text{ Joule} = 1\,kg\,m^2/s^2$

Example Calculation:

• What is the KE of a $0.454\,kg$ ball moving at a velocity of $201\,m/s$?

Conservation of Energy and Thermal Dynamics

Conservation of Energy: Energy can be converted from one form to another, but it is never created or destroyed. Examples of conversions include:

• Potential energy into kinetic energy.

• Potential energy into heat (thermal energy).

• Potential energy into light or electrical energy.

Thermal Energy Specifics: Thermal energy is the kinetic energy associated with the random motion of:

• Atoms: Found in monoatomic gases and liquids.

• Molecules: Found in diatomic gases and compounds.

Heat vs. Temperature:

• Heat: The transfer of thermal energy between two bodies that are at different temperatures.

• Temperature: A measure of thermal energy. In this context, Temperature is proportional to Thermal Energy.

Energy Changes in Chemical Reactions

Process Definitions:

• Exothermic Process: Any process that gives off heat, thereby transferring thermal energy from the system to the surroundings.

• Endothermic Process: Any process in which heat must be supplied to the system from the surroundings.

Reaction Examples:

• Exothermic:     * $2H_2(g) + O_2(g) \rightarrow 2H_2O(l) + \text{energy}$     * $H_2O(g) \rightarrow H_2O(l) + \text{energy}$

• Endothermic:     * $\text{energy} + 2HgO(s) \rightarrow 2Hg(l) + O_2(g)$     * $\text{energy} + H_2O(s) \rightarrow H_2O(l)$

Units of Energy Measurement

• calorie (cal): A historical unit of measurement. It is the energy required to raise $1\,gram$ of water by $1\,^{\circ}C$ (or $1\,K$).

• Calorie (Cal): Written with a capital C, this equals $1000\text{ calories}$ or $1\text{ kilocalorie (kcal)}$. This is the unit used to express food energy content.

• Joule (J): The SI unit of heat, work, and energy. It is defined as the energy required when a force of $1\text{ newton}$ moves an object $1\text{ meter}$.

Describing Heat Flow: Heat Capacity and Specific Heat

Heat Capacity (C): The quantity of heat ($q$) that an object absorbs when its temperature increases by $1\,^{\circ}C$ (or $1\,K$).

• A smaller object (e.g., a small iron pot) has a smaller heat capacity than a larger object of the same material.

• General Limitation: It must be determined uniquely for every specific object, making it less universal for calculations.

Specific Heat Capacity (c or s): Often referred to simply as "Specific Heat," this is the quantity of heat required to raise the temperature of exactly $1\,gram$ of a substance by $1\,^{\circ}C$ (or $1\,K$).

• Specific heat corrects for the mass of the substance.

• It is a material-specific property, not dependent on the size or mass of the object.

Comparing Specific Heats of Materials:

• Specific Heat Values ($J/g\,^{\circ}C$):     * Liquid Water: $4.18$     * Solid Water (ice): $2.11$     * Water vapor: $2.00$     * Dry air: $1.01$     * Basalt: $0.84$     * Granite: $0.79$     * Aluminum: $0.897$     * Iron: $0.449$ (or $0.45$)     * Copper: $0.38$     * Lead: $0.13$

Comparative observations:

• Materials with lower specific heat (like lead or iron) heat up and cool down faster than materials with high specific heat (like water).

• Water ($4.184\,J/g\,^{\circ}C$) heats up much slower than Aluminum ($0.897\,J/g\,^{\circ}C$).

Quantitative Heat Calculations

Core Formulas:

1. Specific Heat ($s$): The amount of heat ($q$) required to raise the temperature of one gram of the substance by one degree Celsius.

2. Heat Capacity ($C$): The amount of heat ($q$) required to raise the temperature of a given quantity ($m$) of the substance by one degree Celsius.

3. Heat Equation: $q = m \times s \times \Delta t$

Problem Examples:

1. Iron Bar Cooling: Calculation of heat given off when an $869\,g$ iron bar cools from $94\,^{\circ}C$ to $5\,^{\circ}C$.

2. Water Heating (Example 9.1): A flask containing $8.0 \times 10^2\,g$ of water increases from $21\,^{\circ}C$ to $85\,^{\circ}C$. Calculate heat absorbed.

3. Copper Heating: $5.0\,g$ of copper heated from $20\,^{\circ}C$ to $80\,^{\circ}C$. Specific heat of Cu is $0.092\,cal/g\,^{\circ}C$. Calculate energy used.

4. Mass of Water Sample: A sample increases from $20\,^{\circ}C$ to $46.6\,^{\circ}C$ by absorbing $5650\,calories$. Find the mass ($s = 1.0\,cal/g\,^{\circ}C$).

Calorimetry

Calorimetry is a technique used to measure the amount of heat from a chemical reaction or physical change using an instrument called a calorimeter.

Principles of Calorimetry:

• The calorimeter helps define the System (the chemical reaction) and the Surroundings (the water inside the calorimeter).

• Endothermic ($+q$): The reaction absorbs heat from the surroundings. The thermometer shows a temperature decrease in the water.

• Exothermic ($-q$): The reaction releases heat into the surroundings. The thermometer shows a temperature increase in the water.

Types of Calorimeters:

1. Cup Calorimeter: Used for reactions in solution.

2. Bomb Calorimeter: Used for combustion reactions.

Calorimetry Equations: Assuming no heat enters or leaves the system ($$q_{sys} = 0