Chapter 6: Thermochemistry: Energy Flow and Chemical Change

Merged flashcards from Chapter 6 of Chemistry: The Molecular Nature of Matter and Change.

Thermodynamics

The study of energy and its transformations

Work and heat

Potential energy is converted to kinetic energy in these forms

System

The part of the universe a thermodynamic study focuses on

• Can include the contents of a flask or reaction

Surroundings

Everything outside a system that a thermodynamic study focuses on

• Includes a reaction’s flask and laboratory conditions

Internal energy (E)

The sum of a system’s potential and kinetic energies

Delta E

How a change in energy is denoted

• Equal to E_final - E_initial or q + w

• Described from a system’s perspective as the opposite of the change to the surroundings

Heat (q)

Energy transferred as a result of temperature differences until an equilibrium is released

• Flows out are exothermic

• Flows in are endothermic

Exothermic process

A process that releases heat to the surroundings and loses energy, such as a pot cooling

Endothermic process

A process that absorbs heat from the surroundings and gains energy, such as ice melting

Work (w)

Energy transferred when an object moves a force

• May be positive (having this done) or negative (doing this)

Negative work

How a system doing work on the surroundings is notated, as the system loses energy

Positive work

How a system having work done on it is notated, as the system absorbs energy

Law of conservation of energy (first law of thermodynamics)

This states that the total energy of the universe is constant and conserved

• ΔE_universe = ΔE_system + ΔE_surroundings = 0

Joule (J)

The SI unit of energy, equal to 1 (kg * m²)/s²

• Relates the force of an object to the distance the mass moves

calorie (cal)

An older unit of energy defined as the quantity of energy needed to raise 1 g of water by 1 degree C

• Now related as equal to 4.184 J

• Different from nutritional Calories

British thermal unit (Btu)

The quantity of energy required to raise 1 lb of water by 1 degree F, often used in appliances

• Equal to 1055 J

State function

A property dependent only on the current state of the system

• Includes the internal energy (E)

Pressure-volume work (PV work)

Work in the form of volume change given external pressure

• w = -PΔV

• For a positive ΔV, w becomes negative

• For a negative ΔV, w becomes positive

• Converted using the ratio 1 atm L = 101.3 J

Enthalpy (H)

A thermodynamic quantity defined as the sum of internal energy and pressure-volume at constant pressure

• Denoted as H = E + PV

Change in enthalpy (ΔH)

The heat released or absorbed of a system at constant pressure

• Change is denoted as ΔH = ΔE + PΔV

  • This can expand to q + w + PΔV

  • And substituting w = -PΔV, we get ΔH = q + w - w = q_P

  • Essentially equal to ΔE due to greater magnitude of heat change

Heat at constant pressure (q_P)

The heat released or absorbed in a reaction at constant pressure

• Roughly equal to ΔH and ΔE due to greater magnitude

• May be positive or negative depending on absorption or emittance

Enthalpy diagram

A diagram demonstrating the flow of energy in a enthalpy reaction

• Goes down as heat is released exothermically

• Goes up as heat is absorbed endothermically

Specific heat capacity (c)

The quantity of heat required to change the temperature of 1 gram of a substance by 1 K

• c = q / mΔT

• Can be rearranged to determine the heat absorbed or released, using q = mcΔT

• Very high for water, requiring much more energy than metals

Molar heat capacity (C_m)

The quantity of heat required to change the temperature of 1 mol of a substance by 1 K

• C_m = q / (mol ΔT)

• Expressed as J / mol K

Calorimeter

A device used to measure heat changes in a process

Constant-pressure calorimetry

Calorimetry based on the idea that heat released by the system is absorbed by the water (-q_system = q_water or q_soln = -q_rxn)

• Using q = mcΔT, one can solve for the variable needed, noting ΔT to determine sign

• Uses the same T_final in calculations for ΔT due to equilibrium

Constant-volume calorimetry

Used to measure the heat of combustion reactions, usually involving fuels or foods

• Uses the formula -q_reaction = q_water + q_calorimeter

  • q_calorimeter = C_calorimeter * ΔT_calorimeter if the calorimeter absorbs energy

  • C_calorimeter represents the heat capacity of the calorimeter

Heat released at constant volume (q_V)

What is measured in constant-volume calorimetry, representing the heat absorbed by the calorimeter

• Equal to ΔE

Thermochemical equation

A balanced equation that also includes the enthalpy change of the reaction, referring only to the specific amounts and states of matter

• Can be reversed or multiplied proportionally to obtain specific amounts of substance or heat

Direction (reversing a reaction inverts the sign of H) and magnitude (scaling by n results in a scaling of H)

What are the two operations for manipulating thermochemical equations?

Phase transitions have non-zero enthalpy changes, resulting in physical state changes

Why must states of matter be specified in a thermochemical equation?

• Endothermic: q > 0, ΔH > 0

• Exothermic: q < 0, ΔH < 0

How do the terms “heat absorbed” and “heat released” translate into algebraic signs for q_P and ΔH?

Hess’s law

The enthalpy change of a process is the sum of the changes of its steps

• Expressed as ΔH_overall = ΣΔH

• Used with equation manipulation

1. List out the steps needed with their known ΔH values.

2. Manipulate the equations.

3. Find ΣΔH for the overall equation and steps.

How do I utilize Hess’s Law?

Use fractional coefficients or reverse the equation.

How can I manipulate a chemical equation?

State function

A property that depends solely on a current or instantaneous state of a substance

Standard state

A standardized set of conditions to study and compare reactions

• Gases: 1 atm and ideal behavior

• Aqueous solutions: 1 M concentration

• Pure substance: Most stable form at 1 atm and the temperature of interest, usually 25°C

Standard enthalpy of reaction (ΔH°_rxn)

The enthalpy change of a reaction given standard states

Formation equation

An equation where 1 mol of compound forms from its elements

• Often scaled to only form 1 mol of product

Standard enthalpy of formation (ΔH°_f)

The enthalpy change for a formation equation given standard states

• Elements in standard state: ΔH°_f = 0

• Halogens: evaluated in molecular form, allowing ΔH°_f = 0

• Allotropes: varying elemental forms with only one with ΔH°_f = 0

• Most have negative values of ΔH°_f due to forming through exothermic reactions

1. Write the elements and reactants to form 1 mol of product

2. Ensure all elements are in their standard states

3. Balance the equations

4. Find ΔH°_f

How do I write a formation equation?

1. Note that elements must decompose, so take -ΣnΔH°_f from the reactants

2. Then, elements must form products, so take ΣnΔH°_f from the products and add it to above

3. Add ΣnΔH°_f(products) - ΣnΔH°_f(reactants)

How do I determine ΔH°_rxn from values of ΔH°_f?