Essential Definitions

System

• System is the part of universe under study

• Everything else is called surroundings

• Open systems transfer energy and matter to surroundings

• Closed system can transfer energy but not matter

• Isolated system has no energy or matter transfer

State Functions

• State functions describes an aspect of a chemical system

  • Pressure, volume, temperature, and moles are state functions in the ideal gas law

• Thermodynamics state functions can be internal energy (E), enthalpy (H), entropy (S), and Gibbs free energy (G)

• Changes in state function depend on the initial and final state of a system

• Not state functions are heat (q) and work (w) because they rely on the steps used to transform matter from initial to final

Standard State

• Extensive properties of matter change as the amount of sample changes

  • Such as ΔH, ΔS, ΔG, and ΔE

• To make them intensive properties, the temperature, pressute, mass, and physical state must be defined

• Standard state is 1 atm, 25 celsius, and 1 mol of compound present

  • This gives them a superscript of zero to note that they are at standard state

  • ΔH°, ΔS°, ΔG°, ΔE°

Exo- and Endo- Prefixes and Sign Conventions

• Exo-, or exothermic, means energy is being lost from the system to the surroundings

  • Negative sign for quantities

  • ΔH is negative if exothermic

• Endo-, or endothermic, means energy is being gained from surroundings

  • absorbs heat energy

  • ΔH is positive

Types of Energy

• Energy can be heat, light, chemical, nuclear, electrical, or mechanical

• Law of conservation of energy states energy cannot be created or destroyed. All forms of energy can be converted to heat energy

• Can be called kinetic, KE, or potential, PE, energy

  • Kinetic is the energy matter has because of its motion

  • potential energy is stored and released under certain conditions

    • gravitational energy is PE

    • electrostatic attraction between oppositely charge ions is PE

Measurement of Energy

Specific Heat

• Calorie is the amount of heat needed to raise the temperature of 1 gram of something by 1 degree celsius. 1 calorie = 4.184 joules

  • The specific heat of water is the energy needed to raise 1 gram of water by 1 degree celsius = 4.184 J/g*C

• Specific heat equation

  • q = sp.ht x g x ΔT

  • heat energy = (specific heat)(mass in grams)(temperature change)

First Law of Thermodynamics

• First law of thermodynamics states energy is always conserved

• ΔE = q + w

  • change in energy = heat + work

  • if q is positive, heat is added to the system

  • if w is positive, work is done on the system

Work

• Work is the force applied to an object as it moves a certain distance

  • work = force x distance moved

• Can also be found over a given area

  • work = pressure x area x volume change

Definition of qp, qv, ΔE, and ΔH

• First law of thermodynamics

  • ΔE = qp - PΔV

  • heat, q, have the subscipt, p, to indicate the pressure must be held constant

• When heat energy is measure in a calorimeter that does not allow the volume to change, PΔV is zero, resulting in ΔE = qv, where the volume is held constant

  • Calorimeters that do not let volume change are called bomb calorimeters

• Enthalpy change is ΔH where H is the heat content of the substance and the Δ is the difference

  • ΔH = Hproducts - Hreactants

• ΔE = ΔH - PΔV

Standard Enthalpy Changes and the Standard Heat of Reaction

• ΔH is an extensive property and to make it intensive, the amount of chemical reacting must be specified

• Standard heat of reaction, ΔH°, is the heat produced when the number of moles in the balanced equation reacts. in kJ.

• If a formation reaction occurs, ΔH°f, where only one mole of product is the equation, the units will be kJ/mol

Hess’s Law

• When converting coefficients in a balanced equation by multiplying, dividing, or any mathematic conversion, that same conversion must be done to the heat of reaction

• Summarized law

  • If the coefficients are multiplied by a constant, that ΔH°reaction is multiplied by the same constant

  • If two or more equations are added to create an overall reaction, the heats of the equations can be added to obtain the overall reaction heat

• When trying to use Hess’s Law, follow the principles:

  • Focus on complex molecules first

  • focus on atoms and molecules that occur in just one reaction

  • Focus on atom and molecules that are in the overall reaction

  • Focus on finding atoms and molecules to cancel out unneeded ones from equations

Formation Reactions and Heats of Formation

• Formation reactions are when reactants are in their standard state and there is only one mole of product. Examples:

  • Fe (s) + ½O2 (g) → FeO (s)

  • 2K (s) + ½H2 + 2O2 (g) + P (s) → Fe2O3 (s)

• Heat of formation for elements is always zero

Entropy and the Second Law of Thermodynamics

• Entropy is the number of different ways a system can arrange the particles

  • Increasing number of particles increases entropy

  • Increasing volume increases entropy

  • Changing state (liquid to gas) increasing entropy

• Ice has the least entropy while gaseous water has the most

• Entropy is assigned the symbol S with the units J/C

• Standard entropy is S° with the units J/C mol

• Entropy principles

  • Gas formation increases entropy greatly. A higher Δng means a greater entropy

  • If Δng = 0, solid to liquid phase is a high contributor to entropy increase

    • solid formation decreases entropy

  • increase in temperature increases entropy

Gibbs Free Energy, ΔG°

• Free energy change, ΔG°, is the max amount of energry available from any chemical reaction

  • Driven by entropy and enthalpy values

• ΔG° = ΔH° - TΔS°

• Second law of thermodynamics is any physical or chemical change must result in the increase of entropy in the universe

Free Energy and Equilibrium

• When not at standard state, ΔG is used not ΔG°

• ΔG = ΔG° + RT ln(Q)

  • Q is reaction quotient

  • R is 8.314 J/mol K

• ΔG° = -RT ln(K)

  • K is equilibrium constant

• If K > 1, reaction will move forward

• If K < 1, reaction will move reverse