Unit 1: Thermochemistry

Unit 1: Thermochemistry Introduction and Some Definitions Internal Energy The First Law of Thermodynamics Enthalpy and Enthalpy Changes Calorimetry Hess’s Law Using Enthalpy’s of Formation

Thermochemistry – Some Definitions • Most daily activities involve processes that either use or produce energy: • Metabolism of food • Burning fossil fuels • Photosynthesis • Pushing a bike up a hill • Thermodynamics: • the study of the energy and its transformations

Thermochemistry – Some Definitions • Thermochemistry: • A branch of thermodynamics • the study of the energy absorbed or released as heat during a chemical reaction or process • Objects (including chemicals) can have two types of energy: • Kinetic energy • Potential energy

Thermochemistry – Some Definitions • Kinetic energy • Energy of motion • Thermal energy • a type of kinetic energy a substance possesses because of its temperature. • Potential energy • “stored” energy or energy of position • Energy that results from attractions and repulsions between objects

Thermochemistry – Some Definitions • Chemical energy: • A type of potential energy stored within a substance • Results from electrostatic forces between charged particles within the substance as well as from the arrangement of the atoms (or ions) within the substance

Thermochemistry – Energy Units • Units of Energy: • SI unit = joule (J) • A very small quantity • ~ the energy required to lift a small apple one meter into the air • Kilojoule (kJ) • 1 kJ = 1000 J

Thermochemistry – Energy Units • Calorie (cal) • Originally defined as the amount of energy needed to raise the temperature of 1 g of water from 14.5oC to 15.5oC • 1 cal = 4.184 J (exactly) • Kilocalories (kcal) • 1 kcal = 1000 cal = 1 Cal (food calorie) • British Thermal Unit (BTU) • 1 BTU = 1054.35 J

Thermochemistry – Energy Units • You are responsible for knowing the conversion factors shown in red on the previous slides. • You must be able to use dimensional analysis to convert from one energy unit to another.

Thermochemistry – Energy Units Example: Convert 7.63 kcal to BTU using the relationships from the previous slides.

Thermochemistry – Energy Units Example:A particular furnace produces 9.0 x 104 BTU/hr of heat. Use dimensional analysis to calculate the number of kcal of heat delivered by the furnace after running for 2.50 hours.

Thermochemistry – More Definitions • When studying the amount of heat gained or lost during a process or reaction, chemists focus on a limited, well-defined part of the universe to study: • System: • The part of the universe singled out for study • Typically the chemicals involved in the reaction or process • Surroundings: • Everything else

Thermochemistry – More Definitions The system is usually the chemicals in the flask/reactor. The system The flask and everything else belong to the surroundings.

Thermochemistry – More Definitions • A system can be either open or closed. • Open system: • A system that can exchange both matter and energy with the surroundings • Closed system: • A system that can exchange energy with the surroundings but not matter. A cylinder with a piston is one example of a closed system.

Internal Energy • The internal energy (E) of a system is the sum of the kinetic and potential energy of all components of a system. • For the molecules in a chemical system, the internal energy includes: • The motion and interactions of all of the molecules • The motion and interactions of the nuclei and electrons found in the molecules

Internal Energy • Internal energy (E) is an extensive property. • Depends on the mass of the system • Internal energy (E) is a state function. • A property of the system that is determined by specifying its condition or state in terms of T, P, location, etc. • Depends only on its present condition andnot how it got there

Heat and Work • The internal energy (E) of a system can change when the system gains energy from or loses energy to the surroundings as either heat (q) or work (w). • Work (w): • Energy used to move an object against a force • Lifting your backpack • Hitting a baseball with a bat

Heat and Work • Heat (q): • Energy used to increase the temperature of an object • The energy transferred from a hotter object to a colder one • Energy: • The capacity (ability) to do work or transfer heat

system (+) q (+) w Heat & Work – Sign Conventions • Energy can be transferred between the system and the surroundings as either heat or work. • Energy gained by the system is always designated using a positive sign. • A reaction or process in which the system gains heat from the surroundings is endothermic.

system q (-) w (-) Heat & Work – Sign Conventions • Energy lost by the system is always designated using a negative sign. • A reaction or process in which the system loses heat to the surroundings is exothermic.

Internal Energy Changes • The change in the internal energy (DE) that occurs when energy is gained from or lost to the surroundings: D E = Efinal - Einitial DE = change in internal energy Efinal = final energy of system Einitial = initial energy of system

Internal Energy Changes • A reaction or process that experiences a net gain of energy from the surroundings is referred to as endergonic. • Efinal > Einitial • DE > 0 (positive)

Internal Energy Changes • A reaction or process that experiences a net loss of energy to the surroundings is referred to as exergonic. • Einitial > Efinal • DE < 0 (negative)

Internal Energy Changes • From a practical perspective, the change in internal energy (DE) of the system is found by measuring the amount of heat gained or lost by the system and the amount of work done on or by the system: DE = q + w Where q = heat w = work Be sure to use the correct sign for q and w!

Internal Energy Changes Example: Calculate the change in internal energy of the system for a process in which the system absorbs 197 J of heat from the surroundings while doing 73 J of work.

Internal Energy Changes Example: Calculate DE for a system when the system loses 72 kJ of heat while the surroundings do 193 kJ of work on the system.

First Law of Thermodynamics • Energy can be transferred between the system and the surroundings as heat and/or work. • Energy can also be converted from one form to another. • Kinetic energy  Potential energy • First Law of Thermodynamics: Energy can be converted from one form to another, but it cannot be created or destroyed. • Any energy lost by the system must be gained by the surroundings and vice versa.

Enthalpy • Many reactions or chemical processes occur in open containers (i.e. at constant pressure). • The amount of heat gained or lost under constant pressure conditions (qp) is often referred to as the enthalpy change (DH) • Enthalpy (H): • An extensive property (one that depends on the amount of substance present) that is defined by the equation H = E + PV

Calculating the Amount of Work • Two common types of work done by chemical systems: • Electrical work • Redox reaction incorporated in galvanic cells (Unit 5) • Mechanical work (P-V work) • Work done by expanding gases • For example, the expanding gases in a cylinder of a car engine

Calculating the Amount of Work • The amount of P-V work done at constant pressure can be found: • w = -P DV where P = pressure DV = change in volume • The negative sign indicates that work is being done by the system

DE = DH – P DV • For a process occurring at constant pressure in which the only work done is PV work, DE = qp + w DH = qp w = -P DV DE = DH – P DV (at constant pressure)

DE = DH – P DV Example: If the volume of a cylinder increases from 2.13 L to 3.50 L at a constant pressure of 1.33 atm while it absorbs 2.515 kJ, what is the change in internal energy of the system? (Note: 1 atm.L = 101.3 J)

Unit 1: Thermochemistry