Lectures 21-26 Flashcards
energy
quantitative property that provides the ability to ‘do work’ or ‘supply heat’
potential energy (locked-in energy)
stored energy
kinetic energy (movement)
energy in motion
thermodynamics
science of the relationship between heat and other forms of energy
1 J =
1 Nm
1 J =
1 kg m^2s^-2
calorie (cal)
energy required to raise the temperature of 1g of water by 1 degree Celsius (1 cal = 4.184 J)
Calorie (C)
used to represent energy content of our food; 1 C = 1000 cal
The Law of Conservation of Energy
energy can be converted from one form into another, but it CANNOT be created or destroyed. The total energy of the universe is constant
universe
system + surrounding
open system
matter can be transferred through opening in the flask; heat can be conducted through flask walls
closed system
flask is stoppered so no matter can be transferred; heat can be conducted through the flask walls
isolated system
heat transfer is prevented by the vacuum flask; matter cannot be transferred
isothermal change
heat is exchanged between system and surroundings, so their temperatures are equal
adiabatic change
no heat exchange between system and surroundings, so their temperatures may not be equal
heat (q)
energy transferred from one system to another due to a temperature difference
temperature (T)
measure of ‘how hot’ something is (kinetic energy) and ability to transfer heat to other systems or surroundings
extensive property
a property of matter that changes as the amount of matter changes
is heat an extensive or intensive property?
extensive
intensive property
property of matter that does not change as the amount of matter changes
is temperature an extensive or intensive property?
intensive
work (w)
involves energy exchange as a result of motion against an opposing force (F)
work equation
w = Fd
when a system does work on the surroundings,
it loses energy
when the surroundings do work on the system,
it gains energy
volume work equation
w = -p(Ext)deltaV where p(ext) is external pressure
endothermic reaction
chemical process that absorbs heat
exothermic reaction
chemical process that releases heat
activation energy
energy required to break the bonds of the reactants
state functions
depend only on the amount of substance and its conditions; pathway does not matter
internal energy (U)
the sum of all its kinetic and potential energies of all the atoms, ions and molecules within the system
internal energy is an
extensive property and a state function
deltaU equation
deltaU = U(final) - U(initial)
change in internal energy in a closed system
deltaU = q + w
system gains internal energy
deltaU > 0
system loses internal energy
deltaU < 0
K and Celsius interchangeable for
SHC calculations
enthalpy change (deltaH)
equal to the heat transferred between the system and surroundings during a process that occurs at constant pressure
heat transferred under constant pressure
q(p) (deltaH = q(p))
enthalpy is a
state function
deltaH for exothermic
< 0
deltaH for endothermic
> 0
relationship between internal energy and heat gained/lost
deltaH = deltaU + pdeltaV
for systems containing no gases, deltaH
approximately equals deltaU
for systems containing gases, deltaH
= deltaU + deltan(gas)RT
molar enthalpy change of fusion/latent heat of fusion
energy required to melt 1 mol of a pure substance at its melting point
molar enthalpy change of vaporisation/latent heat of vaporisation
energy required to vaporise one mole of a pure substance at its boiling point
molar enthalpy change of fusion for water
6.00 kJ mol^-1
molar enthalpy change of vaporisation for water
40.65 kj mol^-1
heat capacity
tells us how much heat energy we need to put into a substance to raise its temperature
specific heat capacity (Cs)
heat needed to raise the temperature of 1g of substance by 1 K
specific heat capacity equation
Cs = q/mdeltaT
molar heat capacity (Cm)
heat required to raise the temperature of 1 mol of substance by 1 K
molar heat capacity equation
Cm = q/ndeltaT
for gases, heat capacity depends on whether measurement is carried out at
constant pressure (Cp) or constant volume (Cv)
molar heat capacity at constant pressure
Cp = q(p)/ndeltaT J K^-1 mol^-1
molar heat capacity at constant volume
Cv = q(v)/ndeltaT J K^-1 mol^-1
enthalpy change of a reaction (deltarH)
the difference between the sum of enthalpies of the products and the sum of enthalpies of the reactants
solution calorimetry
performed under constant pressure
bomb calorimetry
performed under constant volume
equation for heat transferred by a particular amount of substance
q = C x m x deltaT
as bomb calorimeters maintain constant volume, deltaU =
q(v) (no work term)
enthalpy change of formation
enthalpy change when 1 mol of a substance is formed under
standard conditions, from its constituent elements in their standard states
standard enthalpy change of a reaction (deltarH^o)
the enthalpy change when all the reactants and products are in their standard states
standard conditinos
standard conditions
1 M, 1 bar, 298 K
Hess’s Law
states if a reaction can be written as the sum of two or more steps, its enthalpy change of reaction is the sum of the enthalpy changes of reaction of each of the steps
standard enthalpy change of formation for a substance in its standard state
zero
standard enthalpy change of formation equation
v(i)standard enthalpy change of formation of products - v(i)standard enthalpy change of reactants
homolytic dissociation
cutting through a bond right down the middle
bond-breaking costs energy and is therefore
endothermic
bond energy (D)
the energy required to break 1 mol of bonds
calculating enthalpy change using bond energies
deltarH = D(bonds broken) - D(bonds formed)
summary of ways to calculate enthalpy changes
using steps (Hess’s Law), using enthalpy changes of formation equation, using bond energies
when temperature increases, the enthalpy of a substance
increases
calculating deltaH using molar heat capacity at constant pressure
deltaH = C(p) x deltaT
calculating deltaC(p)
v(i)C(p)products - v(i)C(p)reactants
Kirchoff Equation
standard enthalpy change at T2 = standard enthalpy change at T1 + deltaC(p)(T2 - T1)
system absorbs heat
q > 0
system gives off heat
q < 0
surroundings does work on system
w > 0
system does work on surroundings/work done by system
w < 0
spontaneous process
process which, once started, occurs without any external intervention or action
some spontaneous processes involve no
exchange with heat with surroundings at all
some spontaneous reactions occur
very slowly
entropy
degree of disorder in a system, its surroundings and the universe
higher the degree of disorder,
higher the entropy
probability of a disordered state is
higher than the probability of an ordered state
Second Law of Thermodynamics
spontaneous processes are those that increase the total entropy of the universe
to calculate SPONTANEOUS change of the universe
deltaS(universe) = deltaS(system) + deltaS(surrounding) > 0
microstate
each possible arrangement of gas molecules
macrostate
equivalent microstates
probability of macrostate
add up probabilities of its microstates
when the number of atoms is large, we reach the
thermodynamic limit
as number of atoms increases,
distribution of matter with a high degree of order is very unlikely
driving force in dispersal of matter
maximising system’s entropy
maximally dispersed
more disordered
exception to mixtures always leading to greater disorder (example)
LiOH crystal lattice
Boltzmann Equation
S = kBln(W)
Boltzmann constant, kB
R/N(A)
W
number of microstates in a given macrostate
more microstates in a given macrostate, the higher the
entropy
how to calculate W
number of possible arrangements for one particle (w) to the power of the number of particles
entropy sign of UNIVERSE governs whether reaction is
spontaneous or not (even if reaction entropy change is negative)
reversible process
system changes in such a way that the system and surroundings can be put back in their original states by exactly reversing the process
in a reversible system, the entropy of the
universe does not change
phase changes are
reversible
Third Law of Thermodynamics
there is no disorder (S = 0) in a perfect crystal at 0 K
entropy increases with
temperature
entropy is a
state function
how to calculate deltaS(sys)
q(rev)/T J K^-1
standard molar entropy (S^o)
defining entropy for 1 mol of substance under standard conditions
entropies of substances increase as they change
from solid through to liquid and then gas (more freedom of motion)
entropies of substances composed of larger molecules have
higher S^o than those made up of smaller molecules
substances whose molecules have a more complex structure have a
higher S^o than substances with simpler molecules
within the same phase, the entropy of 1 mol of substance for a temperature T2
S^o(T2) = S^o(T1) + C(p)ln(T2/T1)
deltarS < 0 from
gas to liquid
how to calculate deltarS^o
v(i)S^o(products) - v(i)S^o(reactants)
if deltaS(univ) < 0, the reaction is
spontaneous in the opposite direction
if deltaS(univ) = 0, the system is at
equilibrium
deltaS(sys) is the same as
deltarS
how to calculate deltaS(surr)
-deltarH(sys)/T
how to calculate delta(vap)S^o
delta(vap)H^o/T(boiling point)
how to calculate delta(fus)S^o
delta(fus)H^o/T(melting point)
above freezing point
spontaneous melting
below freezing point
melting not spontaneous
Gibb’s free energy change (deltaG) equation
deltaH(sys) - TdeltaS(sys)
Gibb’s free energy change
thermodynamic function that brings together the properties of enthalpy and entropy for a system
units for Gibb’s free energy
J mol^-1 or kJ mol^-1
Gibb’s free energy is a
state function
Gibb’s free energy change for a reaction (deltarG)
G(products) - G(products)
if deltarG < 0, reaction is
spontaneous
if deltarG is > 0, reaction is
non-spontaneous
if deltarG = 0, reaction is
at equilibrium
if reaction is exothermic and has negative deltaS, it will be spontaneous
at low temperatures
if reaction is exothermic and has positive deltaS, it will be spontaneous
at all temperatures
if a reaction is endothermic and has positive deltaS, it will be spontaneous
at high temperatures
if a reaction is endothermic and has negative deltaS, it will be non-spontaneous
at all temperatures
product is favoured,
reaction is entropy driven
standard Gibbs free energy of formation (deltafG^o)
change in free energy when 1 mol of pure substance is formed (in its standard state) from its component elements (in their standard states) at 1 bar pressure
standard Gibbs free energy of formation (deltafG^o) equation
v(i)deltaG(products) - v(i)deltaG(reactants)
elements in their standard states have deltafG^o = 0 and deltafH^o = 0, but their So at 298 K is
nonzero
Hess’s Law also applies to
Gibb’s free energy change
free energy work is equivalent to
maximum non-expansion work that can be obtained from a change; maximum work after change in volume has been accounted for
maximum amount of energy that can be used to do work
deltarG^o = w(max)
product favoured
G < 0
reactant favoured
G > 0
