UNIT 3 Flashcards
Define thermodynamics
the study of energy and its conversion from one form to another
Define thermochemistry
a branch of thermodynamics, which highlights how heat is involved in chemical and physical transformations
there are two basic kinds of energy:
potential energy (Ep): stored energy
kinetic energy (Ek): the energy of motion
Ek =
1/2mv^2
internal energy, E, of a system
the sum of all energy (potential and kinetic) of everything in that system
chemical bond energy
potential energy stored in molecular bonds
energy of molecular motion
- translational, rotational, and vibrational
- sometimes also called thermal energy
- temperature is a common measure of the energy of molecular motion
define energy
the capacity to do work or supply heat
change in energy of a system =
work done on the system + heat flow in to the system
ΔE = w + q
energy is measured in
Joules or kiloJoules
1st law of thermodynamics
energy cannot be created or destroyed
ΔE(universe) = 0
the total energy of the universe is conserved
system
chosen to include whatever you are focusing on (and things that cannot usefully be separated from that)
surroundings
everything not in the system
universe =
system + surroundings
how can we determine the energy change of a system of interest and why?
by measuring the energy change in the surroundings; if energy leaves the system, it must enter the surroundings (and vice versa)
any change in the energy of a system must be accompanies by
an equal magnitude change in the energy of the surroundings, but the signs of these changes must be opposite
ΔE(universe) =
ΔE(system) + ΔE(surroundings) = 0
in thermodynamics, signs are defined from
the system’s point of view/perspective
ΔE of a reaction
= E(products) - E(reactants)
if a reaction releases energy
- system loses energy
- reaction produces heat or does work
- reactants higher in energy than the products
ΔE < 0
state function
values that depend on the state of the substance, and not on how that state was reached. For example, density is a state function, because a substance’s density is not affected by how the substance is obtained.
examples of state functions
E, V, P
examples of non-state functions
q and w
heat (q)
heat (q), also called thermal energy, will flow from higher-temperature objects to lower-temperature objects. eventually, the objects will reach the same temperature
work (w)
mechanical work is the product of force (F) and distance (d)
w = F x d
the larger the required force, or the longer the distance on an object is moved, the more work is done on it
example of energy transfer as work (w) only (q=0)
gas forming reaction in an insulated container that is attached to a piston-cylinder assembly that pushes against something
- the system pushes the piston out, doing work on the surroundings
- system releases (loses) energy
- ΔE = w < 0
expansion work
the work done when the volume of a system changes in the presence of an external pressure
expansion work is also often known as
pressure volume (PV) work because the amount of work done depends on both P and V
at a constant pressure, P=P(surr), w=
-PΔV = -P(V(final) - V(initial))
more work is done when the volume change (ΔV) is larger and/or when pushing against a higher external pressure
why is there a negative sign for expansion work?
- required to fit with the convention that signs reflect the system’s perspective
- when a system expands against a pressure:
-> P, which is the constant pressure of the surrounding, is always >0
-> the system does work. it loses internal energy by doing this work
-> it must be that w<0, ΔE<0
we calculate: w = -PΔV <0
enthalpy (ΔH)
defined as
H = E + PV
another accounting system for keeping track of energy, similar to internal energy, but does not include expansion work
change in enthalpy (ΔH)
ΔH = ΔE + PΔV
- amount of heat absorbed or released in a transformation (real or imagined reaction) at a constant pressure.
- this means enthalpy change (ΔH) is often straightforward to determine via experiment
the sign of ΔH shows
whether heat is produced or consumed in a reaction.
ΔH<0
exothermic
- evolved heat flows out of the system into the surroundings
- heat is a product of the reaction
ΔH>0
endothermic
- heat flows into the system from the surroundings
- heat is a reactant
is enthalpy a state function?
yes.
ΔH = H(products) - H(reactants) = H(final) - H(initial)
differences and similarities between enthalpy (ΔH) and internal energy change (ΔE)
- differ in whether they account for the expansion (PV) work done on the surroundings or done by the surroundings on the system
- most reactions involve little (if any) PV work. this is especially true for reactions involving only solids and liquids, or reactions in which the amount of gas does not change
thus, often (but not always) ΔH = ΔE
effect of multiplying a reaction on ΔH
increases ΔH by the same factor
effect of reversing a reaction on ΔH
changes the sign of ΔH for a reaction
heat capacity
amount of heat (q, or ΔH) required to raise the temperature of an object or substance by 1’C (or 1K)
heat capacity = q/ΔT
things with high heat capacity
- require a lot of heat to increase in temperature and give off a lot of heat when they cool down.
- act as a good sponge for (or, store of) thermal energy
types of heat capacity
specific heat capacity (c): heat capacity per mass (1.00g)
c = q/(mass x ΔT)
units of J/gK
molar heat capacity (C or Cm): heat capacity per mole
C = q/(amount (mol) x ΔT)
units of J/molK
effect of energy from heat, q
excites both translational motion of molecules and vibrations and rotations within and between molecules
calorimetry
the science of measuring the heat exchanged in chemical reactions
in calorimetry, the heat of reaction (qrxn) is measured
indirectly, by means of a calorimeter.
- if the reaction produces heat, the temperature of the surroundings increases and vice versa
q(rxn)
= - q(calorimeter)
= -(q(vessel) + q(solution) + q(other))
constant-pressure, or ‘coffee-cup’ calorimetry
q(system) = - q(calorimeter)
- reaction or physical transformation is done in an insulated container at a constant pressure in a bath (usually water) of known heat capacity
- heat generated thus tells us about the enthalpy of reaction
q = qp = ΔE + PΔV = ΔH
- heat capacity of the calorimeter (often ~equal to that of the surrounding water because absorption of heat by the vessel is minimal) provides the link between change in temperature and heat gained or lost, from which you can determine the enthalpy change of the system
Heat capacity (calorimeter) = q(cal)/ΔT(cal)
second type of calorimetry
constant volume or ‘bomb’ calorimetry:
- measures the heat change at constant volume such that q = qv = ΔE + PΔV = ΔE
- commonly used to measure heat of combustion reactions
Hess’s Law
the enthalpy change for a process is equal to the sum of the changes for individual steps (1, 2, …n) of the process
ΔH(overall) = ΔH1 + ΔH2 + …. + ΔHn
standard states
- pure substance: most stable form at 1atm
- gas: 1atm and ideal behaviour
- substance in aqueous solution: 1M concentration all at a specified temperature, which is usually 25’C
standard enthalpy of reaction
ΔH’(rxn)
- all reactants and products are in standard states
ΔH’f
standard heat of formation
- enthalpy change for the formation of 1 mole of substance in its standard state from its constituent elements in their standard states
we can use our database of ΔH’f’s to calculate ΔH (rxn)
ΔH’(rxn) = ΔH’f(products) - ΔH’f(reactants)
- ΔH of each reactant or product must be multiplied by its stoichiometric coefficient in the balanced rxn equation
ΔH’(rxn) =
bonds broken - bonds made
Spontaneous change
one that occurs without a continuous input of energy from outside the system (though activation energy may be required to initiate it)
If a change is spontaneous in one direction
it will be non-spontaneous in the reverse direction
how may non-spontaneous reactions and processes be driven?
with continual input of energy
is enthalpy change a predictor of spontaneity?
many exothermic processes are spontaneous (eg combustion reactions)
examples of spontaneous endothermic transformations
- melting of ice (at higher temps)
- dissolution (of some solids at some concentrations and temps)
all show an increase in the freedom of motion of particles in the system
relationship between freedom of motion and spontaneity
an increase in freedom of motion (dispersal of energy) favours spontaneity
entropy, S
measure of energy dispersal, or freedom of motion, in a system
positive value of ΔS indicates
negative value of ΔS indicates
increased dispersal of energy
decreased dispersal of energy
is entropy a state function?
yes
ΔS = S(final) - S(initial)
units of entropy
J/K
S’
standard molar entropy: entropy of 1 mole of the pure substance in its standard state
entropy trends
- effect of physical state
- effect of particle numbers
- effect of molecular complexity
- effect of temperature
effect of physical state
S of solids < S of liquids «_space;S of gas
solids - less energy dispersed, lower entropy
effect of particle numbers
more molecules have higher entropy than fewer molecules
effect of molecular complexity
entropy increases with chemical complexity and flexibility
- this only holds for substances in the same physical state
- the effect of physical state dominates the effect of molecular complexity
effect of temperature
as temperature increases, entropy increases
- higher temperature means more freedom of molecular motion
draw graph for temperature vs entropy
discontinuous jumps at phase changes
entropy change upon dissolution
- salt gains entropy as it is dispersed
- water loses entropy as it is ordered around the ions
- net entropy change depends on the relative magnitudes of entropy changes in both solute and solvent entropy
why does increased entropy favour spontaneous change?
high-entropy configurations can be achieved in more ways than low-entropy configurations. they are therefore more likely to occur
draw a diagram for spontaneous expansion of a gas and explain it
- arrangement B has higher entropy (is entropically favoured)
- gas expands when stopcock is opened because there are more ways to achieve the configuration on the right (B) than on the left (A)
absolute entropies can be calculated from
the number of micro states (W) a system may occupy
2nd law of thermodynamics
spontaneous reactions proceed in the direction that increases the entropy of the universe (system + surroundings)
ΔS(Universe) = ΔS(sys) + ΔS(surr) > 0
thus, any decrease int he entropy of the system must be offset by a larger increase in the entropy of the surroundings for that process to be spontaneous
2nd law has profound implications:
- isolated systems always evolve toward higher entropy states
- entropy of the universe is always increasing
how are temperature, heat flow, and entropy linked?
- if heat flows into a system from the surroundings, the entropy of the system increases. the surroundings lose entropy
- the amount by which a given amount of heat flow changes entropy depends on temperature. if it is low, the effect on entropy can be enormous
3rd law of thermodynamics
a perfect crystal has zero entropy at absolute zero
S(sys) = 0 at 0K
has flawless alignment of all its particles. at absolute zero, particles have minimum energy so there is only one micro state
contrast between S and H
- entropy scale is anchored to an absolute value
- enthalpy does not have an absolute 0
standard entropy of a reaction ΔS’(rxn)
entropy change that occurs when all reactants and products are in their standard states
ΔS’(rxn) = S’(products) - S’(reactants)
remember to include coefficients
Gibbs free energy change ΔG
evaluates spontaneity as a function of enthalpy and entropy of the system alone
ΔG(sys) = ΔH(sys) - TΔS(sys)
ΔG < 0 process is spontaneous
ΔG = 0 process is at equilibrium
ΔG > 0 process is non-spontaneous
- lowering free energy is the driving force of chemical reactions
- negative ΔH(sys) and positive ΔS(sys) favour spontaneity
- entropic contribution to free energy change (-TΔS) is increasingly important at higher temperatures
give another way to define ΔG
the maximum useful work that can be done by a system as it undergoes a spontaneous process at constant temperature and pressure
ΔG = w(max)
also the minimum work that must be done on a system to drive the occurrence of a non-spontaneous process
is ΔG a state function?
yes
describe the extensive property of ΔG
scales linearly with amount
ΔG’f
standard free energy of formation of a compound from its constituent elements in their standard states
ΔG’f of an element in its standard state is
0
thermodynamics vs kinetics
ΔG tells us whether a reaction will/won’t proceed
the free energy of activation (including Ea) tells us how fast a reaction proceeds
ΔG and spontaneity
reaction is spontaneous when ΔG(rxn)<0
how do we make a non-spontaneous reaction happen?
must be driven by coupling the non-spontaneous reaction with a spontaneous reaction of sufficiently favourable ΔG
how are free energies and equilibrium position linked?
ΔG depends on how much product and reactant is present at that instant (Q0 compared to their equilibrium values (K), the temperature and the gas constant:
ΔG = RTln(Q/K)
how does the magnitude of ΔG tell us how far out of equilibrium the mixture is?
ΔG < 0 ; Q < K ; ln(Q/K) < 0 - process proceeds (forward spontaneously)
ΔG = 0 ; Q = K ; ln(Q/K) = 0 - process is at equilibrium
ΔG > 0 ; Q > K ; ln(Q/K) > 0 - reverse process proceeds spontaneously
if Q and K are very different
ΔG has a very large value (negative or positive). The reaction releases or absorbs a large amount of free energy as it proceeds to equilibrium
if Q and K are nearly the same,
ΔG has a very small value (negative or positive). The reaction releases or absorbs very little free energy as it proceeds to equilibrium.
how are thermodynamic Q and K different to Qc and Kc
in the thermodynamic Q and K, all substances are referenced to their own standard states, So, Q and K may include mixed states
equation linking ΔG, gas constant, temp, and K
-RTln(K)
another formula for ΔG
= ΔG’ + RTln(Q)
draw the two free energy hills with ΔG
features of free energy hills
- slope at any given point tells us the value of ΔG for that mixture
- as the system approaches equilibrium, ΔG approach 0
- at equilibrium, the free energy is at a minimum.
Define electrochemistry
the interconversion of chemical and electrical energy
- spontaneous reactions can produce electricity and electricity can cause non-spontaneous reactions to occur
redox reactions involve
the movement of electrons from one reagent to another
oxidation
loss of electrons
reduction
gain of electrons
oxidising agent
- reduced
- takes electrons from the substance being oxidised
reducing agent
- oxidised
- gives electrons away
how to balance redox reactions
- write down the two half-reactions
- balance the atoms and charges in each half-reaction
- first balance atoms other than O and H, then O, then H
- charge is b balanced by adding electrons (e-) to the reactant side of the reduction half-reaction and to the product side in the oxidation half-reaction
- electrons should cancel out in net reaction. if necessary, multiply one or both half-reactions by an integer so that number of e- gained in reduction = number of electrons lost in oxidation
- add the balanced half-reactions and include states of matter
two types of electrochemical cells
- voltaic cells
- electrolytic cells
voltaic/galvanic cells
spontaneous chemical reaction (ΔG<0) generates an electric current
- batteries contain one or more voltaic cells
- voltaic cell does work on the surroundings, converting higher energy reactants in the cell into lower energy products
electrolytic cells
non-spontaneous reactions (ΔG>0) are driven by electric current
- external power source supplies free energy to run electrolytic cells. the surroundings thus do work on the cell. lower energy reactants are converted to higher energy products in the cell
- used for electroplating, purification of metals, and more
electric current is
flow of electrons
electrons only flow if
the driving force (free energy change) is large enough
electrodes
usually metal strips/wires connected by an electrically conducting wire
anode
electrode where oxidation takes place
cathode
electrode where reduction takes place
describe salt bridge
U-shaped tube that contains a gel permeated with a solution of inert electrolyte (contains positive and negative spectator ions).
function of salt bridge
The salt bridge keeps half cells electrically neutral because ions flow in and out of the salt bridge, counteracting charge build-up due to electron flow.
conventions for notation for a voltaic cells
- anode components are written on the left
- cathode components are written on the right
- components of each half cell are written in the same order as in their half reactions
- single line shows a phase boundary between the components of a half cell
- double line shows that the half cells are physically separated
addition notation for voltaic cell that is more complex
any inactive (inert) electrode is specified
a comma is used to show components that are in the same phase
why does a voltaic cell work?
- differing abilities of metals to gain electrons gives rise to a voltage drop
- this is also known as electromotive force (EMF) or cell potential
multimeter
can measure voltage
cell potential (Ecell)
the difference in electrical potential between two electrodes
units for Ecell
measured in Volts (V), where
V = J/C (Colomb is SI unit of charge)
difference in Ecell for voltaic and electrolytic cells
Voltaic: Ecell>0 (spontaneous e- flow)
electrolytic: Ecell<0 (non-spontaneous e-flow)
when Ecell = 0,
the redox reaction has reached equilibrium so the cell can do no more work
define standard potential (E’cell)
cell potential under standard state conditions
how to calculate Ecell from Ehalf-cell
the cell potential of any electrochemical cell is the sum of the half cell potentials for the oxidation and reduction half cells
Ecell = E ox + E red
what does it mean to say that E is an intensive property?
Bigger cells with more moles of redox components will last longer, but will have the same standard output voltage
which half cell will form the oxidation half cell?
the lower Ered (ie more negative/less positive = worse at being reduced)
use of standard hydrogen electrode
- potentials are determined experimentally from the difference in potential between two electrodes
- the reference point is called the standard hydrogen electrode
Standard Hydrogen Electrode
consists of a platinum electrode in contact with H2 gas (1atm) and aqueous H+ ions (1M)
standard hydrogen electrode is assigned (arbitrarily) a value of
exactly 0.00 V
how is Ecelll related to ΔG?
ΔG = -nFE(cell) = -RTln(Keq)
- ΔG is in J/mol
- n is in mol and is the no of moles of electrons transferred per mole of the reaction
- F is faraday’s constant
at equilibrium, ΔG = and Ecell =
0
Nernst Equation
E = E’ - RT/nF (lnQ)
E = E’ - 0.0592V/n logQ
E - The cell potential (electromotive force, EMF) under non-standard conditions, measured in volts (V).
The standard cell potential, which is the voltage of the electrochemical cell under standard conditions
how does cell potential depend on the relative concentration of reactants and products?
when Q<1, lnQ<0 so Ecell>E’cell
when Q=1, lnQ=0 so Ecell = E’ cell
when Q>1, lnQ>1 so Ecell<E’cell
two consequences of cell potential depending on the relative concentration of reactants and products:
- as a cell is operated, concentration of reactants will decrease and products will increase. thus, cell potential will decrease over time
- concentration cells - capture the electrical energy from a concentration difference
how is a concentration cell created
- has the same half-reaction in both cell compartments, but with different concentrations of electrolyte
- there is a potential difference between cells, which drives current flow until both compartments have an equal concentration of ions
- Ecell>0 as long as the half-cell concentrations are different
- once the concentrations equalise, Ecell = 0 and current stops flowing
write equations for corrosion
how to protect against corrosion
Galvanisation: coating of iron with zinc.
Zinc is more easily oxidised than iron, giving up electrons to it. Iron is more easily reduced than zinc.
