General Chemistry Ch. 6 - Thermochemistry
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| Thermodynamics | The study of heat (thermal energy) and its transformations.
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| Thermochemistry | The branch of thermodynamics that deals with the head involved in chemical and physical change, and especially with the concept of enthalpy.
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| When energy is transferred from one object to another, it appears as | Work and/or heat
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| System | The part of the universe that we’re focusing on
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| Surroundings | Everything else relevant to the change, except the system
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| 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.
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| Calculation for change in (delta) E | Delta E = E_final – E_initial = E_products – E_reactants.
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| 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
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| Relationship between E_final and E_inital if: lost energy to surroundings | E_final < E_inital; delta E < 0
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| Relationship between E_final and E_inital if: gained energy from surroundings | E_final > E_initial; delta E > 0
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| Heat | (symbol q); the energy transferred between a system and its surroundings as a result of a difference in their temperatures only
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| 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
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| 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.
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| How are the signs of q & w derived? | Energy coming into the system is positive, energy going out of the system is negative
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| 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
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| Pressure-volume work | The type of work in which a volume changes against external pressure (e.g. piston movement).
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| 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.
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| First law of thermodynamics | The total energy of the universe is constant. This law is also known as the law of conservation of energy
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| E_system + E_surroundings = … | 0
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| Standard unit of energy | Joule (kg*m^2/s^2)
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| 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
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| 1kJ = how many cals? | 1 kJ = 1000 J = .2390 kcal = 239 cal
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| Nutritional Calorie | Expressed with a capital C; it is equal to one kcal.
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| 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.
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| 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.
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| Equation for w | w = -PdeltaV
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| 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
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| In verbal summary, the change in enthalpy equals… | The heat gained or lost at constant temperature.
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| 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.
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| 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.
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| 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.
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| Exothermic process | deltaH is negative. H_final is less than H_initial. Energy is released to the surroundings.
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| Endothermic process | Absorbs heat and results in an increase in enthalpy of the system.
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| 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)
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| Heat of fusion | When 1 mol of a substance melts, the enthalpy change is called the heat of fusion (deltaH_fus)
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| Heat of vaporization | When 1 mol of a substance vaporizes, the enthalpy change is called the heat of vaporization (deltaH_vap)
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| 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.
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| 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
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| 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)
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| 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)
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| 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.
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| Two common types of calorimeters | The constant pressure and constant volume calorimeters.
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| Constant pressure calorimetry | E.g. a coffee-cup calorimeter, already know from lab.
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| c of solid added to constant-pressure calorimeter | c = -[(c_h2o x mass_h2o x deltaT_h2o)/(mass_solid x deltaT_solid)]
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| 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.
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| Constant volume calorimetry | More accurate than constant pressure calorimeters. The heat capacity of the ENTIRE calorimeter must be known. E.g. the bomb calorimeter.
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| 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.
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| 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.
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| The sign of deltaH depends on | Whether it’s exothermic or endothermic. Also, reverse reactions have the opposite signs as forward reactions.
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| Hess’ law of heat summation | The enthalpy change of an overall process is the sum of the enthalpy changes of its individual steps.
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| 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
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| 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.
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| Formation equation | 1 mole of a compound forms from its elements
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| Standard heat of formation | The enthalpy change for the formation when all the substances are in their standard states.
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| Heats of formations of elements in their standard states | 0
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| Heats of formations of most compounds | Negative; they’re exothermic. Note: MOST, not all.
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| 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.
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