Midterm 2 Flashcards
Energy
capacity to do work or transfer heat
work
energy that causes a mass to be moved by applying a force
work=-pΔV
Heat
q.. energy that causes the temperature to go up
Matter possesses energy in the form of
1) kinetic energy
2) thermal energy (heat)
3) potential energy
kinetic energy
energy of motion of an object
thermal energy
kinetic energy of molecules .. more motion, hotter the object
potential energy
position of object.. stored energy due to position or composition
difference between thermal and kinetic energy
kinetic moves in one direction.. thermal is individual molecules (which are all moving in different directions)
calories
defined such that 1cal=4.184J
1 calorie = 1000 cal = 1 kcal
open system
can exchange matter and energy with the surroundings
closed system
can only exchange energy with the surroundings (ex heat)
isolated system
neither matter nor energy are exchanged with the surroundings
First law of thermodynamics
the total energy of the universe is conserved.. any energy lost by the system is gained by the surroundings and vice versa.
ΔU=q+w
q+=heat added, q-=heat released
w+=work done on, w-=work done by system
U
internal energy - total energy possessed by a system
ΔU=q+w
ΔU=Uf-Ui or ΔU=Uproducts-Ureactants
state function
A property of a system that is determined by the state or condition of the system and not by how it got to that state. It’s value is fixed when P,V,T,G&S are specified
(non-state = g, w)
enthalpy
ΔH
-the heat change at constant pressure is the change of enthalpy of the system
thermochemical equation
1) depends on amounts of reactants and products
2) ΔH forward = -ΔH reverse
3) ΔH depends on the physical state of reactants and products
balanced chemical reaction and the value of ΔH.. **Provides relationship between amounts of chemicals and the enthalpy change
heat capacity
amount of heat required to raise the temperature of an object by 1K (or 1°C)
-depends on size of object bonding, complexity, physical state
specific heat capacity
It is the amount of heat required to change the temp of a gm of a substance by 1°C.
Cs… can be determined experimentally
q=CsXmass(g)XΔT
calorimetry
measure of heat flow
constant pressure calorimetry
measures ΔHrxn
heat flows between reaction and solution at constant pressure
at constant P, qrxn =ΔHrxn
constant volume calorimetry
measures ΔUrxn
heat from chemical reaction absorbed by water and all components of calorimeter
at constant V, qrxn = ΔUrxn
Hess’s Law
The enthalpy change of an overall process is the sum of the enthalpy changes of its individual steps - combines the individual reactions so their sum gives the desired reaction
-can be used for both chemical and physical changes
-magnitude of ΔH depends on T, P and physical state of prod and reacts.
standard enthalpy change
ΔH°rxn. (°means under standard conditions)..
the enthalpy change for a reaction with reactants and products in their stable forms at 1atm and 298K.
standard enthalpy of formation ≈
ΔH°f
the enthalpy change that accompanies the formation of one mole of a substance from its elements, with all substances in their standard states
for pure elements in their most stable form at standard conditions.. ΔH°f=0
Steps to take using ΔH°f to calc ΔH°rxn
- decompose reactants into elements
- recombine elements into products
ΔH°rxn =ΣnΔH°f (products) - ΣnΔH°f (reactants)
fuel value
energy released when one gram of substance is combusted
foods and fuel value
carbs ≈ proteins < fats
fats easily stored. excess carbs not readily stored, and protein used to build
Fuels and fuel value
natural gas>gasoline>coal (E/mass)
Spontaneity
physical and chemical change tends to favour one direction over the other..
spontaneous process occurs without ongoing outside intervention
nonspontaneous process
can occur only when the surrounding does work on the system or transfers energy to or from it
Reversible process
the change occurs in such a way that both the system and the surroundings can be restored to their original states by exactly reversing the change
in a reversible change, the max possible work is done
Irreversible process
the system and the surroundings cannot both be returned to original state
Spontaneous reactions are always irreversible
Entropy
(S) a function that measures randomness in a system
entropy changes
ΔS
JK^-1
depends on heat and temperature
phase changes are examples of isothermal processes
ΔS = qrev / T = ΔH(phase change)/T
second law of thermodynamics
In any spontaneous process, the total entropy of the universe increases
ΔSuniv>0 (for a spontaneous process)
Entropy increases for these processes
solid->liquid->gas #moles or elements of products higher than number of reactants
Third law of thermodynamics
The entropy of a pure crystaline substance at 0K is zero
- sets a baseline for further experiments because we know zero at 0K.
- Absolute value of entropy, S, can be determined by measuring how much entropy rises as a crystal is warmed from 0K.
Entropy changes in chemical reactions
1) heating increases molecular motion and entropy
2) phase changes are isothermal, entropy changes sharply
Standard molar entropy
the entropy value for a mole of a substance in its standard state
- gases>liquid>solid
- increases with molar mass
- increases with number of atoms
Change in entropy
ΔS°sys= ΣnS°(products)- ΣnS° (reactants)
ΔSsurroundings= -qsys/T at constant P for isothermal process.
Gibbs Free Energy
Thermodynamic state function that combines enthalpy and entropy, G=H-TS
the process is spontaneous and ΔG must be negative.. ΔG determined the direction of spontaneous change
In any spontaneous process at constant T and P, free energy decreases, ie ΔG is negative
ΔGsys=ΔHsys-TΔSsys or
Gibbs determining positive or negative
ΔG>0 the reaction is not spontaneous in the forward direction. work must be done to make it occur
ΔG=0 the reaction mixture is at equilibrium
Standard free energy of formation
Defined for formation of 1 mole of substance from its composed elements in their standard states.
temperature = 298 K pressure = 1 atm (gases) pure liquid or solid unit molarity (solutions)
elements in most stable form
ΔG°f=0
When ΔH & ΔS have same sign..
the sign of ΔG depends on how big they are
Free energy and the equilibrium constant
Most chem reactions do not occur under standard conditions, and need to obtain ΔG from ΔG°
ΔG=ΔG°+RTlnQ
R=gas constant
Q=reaction quotient
Q depends on actual, nonstandard conditions and is useful in predicting the direction of a reaction
Q predicting direction of a reaction
QK reaction is spontaneous from right to left
Q=K equilibrium
The direction of sponteneity can be changed by altering the concentrations of reactants and prods, which changes Q, and the value of ΔG
Relationship between ΔG and K
At equilibrium ΔG=0 and Q=K
Equilibrium:concentrations
concentrations of reactants and products become constant over time at equilibrium, concentration of reactants and products are not equal but they are unchanging.. reaction continues even @ equilibrium
Equilibrium constant
Kc = equilibrium constant when concentrations are in M
Evaluating Kc
insert equilibrium concentrations into kc expression
equilibrium constants and pressure
Kp - same as kc but with P instead of concentrations
Magnitude of equilibrium constants
the value of Kc (or Kp) indicates how far the reaction will proceed towards products
10^-3 - 10^3 = significant concentration of both products and reactants
direction of net reaction
depends on the relative values of Qc & Kc
When Q>K, there are more products than reactants. To decrease the amount of products, the reaction will shift to the left and produce more reactants.
When Q<K there are more reactants than products. To decrease the amount of reactants, shift to the right to produce more products.
Q=K reactants and products are at equilibrium
Le Chateliers principle
If an equilibrium is distrubed by changing concentrations, T, or P, the equilibrium will shift to try to counteract the change
[A] = will shift to use up or replace added/taken away concentrations
V & total pressure = partial pressures change and Qp & Kp are not at equilibrium.
^ in total pressure = reaction shifts toward fewer moles of particles
Free energy and equilibrium constant
direction of spontaneity can be changed by altering the concentrations of reactants & products, which changes Q and therefore changes the value of ΔG
Temp effect of equilibrium
values changes with temp..
endothermic = heat is a reactant exothermic = heat is a product
System adjusts to counteract this heat change
Q
Reaction quotient:
the value that is obtained when concentrations of reactants and products are inserted into the equilibrium expression. If the concentrations are at equilibrium: Q=K, if not Q does not = K
