Unit 5 Thermochemistry
(Tro textbook Chapter 6)
In this unit we will study energy in terms of chemical potential energy and how that energy can change form to accomplish work. We will focus specifically on heat and energy transfer.
Vocabulary for you to learn that will make this unit more understandable:
Energy (E) – the ability to do work or produce heat
Work- force acing over a distance (it always involves a transfer of energy)
1st Law of Thermodynamics- also known as The Law of Conservation of Energy. States that energy can neither be created nor destroyed; it can only be converted from one form to another. This means that the total amount of energy in the universe is constant.
Energy can be classified in two ways:
Potential energy – energy due to position or composition. In chemistry this is usually the energy stored in bonds.
Kinetic energy – energy due to the motion of an object, usually of particles; it is proportional to Kelvin temperature; kinetic energy depends on the mass and the velocity of the object: KE = ½ mv2
Heat (q) – involves a transfer of energy between two objects due to a temperature difference between the two objects. Heat energy “flows” from a warmer object to a cooler one, but, remember, temperature is not a measure of energy—it just reflects the motion of particles.
Temperature- a property that reflects random motions of the particles of a particular substance
Exothermic- reaction (or process) that releases heat energy
Energy flows OUT of the system. Potential energy is converted to thermal energy. The products have a lower potential energy that the reactancts.
Endothermic- reaction (or process) that absorbs heat energy
Energy flows INTO the system. Thermal energy is converted to potential energy. The products have a higher potential energy that the reactancts.
In considering a chemical reaction, the reaction is our “system”. The surroundings are everything else, including things like the container the reaction occurs in, the room it sits in, etc.
Internal energy (E) of a system is the sum of the kinetic and potential energies of all the particles in a system.
Thermodynamic quantities always consist of a number and a sign (+ or -). The sign represents the systems point of view. (Engineers will use the surroundings point of view)
Exothermic: -q (system’s energy is decreasing)
Endothermic: +q (system’s energy is increasing)
Example: Calculate the ΔE if q = -50 kJ and w = +35 kJ.
For a gas that expands or is compressed, work can be calculated by: w = -PΔV units of w = L*atm
1 L*atm = 101.325 J
Example: Calculate the work done if the volume of a gas is increased from 15 mL to 2.0 L at a constant pressure of 1.5 atm.
Enthalpy and Calorimetry
Enthalpy (ΔH) concerns the heat energy in a system.
ΔH = q at constant pressure only
At constant pressure, the terms “heat of reaction (Hrxn) and change in enthalpy (ΔH) are able to be used interchangeably.
Note: any reactions completed at atmospheric pressure (read: in an open container) are at constant P
We have 6 ways to calculate enthalpy.
1. Graphically
2. Stoichiometry
The change in enthalpy of a system can be calculated using:
ΔH = Hproducts - Hreactants
For an exothermic reaction, ΔH is negative
For an endothermic reaction, ΔH is positive
Example: For the reaction 2Na + 2H2O à 2 NaOH + H2 ΔH = -386 kJ
Calculate the heat change that occurs when 3.5 grams of Na reacts with excess water.
Example: Carbon monoxide burns in air to produce carbon dioxide according to the following balanced equation:
2 CO(g) + O2(g) à 2 CO2(g) + 566 kJ
How many grams of carbon monoxide are needed to yield 185 kJ of energy?
3. Hess’s Law
Hess’s Law- states that the change in enthalpy from products to reactants, ΔH, is the same whether the reaction occurs in one step or in several steps.
ΔH is NOT dependent on the reaction pathway
The sum of the ΔH for each step equals the ΔH for the total reaction
If a reaction is reversed, the sign of ΔH is reversed
If the coefficients in a reaction are multiplied by an integer, the value of ΔH is multiplied by the same integer.
Example: Given the following reactions and their respective enthalpy changes, calculate ΔH for the reaction: 2C(s) + H2(g) à C2H2(g)
C2H2 + 5/2 O2 à 2 CO2 + H2O ΔH = -1299.6 kJ/molC2H2
C + O2 à CO2 ΔH = -393.5 kJ/molC
H2 + ½ O2 à H2O ΔH = -285.9 kJ/molH2
Example: Heat of combustion of C to CO2 is -393.5 kJ/molCO2, whereas that for combustion of CO to CO2 is -283.0 kJ/molCO2. Calculate the heat of combustion of C to CO.
4. Standard Enthalpies of Formation
Standard enthalpy of formation (ΔHfo) – the change in enthalpy that accompanies 1 mole of a compound from its elements with all substances in their standard states at 25oC.
The degree sign on a thermodynamics function indicates that the process it represents has been carried out at standard state conditions.
Standard states
For gases, P = 1 atm
For a substance in solution, the concentration is 1M
For a pure substance in a condensed state (liquid or solid) the standard state is the pure liquid or solid
For an element, the standard state is the form under which the element exists under conditions of 1 atm and 25oC
Values of ΔHfo can be found in the textbook’s Appendix
ΔHorxn = ∑ ΔHfoproducts - ∑ ΔHforeactants
Example: The standard enthalpy change for the reaction CaCO3(s) à CaO + CO2 is 178.1 kJ. Calculate the ΔHfo of CaCO3.
- Constant Pressure Calorimetry (coffee cup calorimetry)
Terms to know:
Heat capacity – energy required to raise temp. by 1 degree (Joules/ oC)
Specific heat capacity (c) – same as above but specific to 1 gram of substance
Molar heat capacity -- same as above but specific to one mole of substance
(J/mol K or J/mol oC )
Energy (q) released or gained -- q = m·c·DT
q = quantity of heat ( Joules or calories) m = mass in grams ΔT = Tf - Ti (final – initial)
c = specific heat capacity ( J/goC)
Specific heat of water (liquid state) = 4.184 J/g°C ( or 1.00 cal/g oC)
Water has one of the highest specific heats known! That is why the earth stays at such an even temperature all year round! Cool huh?
Heat lost by substance = heat gained by water
(if this does not happen, calculate the heat capacity of the substance)
Units of Energy:
§ calorie--amount of heat needed to raise the temp. of 1.00 gram of water 1.00 C
§ kilocalorie--duh!; the food calorie with a capital C.
§ joule--SI unit of energy; 1 cal = 4.184 J
Calorimetry- the science of measuring heat flow in a chemical reaction
It is based on observing the temperature change when a body absorbs or discharges heat.
The instrument used to measure this change is the calorimeter.
Constant pressure calorimetry- pressure remains constant during the reaction or process
Simple calorimetry- used to determine heats of reaction (ΔHrxn)
The primary reaction to calculate heat changes in a system is q = mcΔT. (the Mcat equation)
The energy released as heat = (mass of solution) (heat capacity) (increase in temperature)
Example: A coffee cup calorimeter contains 150 grams H2O at 24.6oC. A 110 gram block of molybdenum is heated to 100oC and then placed in the water in the calorimeter. The contents of the calorimeter come to a temperature of 28.0oC. What is the heat capacity per gram of molybdenum?
6. Constant Volume Calorimetry (Bomb Calorimetry)
In terms of calorimetry, we can describe certain properties of the reaction as an:
Extensive property- this depends on the amount of substance (Ex. ΔHrxn)
Intensive property- doesn’t depend on the amount of substance (Ex. Temperature, specific heat)
Calorimetry can be done in a closed, rigid container. This is called constant volume calorimetry (bomb calorimeter). No work can be done since the volume doesn’t change.
Heat evolved = (change in temp) (Heat capacity of calorimeter) = ΔT Ccal
Specific Heat of the calorimeter (Ccal) – energy required to change the temperature inside the calorimeter 1oC
Example: Camphor (C10H16O) has a heat of combustion of 5903.6 kJ/mol. When a sample of camphor with mass of 0.1204 g is burned in a bomb calorimeter, the temperature increases by 2.28°C. Calculate the heat capacity of the calorimeter.
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