Unit 7: Thermochemistry Flashcards
define energy
the capacity to do work or transfer heat
what is the equation for change in energy (delta energy)
delta E = Efinal - Einitial = q +w
what is 1 calorie in joules
4.184 J
define system and surroundings
- system: object(s) being studied
- surroundings: everything outside of the system
describe what it means when the change in energy is <0
- initial energy is higher than final energy
- energy of reactants higher than energy of products
- energy is lost to surroundings
describe what it means when the change in energy is >0
- final energy is higher than initial energy
- energy of products is higher than energy of reactants
- energy gained from surroundings
define state functions
- depend on initial and final states of a system
- uppercase variables: V, E, P, T
define path functions
- depend on the path a system takes to get from the initial to the final state
- lowercase variables: q, w, t, d
which type of function is easier to measure
state functions
define internal energy
energy contained within a system
define kinetic energy
energy associated with motion
what are the three types of kinetic energy
- translation
- vibration
- rotation
define potential energy
energy associated with position
what are the two types of potential energy
- chemical bonds
- intermolecular forces
what is the order of phase states from lowest energy to highest energy
- lowest energy: solid
- middle energy: liquid
- highest energy: gas
is energy absorbed or released when water evaporates
- liquid to gas: lower energy to higher energy
- energy of products (gas) greater than energy of reactants (liquid)
- energy is absorbed
define specific heat (C)
amount of heat required to raise the temperature of ONE GRAM of a substance one degree celsius
define molar heat capacity (Cm)
amount of heat required to raise the temperature of ONE MOLE of a substance one degree celsius
does water have an unusually high or low specific heat
high - 4.814 J/gC
define the variables in q=mCAT
- q = heat (J)
- m = mass (g)
- C = specific heat (J/gC)
- AT = temperature change, final-initial (C)
how is the heat of the system related to the heat of the surroundings
- q(system) = -q(surroundings)
- heat from system is equal and opposite of heat from surroundings
- system with either gain or lose heat while the surroundings do the other until the temperatures of both are equal
define calorimetry
the measurement of heat changes
describe the equation used in calorimetry
- q(reaction) = -q(calorimeter)
- q(reaction) = -q(water)
- q(reaction) = -mCAT(water)
what are two types of calorimeteres
- coffee cup calorimeter
- bomb calorimeter
why is the molar energy of a reaction typically reported rather than just the energy of the reaction
- molar energy takes into account the amount of substance that was reacted
- used to compare with other experiments
how do you find the molar energy of a reaction
q(reaction) of Xg / Xg = q(reaction) of 1 mole / molar mass
define enthalpy (H)
- heat flow under constant pressure (same as q but with constant pressure)
- state function
equation for change in H
delta H = Hproducts - Hreactants
define endothermic and it’s relation to H
- energy is absorbed by the reaction
- reaction gets warmer, surroundings get cooler
- change in H > 0
- products have higher energy than reactants
define exothermic and it’s relation to H
- energy is released by the reaction
- reaction gets cooler, surroundings get warmer
- change in H < 0
- products have lower energy than reactants
what is the simple equation for endothermic reactions
energy + reactants = products
what is the simple equation for exothermic reactions
reactants = products + energy
is enthalpy intensive or extensive
- extensive
- scalable: double products = double H
which takes more energy: phase changes or warming a substance in one phase
phase changes
which phase change rakes the most energy
- vaporization
- liquid to gas
how do you determine the energy used to heat a substance in a singular phase
q=mCAT
how do you determine the energy used during a phase change
delta H(fus/vap) + moles of substance
what are the three techniques for heats of reaction calculations
- bond enthalpy
- hess’s law
- heats of formation
which takes energy and which releases energy: making bonds and breaking bonds
- takes energy to break a bond
- energy released when bond is made
define bond enthalpy
how much energy is required to break a bond
are bonds being broken endothermic or exothermic
- endothermic
- energy taken in
are bonds being made endothermic or exothermic
- exothermic
- energy released
how can the heat of a reaction be approximated using bond enthalpies
heat of reaction = sum of reactant bond enthalpies - sum of product bond enthalpies
is the reaction endothermic or exothermic if the heat of reaction is negative when using bond enthalpies technique
- exothermic
- means more energy is on the products side
is the reaction endothermic or exothermic if the heat of reaction is positive when using bond enthalpies technique
- endothermic
- means more energy is on the reactants side
describe hess’s law
- if a reaction is carried out by a series of steps, enthalpy (H) for the overall reaction with be equal to the sum of the enthalpy changes for the individual steps
- taking parts of a chemical reaction, rearranging them to mirror the overall reaction, and adding the enthalpies
what happens to the enthalpy of a reaction when you switch the reactants and the products sides
flip the sign of the enthalpy (H)
define enthalpy of formation (Hf)
enthalpy change for the reaction in which one mole of a compound is made from its constituent elements in their elemental forms
define standard enthalpy of formation (Hf)
enthalpy of formation (Hf) measured under standard conditions (25 C, 1 atm, 1M)
what is the enthalpy of formation value (Hf) for elements in their elemental state
0
describe the equation used to determine the enthalpy of a reaction using enthalpy of formation
enthalpy of reaction (H) = sum of enthalpy of formations for products - sum of enthalpy of formations for reactants
