# Unit 6: Thermodynamics

## Heat and Temperature

• Temperature is the average amount of kinetic energy based on molecular motion of a substance

• Heat is the flow of energy between two substances at different temperatures

• The first law of thermodynamics states energy cannot be created or destroyed and can only transfer or change forms. Energy transfers through molecular collisions.

• Exothermic is when energy is transferred from the system to the surroundings, such as when a chemical reaction occurs in a beaker and the beaker feels hot. Endothermic reactions occur when a system absorbs the energy from the surroundings, such as when a beaker feels cold.

## Enthalpy

### Enthalpy Change, ΔH

• Enthalpy change is the measure of energy that is released or absorbed when bonds are broken/formed

• When bonds are formed, energy is released. When bonds are broken, energy is absorbed.

• If more energy is released than absorbed in a reaction, then it is exothermic and has negative enthalpy change. If more energy is absorbed than released in a reaction, then it is endothermic and has a positive enthalpy change.

## Energy Diagrams

### Exothermic and Endothermic Reactions

• When products have a lower energy than the reactants, the ΔH is negative and it is an exothermic reaction

• When reactants have a lower energy than the products, the ΔH is positive and it is an endothermic reaction

## Catalyst and Energy Diagrams

• Catalysts speed up a reaction by lowering the required activation energy and only changes the activation energy, not the amount of energy of products or reactants in the end or beginning (ΔH).

• This change can only occur when a reaction is not in equilibrium conditions

## Enthalpy of Formation, ΔHf°

• Enthalpy of formation, or heat of formation, is the energy change when one mole of a compound is formed from pure elements under standard conditions (298°K, 1 atm)

• Example: C(s) + 1/2 O2 (g) + 2 H2 (g) → CH3OH

• There is exactly one mole of a product, methanol, which has an odd number of elements, and diatomic in nature.

• ΔHf° of a pure element is defined as zero and stands true for elements that are diatomic in pure states, such as hydrogen or oxygen.

• If ΔHf° of a compound is negative, energy is released when the compound is formed from pure elements, is exothermic, and is more stable than the reactants.

• If ΔHf° is negative, the opposite is true, and it is endothermic.

• ΔH can be found if ΔHf° of product and reactants are known

• ΣΔHf° Products - ΣΔHf° Reactants = ΔH

• The units for reactions enthalpy is kj/mol.

## Bond Energy

• Bond energy is the energy required to break a bond and is always a positive number. When a bond is formed, the energy of formation is equal to the bond energy released.

• ΔH = Σ Bond energy of bond broken - Σ bond energy of bonds formed

• Bonds broken are the reactant bonds and the bonds formed are the product bonds

• The number of bonds broken is determined by how many moles of that substance there is and how many individual bonds there are.

• For 2 H2O molecules, there are two H-O bonds and two of the H2O molecules. So, there are four H-O bonds and to find the bond energy, multiply four by the bond energy of H-O (463 kJ/mol).

## Hess’s Law

• Hess’s Law states that if a reaction can be broken into a series of steps, the sum of the ΔH of those steps are equal to the ΔH of the reaction.

• When manipulating equations for enthalpy calculations…

• If a reaction is flipped, flip the enthalpy value (pos to neg, or neg to pos)

• If you multiply or divide the coefficients of a reaction, multiply or divide the enthalpy value by that same interval

• Adding up the enthalpy values of steps will be the new enthalpy value for the new reaction created from those steps

## Enthalpy of Solution

• When an ionic substance dissolves in water, energy is absorbed because bonds are breaking, but energy is also emitted because new dipoles are forming.

• Step one of dissolving: Bonds are broken in the solute. This step has a positive ΔH because bonds are being broken

• Step two: solvent molecules are separated. The water molecules are spread apart to make room for the solvent ions. This weakens intermolecular bonds, resulting in a positive ΔH

• Step three: New attractions are created. The free-floating ions are attracted to the dipoles in the water molecules, resulting in a negative ΔH despite no bonds being formed.

• Step two and three values are combined to form the hydration energy which is always a negative number. Hydration energy is a Coulombic energy and it’s magnitude increases as the ions increase their charge or decrease the size.

• If adding all three steps together, the enthalpy of solution can be determined.

## Thermodynamics of Phase Change

### Naming the Phase Changes

• Solid to liquid: melting

• Liquid to solid: freezing

• Liquid to gas: vaporization

• Gas to liquid: condensation

• Solid to gas: sublimation

• Gas to solid: deposition.

• Vapor pressure is the amount of pressure given off by molecules as they go from solid/liquid to gas.

### Enthalpy of Fusion

• Enthalpy of fusion is the amount of energy required to cause a solid to melt. Heat of fusion is how much heat given off when a substance freezes. The intermolecular forces in a solid substance are more stable when they are solid so they have a lower energy than liquid forces, so energy is released in the freezing process.

### Enthalpy of Vaporization

• Enthalpy of vaporization is how much energy is required to turn a liquid to a gas. The heat of vaporization is the energy given off when a gas condenses. Intermolecular forces are more stable in a liquid state, so energy is released as the forces stabilize and get stronger.

• When heat is added to a substance, it will either phase change or increase in temperature, not both. Only one can happen at a time. Temperature remains constant during a phase change.

## Calorimetry

• Calorimetry is the measurement of heat changes during a chemical reaction

• The specific heat of substance is how much heat is required to raise one gram of a substance by one degree celsius (or one degree kelvin). The higher or larger a specific heat is, the more heat a substance can absorb without changing temperature. A low specific heat means less heat is required to change the temperature of a substance.

• Specific heat equation is q = mcΔT

• q is heat added (J or cal)

• m is the mass of the substance (g or kg)

• c is specific heat

• ΔT is temperature change (K or C