Thermodynamics: Energy, Heat, and Chemical Reactions

• Chapter 5 discusses thermodynamics, including energy, heat, work, enthalpy, and calorimetry.
• Principles of Chemical Reactivity focuses on energy and chemical reactions.

Key Questions Explored

• How to measure and calculate energy changes in physical changes and chemical reactions
• Connection between energy changes, heat, and work
• Determining if a reaction favors products or reactants at equilibrium
• Conditions for spontaneous chemical reactions or physical changes

Basic Thermodynamic Principles

• Energy has the ability to do work or transfer heat, represented by q.
• Kinetic energy is the energy associated with motion, including thermal, mechanical, and electrical energy.
• Potential energy is stored energy resulting from an object's position, such as gravitational, electrostatic, chemical, and nuclear energy.
• The law of conservation of energy says that energy is neither created nor destroyed, only converted.
• Heat (q) is not temperature; temperature measures kinetic energy.

Units of Energy

• 1 calorie is the heat needed to raise the temperature of 1.00 g of H₂O by 1.0 °C.
• 1000 calories equals 1 kilocalorie (kcal).
• 1 kcal is equivalent to 1 Calorie, often used for food.
• The SI unit for energy is the joule (J).
• 1 calorie is equivalent to 4.184 J.
• James Joule (1818-1889) contributed significantly to the study of energy.

System and Surroundings

• A system refers to the object or collection of objects under study.
• Surroundings encompass everything outside the system that can exchange energy and/or matter with it.
• Energy flows between the system and its surroundings.

Energy Transfer

• An exothermic process happens when energy leaves the system and enters the surroundings.
• The temperature of the system decreases in an exothermic process, with qsystem < 0 and ∆Tsystem = (Tfinal – Tinitial) < 0.
• An endothermic process happens when energy enters the system from the surroundings.
• The temperature of the system increases in an endothermic process, with qsystem > 0 and ∆Tsystem = (Tfinal – Tinitial) > 0

Specific Heat Capacity

• The quantity of energy transferred when heating or cooling an object depends on the material's quantity, magnitude of the temperature change, and identity, including its phase (gas, liquid, aqueous, or solid).
• Specific heat capacity (C) measures the energy (heat) needed to raise the temperature of 1 gram of a substance by 1 Kelvin, with units of J/g ⋅ K.
• Molar heat capacity is the energy transferred as heat to raise the temperature of one mole of a substance by one Kelvin, measured in J/mol ⋅ K.
• The energy gained or lost as heat when warming or cooling a substance is calculated using q = C × m × ∆T.

Thermal Equilibrium

• Energy transfer (as heat) occurs spontaneously from a higher temperature to a lower temperature.
• Transfer happens until both objects reach the same temperature and achieve thermal equilibrium.
• At thermal equilibrium, the object that increased in temperature gained thermal energy, while the one that decreased in temperature lost thermal energy.
• The sum of energy changes in an isolated system is zero due to the law of conservation of energy.
• Thermal energy changes involve water and metal, with water + metal = 0.

Changes of State

• A change of state involves transitions between solid, liquid, and gas.
• When a solid melts, atoms, molecules, or ions gain enough energy to overcome attractions, allowing them to move more freely.
• Temperature remains constant during a change of state; added energy overcomes intermolecular forces.

Enthalpy of Phase Changes

• Sublimation, vaporization, and melting/fusion are endothermic processes with positive ΔH values.
• Deposition, condensation, and freezing are exothermic processes with negative ΔH values.
• ΔH refers to the heat of a phase change, which can be calculated using qphase change = ΔHphase change × n.

Calculating Energy and Phase Changes

• Variables include Heat capacity, energy and moles or mass.
• When converting ice at -50.0°C to steam at 200.0°C, determine what the specific capacities of ice, water and water vapour are.
• The heat of fusion of water is 333 J/g.
• The heat of vaporization is 2256 J/g

First Law of Thermodynamics

• Thermodynamics explores heat and work.
• Work done by or on a system affects its energy.
• An energy balance includes energy transferred as heat and work.
• From the first law of thermodynamics, when a system undergoes a chemical change, the change in internal energy is given by ∆U = q + w.
• U is the change in internal energy, q is the heat and w is the work.
• It is crucial to use the correct sign conventions in the equation.

Sign Conventions

• Energy entering the system is endothermic (+q), while energy leaving is exothermic (-q).
• Work done on the system is (+w), and work done by the system is (-w).
• The effect on Usystem depends on the sign convention: U increases with energy or work to the system and decreases when they leave the system.

Pressure-Volume

• Work is the work done when system volume changes.
• The work to move a piston is calculated by w = F × d.
• Pressure is the force over area, and can be written as P = F/A.
• Substituting P × A for F gives w = (P × A) × d.
• A × d is equivalent to the volume change, where ∆V = Vfinal – Vinitial.
• Work done by a system is negative, so w = -P∆V.

Constant Volume

• In a constant-volume process ∆V = 0, so energy transferred as work is zero.
• At constant volume, the internal energy change is equal to the energy transferred as heat (qv).

Enthalpy

• Enthalpy (“H”) is the heat transferred between a system and surroundings at constant pressure.
• Under constant pressure the change in Enthalpy equals pressure plus work, so; AU = qp + Wp.
• If only type of work is P–V work, then, AU = qp - PAV
• If no PV work is done by the system, the change in enthalpy is the change in internal energy at constant pressure.

State Functions

• State functions like internal energy and enthalpy depend only on initial and final states, not the path.
• The change in altitude is a state function that does not consider the distance travelled.

Enthalpy Measurement

• Enthalpies are state functions and must be measured/reported carefully.
• H(T, P) indicates that enthalpy is a function of temperature and pressure.
• To specify that the enthalpy is taken at standard state conditions, H° is used.
• Standard state conditions include: 1 atm = 760 mm Hg or 760 torr and 298.15 K (25 °C).

Chemical Reactions

• Enthalpy changes accompany chemical reactions.
• For reactions, decomposition can be defined by
• For reverse reaction, represents formation in an amount of water.
• For the amount of propane present in the sample, it can be determined with heat.

Bomb Calorimetry

• Constant volume calorimeters have two types ( constant volume and constant pressure).
• Measure heats of combustion using a bomb.
• Calorimeters all energy are transferred as heat so you can determine heat in any reaction.
• All contents involved in bomb devices are defined as systems to asses heat.

Using Hess's Law

• Not all processes can be directly measured with calorimetry.
• It provides an indirect method that can be used.
• Enthalpy changes for oxidation can be directly studied.
• Manipulated equations must be added together to yield the net equation.
• When 2 mol H2(g) is on the reactant side in the equation 2H2(g) + Q2(g) -> 2 h20(1), multiple with given factors.
• The final equation when adding the components is
• For forming 1 mol CH4, it equals certain factor.

Formations

• When one molecule of compound is formed from elements at standard state conditions, it is known as a formation equation.
• Use AH to form calculate the changes based on above, add product and subtract reaction.

Nitroglycerin Explosions

• Nitroglycerin explosives are very strong, releasing many different gases when detonated
• Can calculate detonatined nitroglycerin, which is used to determine heat formation.