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General Chemistry Ch. 6 - Thermochemistry

Thermodynamics The study of heat (thermal energy) and its transformations.
Thermochemistry The branch of thermodynamics that deals with the head involved in chemical and physical change, and especially with the concept of enthalpy.
When energy is transferred from one object to another, it appears as Work and/or heat
System The part of the universe that we’re focusing on
Surroundings Everything else relevant to the change, except the system
Internal energy, E Each particle in a system has potential energy and kinetic energy, and the sum of these energies for all particles in the system is the internal energy, E.
Calculation for change in (delta) E Delta E = E_final – E_initial = E_products – E_reactants.
Because the total energy must be conserved, a change in the energy of the system is always accompanied by… An opposite change in the energy of the surroundings
Relationship between E_final and E_inital if: lost energy to surroundings E_final < E_inital; delta E < 0
Relationship between E_final and E_inital if: gained energy from surroundings E_final > E_initial; delta E > 0
Heat (symbol q); the energy transferred between a system and its surroundings as a result of a difference in their temperatures only
All other forms of energy transfer (mechanical, electrical, chemical, and so on) other than heat Work (symbol w); the energy transferred when an object is moved by force
Calculation for delta E involving all energies involved. What does w equal? Delta E = q + w. w = -PdeltaV. <- note: it’s negative, which means work is being taken out of the equation.
How are the signs of q & w derived? Energy coming into the system is positive, energy going out of the system is negative
Is delta E positive or negative: the system is a sample of hot water and the surroundings consist of a room at room temperature. What if rather than hot water, it was ice water? Heat flows out from the system until the temperature is equivalent to that of the surroundings; q is negative, w is zero, delta E is negative. If ice: heat would flow into the system and q would be positive, thus delta E = positive
Pressure-volume work The type of work in which a volume changes against external pressure (e.g. piston movement).
Reaction between Zinc and hydrochloric acid in a container attached to a piston-cylinder assembly: positive or negative delta E? What if external pressure was increased on the piston? After the reaction H2 becomes a gas and the volume increases. The energy was lost by the system and converted into pressure-volume work (PV work), thus delta E is negative. If external pressure is increased, E would be positive.
First law of thermodynamics The total energy of the universe is constant. This law is also known as the law of conservation of energy
E_system + E_surroundings = … 0
Standard unit of energy Joule (kg*m^2/s^2)
1 cal = Used to be defined as the quantity of energy required to raise the temperature of one g of water by 1deg C. Now 1 cal = 4.184 J
1kJ = how many cals? 1 kJ = 1000 J = .2390 kcal = 239 cal
Nutritional Calorie Expressed with a capital C; it is equal to one kcal.
State function A property dependent only on the current state of the system, not the path taken to reach that state. The current state = the difference between the final and initial states. Pressure (P) and Volume (V) are also state functions.
What are some non-state functions? Heat and work. Why? Because depending on HOW the change took place, there can be a different distribution of energy lost as heat and energy lost as work. Whereas the energy lost will be the same regardless.
Equation for w w = -PdeltaV
Change in enthalpy equation. How can you use this equation to prove E only refers to heat, not work? deltaH = deltaE + PdeltaV. If you solve for E: deltaE = deltaH – PdeltaV
In verbal summary, the change in enthalpy equals… The heat gained or lost at constant temperature.
Three cases where all of the energy changes occur as transfers of heat 1. Reactions that do not involve gasses; 2. Reactions in which the amount (mol) of gas does not change; and 3. Reactions in which the amount (mol) of gas does change.
In cases where the amount (mol) of gas does change, how is the energy change purely an occurrence of a transfer of heat? Isn’t volume changing? Volume is changing, but q is so much larger than PdeltaV that deltaH is very close to deltaE. Since the great majority of the transfer is due to heat it’s accepted that deltaH = deltaE in this case.
Heat of reaction Always refers to H_final – H_initial. If heat is released to the surroundings, deltaH is negative. If it is absorbed, deltaH is positive.
Exothermic process deltaH is negative. H_final is less than H_initial. Energy is released to the surroundings.
Endothermic process Absorbs heat and results in an increase in enthalpy of the system.
Heat of formation When 1 mol of a compound is produced from its elements, the heat of reaction is called the heat of formation (deltaH_f)
Heat of fusion When 1 mol of a substance melts, the enthalpy change is called the heat of fusion (deltaH_fus)
Heat of vaporization When 1 mol of a substance vaporizes, the enthalpy change is called the heat of vaporization (deltaH_vap)
How do you measure enthalpy if you have no absolute measurement system? Enthalpy can only be measured if there is a change. You can’t say something has “this much enthalpy”. You can only state the enthalpy of a reaction.
Heat capacity Every object has its own heat capacity: the quantity of heat required to change its temperature by 1 K. Heat capacity = q/deltaT
Specific heat capacity The quantity of heat required to change the temperature of 1 gram of a substance by 1 K. specific heat capacity (c) = q/(mass * deltaT)
Molar heat capacity (C) Note: capital C instead of lowercase (which is specific heat). Molar heat is the quantity of heat required to change the temperature of 1 mole of substance by 1 K: C = q/(moles * deltaT)
Calorimeter A device used to measure the heat released/absorbed by a physical or chemical process. The apparatus is the “surroundings” that changes temperature when heat is transferred to or from the system.
Two common types of calorimeters The constant pressure and constant volume calorimeters.
Constant pressure calorimetry E.g. a coffee-cup calorimeter, already know from lab.
c of solid added to constant-pressure calorimeter c = -[(c_h2o x mass_h2o x deltaT_h2o)/(mass_solid x deltaT_solid)]
Problem with constant pressure calorimeters In the coffee-cup calorimeter, we assume all heat is gained by water, but some must be gained by the thermometer, stirrer, etc.
Constant volume calorimetry More accurate than constant pressure calorimeters. The heat capacity of the ENTIRE calorimeter must be known. E.g. the bomb calorimeter.
Thermochemical equation A balanced equation that includes the heat of reaction. The deltaH_rxn that is shown refers to the amounts (moles) of substances and their states of matter in that specific equation.
The magnitude of deltaH in a thermochemical equation depends on The amount and physical state of the substance reacting, and the deltaH per mole of substance.
The sign of deltaH depends on Whether it’s exothermic or endothermic. Also, reverse reactions have the opposite signs as forward reactions.
Hess’ law of heat summation The enthalpy change of an overall process is the sum of the enthalpy changes of its individual steps.
Standard states of gasses, substances in aqueous solutions, and pure substances Gas: 1 atm with the gas behaving ideally; aqueous solution: 1 M concentration; pure substance: 1 atm and the temperature of interest, usually 25deg Celsius
Standard heat of reaction (deltaH(deg sign)_rxn) When the heat of reaction (deltaH_rxn) has been measured with all the reactants and products in their standard states, it is referred to as the standard heat of reaction.
Formation equation 1 mole of a compound forms from its elements
Standard heat of formation The enthalpy change for the formation when all the substances are in their standard states.
Heats of formations of elements in their standard states 0
Heats of formations of most compounds Negative; they’re exothermic. Note: MOST, not all.
How to use heat of formation to determine overall heat of reaction Sum of the heat of formation for all products (including multiplying by coefficients) minus the sum of all reactants.
Created by: Intellex_



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