1Define energy.
Energy is the capacity to do work or produce heat.
2Define kinetic energy.
Kinetic energy is energy due to motion.
3Define internal energy.
Internal energy is the sum of all potential and kinetic energy within a system.
4What is the SI unit of energy?
The joule (J).
5What is another common unit of energy besides the joule?
The calorie (cal).
61 calorie equals how many joules?
4.184 joules.
7State the law of conservation of energy.
Energy cannot be created or destroyed; it can only be converted from one form to another.
8Define a thermodynamic system.
A system is the specific part of the universe being studied.
9Define surroundings.
Surroundings are everything outside the system that can exchange energy with it.
10Define the heat of reaction.
The heat of reaction is the amount of heat absorbed or released during a chemical reaction.
11What is an exothermic process?
A process that releases heat to the surroundings.
12What is an endothermic process?
A process that absorbs heat from the surroundings.
13Define enthalpy (H).
Enthalpy is the heat content of a system at constant pressure.
14What does ΔH represent?
ΔH is the enthalpy change, or heat absorbed or released, at constant pressure.
15What happens to ΔH when a thermochemical equation is reversed?
The sign of ΔH changes.
16What happens to ΔH when a thermochemical equation is multiplied by a coefficient?
ΔH is multiplied by the same coefficient.
17What formula relates internal energy, heat, and work?
ΔE = q + w.
18What formula is used to calculate heat with temperature change?
q = m × s × ΔT.
19What does “m” represent in q = m × s × ΔT?
Mass of the substance.
20What does “s” represent in q = m × s × ΔT?
Specific heat capacity.
21What does “ΔT” represent in q = m × s × ΔT?
Change in temperature (Tfinal – Tinitial).
22Define heat capacity.
The amount of heat required to raise the temperature of an object by 1°C.
23Define specific heat capacity.
The amount of heat required to raise the temperature of 1 gram of a substance by 1°C.
24How is heat absorbed or evolved related to enthalpy of reaction?
q = n × ΔHrxn.
25What is the relationship between qsystem and qsurroundings in calorimetry?
qsystem = −qsurroundings.
26What is Hess’s Law?
The total enthalpy change for a reaction is the sum of the enthalpy changes for individual steps.
27How do you calculate ΔH°rxn using enthalpies of formation?
ΔH°rxn = Σ n ΔH°f(products) − Σ m ΔH°f(reactants).
28What type of process has q > 0?
Endothermic.
29What type of process has q < 0?
Exothermic.
30If the temperature of water in a calorimeter increases, what type of reaction occurred?
An exothermic reaction.
31If the temperature of water in a calorimeter decreases, what type of reaction occurred?
An endothermic reaction.
32What is the formula for work in thermodynamics?
w = −PΔV.
33In a reaction, if the system expands against pressure, is work positive or negative?
Negative; the system does work on the surroundings.
34In a reaction, if the system is compressed, is work positive or negative?
Positive; work is done on the system.
35What does it mean when ΔHrxn is negative?
The reaction releases heat (exothermic).
36What does it mean when ΔHrxn is positive?
The reaction absorbs heat (endothermic).
37What is standard enthalpy of formation (ΔH°f)?
The enthalpy change when one mole of a compound is formed from its elements in their standard states.
38What are the standard states of elements?
The most stable form of the element at 1 atm and 25°C.
39What equation is used in calorimetry problems involving specific heat?
q = m × s × ΔT.
40What are the units of specific heat capacity?
J/(g·°C).
41What does a calorimeter measure?
It measures the heat absorbed or released during a chemical or physical process.
42What is the enthalpy of reaction when multiple reactions are combined?
It is the sum of the enthalpy changes for each individual reaction.
43How is the heat of reaction expressed in a balanced chemical equation?
As ΔH next to the equation, indicating energy released or absorbed per mole of reaction.