Heat, Temperature, and Enthalpy

• Energy is transferred between samples via electric current, light beams, moving pistons, and flames
• All energy is measured in joules (J), including heat
• Energy is neither created nor destroyed; energy transferred between samples is equal

Heat

• Heat is the energy transferred due to temperature differences
• Heat units are joules (J)

Temperature

• Temperature measures how hot or cold something is and depends on the average kinetic energy of particles in an object
• Higher temperatures equate to faster particle movement
• Temperature increases are affected by the sample's mass

Heat vs Temperature

• Heat is energy transferred if samples at varying temperatures come into contact
• Temperature is intensive meaning it does not depend on the amount of a sample
• Heat is extensive, dependent on the amount of sample

Enthalpy

• Enthalpy (H) measures the energy content of a sample
• It's impossible to measure the total energy of matter, only energy content changes
• If pressure is constant, enthalpy increase equals the heat energy received

Enthalpy of Particles

• Enthalpy accounts for the total kinetic energy of the particles in a substance
• Cooling a gold ring causes atoms to move less, decreasing kinetic energy, resulting in an enthalpy decrease
• Absolute zero (0.00 K) doesn't stop all atomic motion, atoms still have minimal total and average kinetic energies that they can have

Molar Heat Capacity

• The molar heat capacity measures the heat required to increase the temperature of 1 mol of a substance by 1 K, having the symbol C, with units J/K•mol
• Molar heat capacity is only accurately measured if no other process like a chemical reaction occurs
• Relate heat (q, in joules) to molar heat capacity: q = nC∆T, where n is the quantity in moles, and ∆T is the temperature change

Molar Heat Capacity Examples

• The molar heat capacity of solid metals have approximately the same molar heat capacities
• The temperature of 1 mol of any solid metal will rise 1 K when it absorbs about 25 J of heat
• M(g/mol) × cp (J/K•g) = C (J/K•mol), where M is molar mass and cp is specific heat

Heat and Motion

• Heating a substance increases entropy and kinetic energy
• Heat produces disorderly particle motion
• Other energy types produce orderly motion or positioning

Molar Enthalpy Change

• Molar enthalpy change occurs heating an iron horseshoe
• Reaction, Fe(s, 300 K) → Fe(s, 1100 K) ∆H = 20.1 kJ/mol, heating 1 mol of solid iron from 27°C to 827°C, increases molar enthalpy by 20 100 joules.
• Only heating or cooling relates molar heat capacity to molar enthalpy change, molar enthalpy change = C∆T

Enthalpy Change Example

• Sample problem: Ice warms from -5.4°C to -0.2°C, Tinitial = -5.4°C = 267.8 K, Tfinal = -0.2°C = 273.0 K, for H2O(s), C = 37.4 J/K•mol
• Ice warms 5.2 K, so ∆H = (37.4 J/K•mol)(5.2 K) = 1.9 × 102 J/mol. The molar enthalpy change is 0.19 kJ/mol

Types of Processes

• Endothermic process heating requires energy = positive enthalpy change
• Exothermic process cooling releases energy = negative enthalpy change

Thermodynamics

• Thermodynamics examines several processes and energy changes accompanying them
• It's possible to measure chemical equations with changes of state, for example, hydrogen gas and bromine liquid react, energy transfers out as heat because the enthalpy 2HBr is less than the enthalpy of H2 and Br2

Calorimetry

• Calorimetry, measures enthalpy changes during reactions in a calorimeter
• Reactions with +∆H are endothermic, while -∆H reactions are exothermic

Hess's Law

• It is possible to use Hess's Law to calculate enthalpy changes
• The overall enthalpy change in a reaction equals the sum of enthalpy changes for each step like the process P4(s) + 10Cl2(g) → 4PCl5(g)

Hess's Law and Algebra

• Chemical equations follow algebra rules for desired outcomes, so equations are added or subtracted, enthalpy changes follow

Standard Enthalpies of Formation

• Formula M (g/mol) × cp (J/K•g) = C (J/K•mol)
• Enthalpy change forming 1 mol of substance from elements in standard states is a standard enthalpy of formation (∆H)
• Reaction entropy is ∆Sreaction = Sproducts – Sreactants

Standard Thermodynamic Temperature

• The standard thermodynamic temperature to standardize the enthalpies of reaction is 25.00°C or 298.15 K

Entropy

• Entropy (S), measures disorder, with units J/K; processes tend to occur with increased entropy (+∆S), measured via the formula ∆Sreaction = Sproducts – Sreactants
• Entropy increases: dispersal through diffusion, dilution in solutions, pressure reduction in gases, temperature increases (not as significant unless phase change occurs), and reactions creating more product moles, more particles, or more gas particles

Gibbs Energy

• To calculate the Gibbs energy, use this equation G = H – TS
• Gibbs Energy helps determine spontaneity; spontaneous reactions occur without continuous outside assistance, having a negative Gibbs energy change
• Nonspontaneous reactions require continuous assistance, with a Gibbs energy change greater than 0

Influence of Temperature

• Spontaneity may be affected by temperature in the equation AG = ∆H – TAS
• Four outcomes of using different values of enthalpy and entropy:
  • Negative enthalpy/​Positive entropy = Spontaneous at all temperatures.
  • Negative enthalpy/​Negative entropy = Spontaneous only if T < ∆H/∆S.
  • Positive enthalpy/​Positive entropy = Spontaneous only if T > ∆H/∆S.
  • Positive enthalpy/​Negative entropy = Not spontaneous at any temperature.