Thermochemistry

I. What is Thermochemistry?
a. This is the study of the transformations of energy, especially heat, during chemical reactions.
II. The Nature of Energy
a. Kinetic Energy: This is the energy related to motion. This can be calculated using the following equation:
i. Ek = 1/2mv2
1. Ek is the kinetic energy (J).
2. m is the mass of the object (g).
3. v is the velocity of the object (m/s).
b. Potential Energy: This is the energy of an object in relation to its position to other objects. This can be calculated using the following equation:
i. Ep = mgh
1. Ep is the potential energy (J).
2. m is the mass of the object (g).
3. g is the gravitational constant (9.8 m/s2).
4. h is the height of the object in relation to the other objects (m).
ii. Electrostatic Potential Energy: This arises from the interactions between charged particles. It is proportional to the charges on the two objects and inversely proportional to the distance between them. The equation that illustrates this relationship is as follows:
1. Eel = kQ1Q2 / d
a. Eel is the electrostatic potential energy
b. k is the constant of proportionality (8.99x109 J-m/C2)
c. Q1 and Q2 represent the electric charges on the two objects (C).
d. d is the distance between the two charged objects (m).
c. Units of Energy:
i. The SI unit of energy is the Joule, which is equal to 1kg-m2/s2.
ii. The calorie is also another common unit which is the amount of energy needed to increase the temperature of 1g of water 1°C.
1. 1 calorie = 4.184 Joules
d. System and Surroundings
i. A system is a specific amount of matter, and the surroundings are everything outside of the system.
ii. A closed system can exchange energy but not matter with the surroundings. Energy can be transferred between the system and the surroundings in the form of heat or work.
e. Transferring Energy; Work and Heat
i. Work is the amount of energy needed to move an object against a force over a distance.
ii. A force is a push or a pull on an object.

iii. Work can be calculated using the following equation:
1. w = Fd
a. w is work.
b. F is the force.
c. d is the distance (m).
iv. Heat is the energy transferred from a hotter object to a colder object.
v. Energy is the amount of work it takes to transfer heat.
III. The First Law of Thermodynamics: This states that energy is conserved.
a. Internal Energy: This is the sum of all potential and kinetic energies of a system. The change in internal energy can be calculated by the following equation:
i. ΔE = Efinal – Einitial
b. A positive value of ΔE indicated a gain in energy from its surroundings, and a negative value indicates a loss of energy to its surroundings.
c. Relating ΔE to heat and work
i. The change in the internal energy of a system equals the sum of the heat transferred and the work done on or by the system, resulting in the following equation:
1. ΔE = q + w
a. q is the heat.
b. w is the work.
ii. Both q and w have signs that indicate the direction of an energy transfer. When heat is transferred from the surrounds to the system q > 0, and when the surroundings do work on the system w > 0.
d. State Functions
i. Internal energy is a state function which depends on the state or condition of the system, and not the path by which it took.
ii. Temperature of a system is a state function.
iii. Heat and work are not state functions because their values are dependent on the way in which a system changes its state.
IV. Enthalpy and Enthalpies of a Reaction
a. Pressure-Volume work occurs within a system when a gas is consumed or produced within a chemical reaction at a constant pressure.
i. w = -P ΔV
1. P is the pressure (atm).
2. ΔV is the change in volume (L).
ii. For this reason, enthalpy, which accounts for heat flow and chemical changes, is also a state function.
b. Enthalpy is equivalent to the internal energy plus the product of the pressure and volume of a system. This can be calculated with the following equation:
i. H = E + PV
1. H is the enthalpy.
2. E is the internal energy
c. If a change occurs at a constant pressure then the equation becomes:
i. ΔH = ΔE + P ΔV
d. For most reactions the difference between ΔH and ΔE is small because the value for PΔV is negligible.
i. If ΔH is positive then the reaction is endothermic. If ΔH is negative then the reaction is exothermic.
e. The enthalpy of a reaction is the enthalpy of the products minus the enthalpy of the reactants, resulting in the following equation:
i. ΔH = H(products) – H(reactants)
f. Balanced chemical equations associated with an enthalpy change are called thermochemical equations.
g. Enthalpy is an extensive property therefore the enthalpy of a reaction is proportional to the amount of reactants.
h. Reversing a reaction changes the sign of ΔH.
i. The enthalpy of a reaction is dependent upon the physical states of the reactants and products.
V. Calorimetry: This is the measurement of heat flow. A calorimeter is used to measure the heat flow.
a. Heat Capacity and Specific Heat
i. The temperature change of a calorimeter is dependent upon the heat capacity which is the amount of heat required to raise its temperature by 1K.
ii. The molar heat capacity is the heat capacity for one mole of a pure substance.
iii. The heat capacity for 1g of a substance is specific heat.
iv. The heat gained or lost can be calculated using the following equation:
1. q = (specific heat) x (grams of a substance) x (ΔT)
b. Constant-Pressure Calorimetry
i. The heat gained by a solution (qsoln) is equal in magnitude but opposite in sign from the heat of the reaction (qrxn). This is represented by the following equation:
1. qsoln = (specific heat of solution) x (g of solution) x (ΔT) = -qrxn
ii. Bomb Calorimetry
1. Constant volume Calorimetry is performed in a bomb calorimeter which is used to measure the heat transferred in combustion reactions. This is represented by the following equation:
a. qrxn = (-Ccal) x (ΔT)
i. Ccal is the heat capacity of the calorimeter (J/°C).
ii. ΔT is the change in temperature (K)..
VI. Hess’s Law: This law states that if a reaction can be described in multiple steps, then ΔH for the overall reaction is a sum of the ΔH values for the various steps.
VII. Enthalpies of Formation
a. Enthalpy Changes
i. The enthalpy of formation (ΔHf) is the enthalpy change for the reaction in which a more complex substance is formed from its original components.
ii. The standard enthalpy of a reaction (ΔH°) is the enthalpy change when all reactants and products are at 1atm and 289K (25°C).
iii. The standard enthalpy of formation (ΔH°f) is the enthalpy change for the reaction that forms 1 mole of a substance at 1atm and 298K.
iv. The most stable state for an element is one in which it is at 298K, 1atm, and ΔH°f is 0.
b. Using Enthalpies of Formation to Calculate Enthalpies of a Reaction
i. The standard heat of reaction is equal to the sum of the standard heats of formation of the products minus the standard heats of formation of the reactants, resulting in the following equation:
1. ΔH°rxn = Σn ΔH°f (products) - Σm ΔH°f (reactants)
a. n and m are variables that represent the coefficients of the chemical equation.
VIII. Foods and Fuels
a. The fuel value of a substance is the heat released when 1g of the substance is combusted.
b. The most common fuels are hydrocarbons (fossil fuels) such as natural gas, petroleum, and coal.
i. Coal is the most abundant fossil fuel.
1. The sulfur within it causes air pollution. Therefore, coal gasification is used to purify this source of energy.
c. Renewable energies, which are unlimited, include solar energy, wind energy, biomass, and hydroelectric energy.