Unit 8 and 9 Thermochemistry and Kinetics

Energy

The capacity to do work or produce heat.

Heat

The transfer of thermal energy between systems.

Specific Heat Capacity

The amount of heat required to raise the temperature of one gram of a substance by one degree Celsius.

Endothermic vs Exothermic

Endothermic processes absorb heat, while exothermic processes release heat.

Enthalpy

A measure of the total heat content in a system.

Hess's Law

The total enthalpy change for a reaction is the sum of the enthalpy changes for the individual steps.

Entropy

A measure of the disorder or randomness in a system.

Convection/Conduction

Convection is the transfer of heat by the movement of fluids, while conduction is the transfer of heat through direct contact.

Heat of Fusion vs Heat of Vaporization

Heat of fusion is the energy required to change a substance from solid to liquid, while heat of vaporization is the energy required to change a substance from liquid to gas.

When to use q=mCΔT vs ΔHvap = mHv or ΔHfus = mHf

Use q=mCΔT for temperature changes, and ΔHvap = mHv or ΔHfus = mHf for phase changes.

Spontaneity

The tendency of a process to occur without outside intervention.

When ΔS and ΔH are + or - (chart)

A chart that indicates the signs of entropy change (ΔS) and enthalpy change (ΔH) for various processes.

Collision Theory

A theory that states that for a reaction to occur, reactant particles must collide with sufficient energy and proper orientation.

Factors that influence rxn rate

Concentration, temperature, surface area, and catalysts are factors that can affect the rate of a reaction.

Rate Determining Step

The slowest step in a reaction mechanism that determines the overall reaction rate.

Equilibrium expressions

Mathematical representations of the concentrations of reactants and products at equilibrium.

Activation Energy/Heat of Reaction

Activation energy is the minimum energy required for a reaction to occur, while heat of reaction is the change in enthalpy during a reaction.

Heat of Formation

The change in enthalpy when one mole of a compound is formed from its elements in their standard states.

Nomenclature

The system of naming chemical compounds.

Q=mCΔT

An equation used to calculate heat transfer, where Q is heat, m is mass, C is specific heat, and ΔT is the change in temperature.

ΔHvaporization = mHv

An equation for calculating the enthalpy of vaporization, where ΔHvaporization is the heat required to vaporize a substance, m is mass, and Hv is the specific heat of vaporization.

ΔHfusion = mHf

An equation for calculating the enthalpy of fusion, where ΔHfusion is the heat required to melt a substance, m is mass, and Hf is the specific heat of fusion.

ΔHrxn = Σ∆H0f (products) - Σ∆H0f (reactants)

An equation for calculating the change in enthalpy of a reaction using the standard enthalpies of formation.

ΔG = ΔH - TΔS

An equation for calculating Gibbs free energy, where ΔG is Gibbs free energy, ΔH is enthalpy, T is temperature, and ΔS is entropy.

Rate = k[A]n and (R2/R1) = (M2/M1)m

Equations used to describe reaction rates, where k is the rate constant, [A] is concentration, and m is the experimental exponent.

Keq = [C]c[D]d / [A]a[B]b

An equation for calculating the equilibrium constant, where [C], [D], [A], and [B] are the concentrations of the products and reactants.

Thermochemistry

Concerned with heat changes that occur during chemical reactions or phase changes.

Chemical potential energy

Energy within chemical substances.

Conduction

Heat transferred by particles colliding into one another, primarily in solids.

Convection

Heat transferred by the circulation of a fluid (or gas), such as in a heating system.

Radiation

Heat transferred by the flow of electromagnetic radiation; can happen in a vacuum.

System

The specific part of the universe that contains the reaction or process you wish to study.

Surroundings

Everything else outside the system.

Exothermic process

Heat flowing out of a system into its surroundings; system loses heat.

Endothermic process

Heat flowing into a system from its surroundings; system gains heat.

Law of Conservation of Energy

In any chemical or physical process, energy is neither created nor destroyed.

Heat of Fusion

Energy required to change from solid to liquid; some attractive forces are broken.

Heat of Vaporization

Energy required to change from liquid to gas; typically higher than heat of fusion.

Thermal Energy

Total energy of particles, including kinetic and potential energy.

Kinetic Energy

Energy of motion of particles; increases as particles move faster.

Potential Energy

Energy of arrangement of particles; increases as particles move farther apart.

Temperature

Measure of the average kinetic energy of all the particles in a sample.

Heating Curve

Graph showing energy (heat) being added as time passes.

Calorie

Quantity of heat needed to raise the temperature of 1 g of pure water by 1 °C.

Calorie (capital C)

Refers to the energy in food; 1 Calorie = 1 kilocalorie = 1000 cal.

Phase changes

Can be exothermic or endothermic; going from solid to liquid to gas requires heat input.

Exothermic reactions

Release energy, usually in the form of heat.

Endothermic reactions

Absorb energy.

Energy storage

Energy is stored in bonds between atoms.

Heat Capacity

The amount of heat needed to increase the temperature of an object exactly 1 °C.

Specific Heat of Water

For water, C = 4.18 J/(g °C) in Joules, and C = 1.00 cal/(g °C) in calories.

Heat Absorbed or Released

Calculated using the formula q = mCpΔT, where q is the heat absorbed or released.

q

The heat absorbed or released during a change in temperature.

Cp

The specific heat of the substance.

m

The mass of the substance in grams.

ΔT

The change in temperature in °C.

Specific Heat Example 1

For 34.4 g of ethanol increasing from 25.0 °C to 78.8 °C, q = 4.52 x 10^3 J.

Specific Heat Example 2

For a 155 g sample of ethanol heated from 25.0 °C to 40.0 °C, Cp = 2.45 J/g·°C.

Specific Heat Example 3

For a 95.4 g piece of copper increasing from 25.0 °C to 48.0 °C, the specific heat is 0.387 J/g°C.

Specific Heat Example 4

A piece of copper absorbs 249.5 J of heat, changing temperature from 21.3 °C to 26.8 °C, mass = 117 g.

Specific Heat Example 5

A 75.8 g piece of titanium absorbs 642.1 J of heat, starting at 21.2 °C, ending at 37.4 °C.

Calorimetry

The measurement of the heat into or out of a system for chemical and physical processes.

Calorimeter

The device used to measure the absorption or release of heat in chemical or physical processes.

Enthalpy (H)

The heat content of a system at constant pressure.

Calorimetry Example 1

A 35.0 g piece of metal heated to 100 °C in 50.0 g of water at 21.4 °C, water temp rises to 25.8 °C.

Calorimetry Example 2

A 37.6 g sample of metal heated to 100.0 °C in 75.0 g of water at 22.3 °C, water temp rises to 29.4 °C, Cmetal = 0.839 J/g°C.

Calorimetry Example 3

A 15.7 g sample of metal heated to 95.0 °C in 35.0 g of water at 21.2 °C, water temp rises to 25.9 °C, Cmetal = 0.634 J/g°C.

Exothermic Reaction

A reaction that produces heat.

Change in Enthalpy (ΔH)

The difference between the enthalpy of the substances at the end and the start of a reaction.

Enthalpy of Reaction (ΔHrxn)

The change in enthalpy for a reaction.

Thermochemical Equation

A balanced equation that includes states of matter and the energy change.

Enthalpy of Combustion (ΔHcomb)

The enthalpy change for the complete burning of one mole of a substance.

Enthalpy of Fusion (Hf)

The heat required to melt one gram of a solid substance.

Enthalpy of Vaporization (Hv)

The heat required to vaporize one gram of a liquid substance.

Heat (q)

The energy transferred due to temperature difference.

Formula for Heat (q)

q = mCpΔT

Formula for Enthalpy of Fusion

ΔH = mHf

Formula for Enthalpy of Vaporization

ΔH = mHv

Molar Enthalpy of Fusion

ΔH = nHf

Energy Released in Reaction

ΔHrxn = Hproducts - Hreactants must result in a negative number.

Combustion of Methane

To liberate 12,880 kJ of heat, 231 g of CH4 must be burned.

Heat from Ammonia Condensation

q = 275 g x 1371.2 J/g = 377080 J or 377 kJ.

Heat to Melt Ice

q = 24.5 g x 334 J/g = 8183 J.

Heat to Vaporize Water

q = 24.5 g x 2260 J/g = 55370 J.

Heat to Melt Ice at -20.0 °C

Total q = 2050 J (heating) + 16700 J (melting) = 18750 J.

Heating Ice to Steam

Energy needed to heat 55.0 grams of ice from -15.0 °C to steam at 150.0 °C.

Step 1 Heating Ice

q = (55 grams)(2.05 J/g°C)(150) = 1,691 J.

Step 2 Phase Change

ΔH = (55 g)(334 J/g) = 18370 J.

Step 3 Heating Liquid

q = (55 grams)(4.18 J/g°C)(100) = 22,990 J.

ΔH

The change in enthalpy for a reaction.

ΔHrxn

The enthalpy of reaction, defined as Hfinal - Hinitial.

Endothermic Reaction

A reaction that absorbs heat, resulting in products having more energy than reactants.

mCpΔT

The formula used to calculate heat, where m is mass, Cp is specific heat capacity, and ΔT is the change in temperature.

Specific Heat Capacity (C)

The amount of heat required to raise the temperature of 1 gram of a substance by 1 °C.

Standard Enthalpy of Formation

The enthalpy change for the formation of one mole of a compound from its elements in their standard states.

ΔHtotal

The total enthalpy change calculated by summing individual enthalpy changes.

Heat Released

The energy released to the environment during a cooling process.

Combustion of Propane

The reaction of propane with oxygen producing carbon dioxide and water, with an associated ΔHrxn of -2044 kJ.

Mass of Water

The amount of water involved in a reaction, affecting the heat transfer calculations.

Temperature Change

The difference in temperature before and after a reaction, used to calculate heat transfer.