Physical chemistry (Horrocks) Flashcards
Define isolated system
exchange neither energy nor matter
Define closed system
exchange energy, but not matter
Define open system
exchange energy and matter
What is the ideal gas equation and there units
pV = nRT
pressure / N m−2 or Pa
volume / m3
R = 8.314 J K−1 mol−1
n = mol
What is van der Waals gas law and what do the letters stand for.
p = (nRT/(V-nb))-a*(n/V)^2
a accounts for attractive interactions between the molecules
b accounts for repulsive interactions (finite size effect)
Advantages and disadvantages of ideal gas equation
Applies approximately to all gases at near-ambient temperatures and
pressures
No molecule-specific parameters; n is the total number of moles of
gas even if the gas is a mixture.
Fails to predict that gases liquefy at low temperature and high
pressure
Advantages and disadvantages of van der Waals gas law
Accounts for liquefaction (partially)
More accurate than the ideal gas law
Requires a knowledge of the a and b parameters for all the gas
species involved – the main reason we do not use it as much as the
ideal gas law.
What is the equation for compression or expansion of a gas
w = −p (Vf − Vi) = −p∆V
What are some key points about the equation for compression or expansion of a gas
work done on the gas is positive when the gas is compressed.
work done by the gas is negative when the gas is compressed.
work done on the gas = - work done by the gas
w / Joules (J)
∆V < 0 for compression; ∆V > 0 for expansion.
define heat
Heat is the transfer of energy in a disorganised form, driven by a
difference of temperature.
define work
Work is the transfer of energy in an organised form.
what do the following letters mean U, H, CV ,m, Cp,m
u = total energy
H = enthalpy
CV,m = heat capacity at constant volume
CP,m = heat capacity at constant pressure
what is the equation for enthalpy
total energy + (pressure*volume)
define molar heat capacity
The heat required to raise the temperature of one mole of substance by
one Kelvin.
what is the total energy change for an isolated system
0
what is the total energy change for a closed system
q + w
heat + work
define system
the portion of the universe under study
define surroundings
everything not part of the system
define thermodynamic variable
Quantities that pertain to the bulk properties of matter and which can, in principle, be
controlled in an experiment
define thermodynamic state
A specification of the values of the thermodynamic variables that uniquely define the condition and thermodynamic properties of a system
define equation of state
equation relating the values of the thermodynamic variables
define state function
A quantity that depends only on the thermodynamic state of a system and not on how it was prepared.
define entropy
a measure of disorder, both of energy and matter
give the 4 definitions of reversible process
- A process for which the system and surroundings can both be returned to their
initial conditions. - A process for which the total change in entropy of system and surroundings together
is zero. - A process which proceeds through a succession of equilibrium states.
- A process for which (∆G)p,T = 0.
define standard state/ conditions
- the most stable form of that substrate at a pressure of one bar
- Temperature is not part of the of the definition, standard state may be any temperature
define reference state
the stable form of the element under standard condition
define standard formation reaction
A reaction in which one mole of the substance is prepared under standard conditions from
the elements in their reference states. The changes in two important state functions H, G
for these reactions are tabulated as ∆f H⊖ and ∆f G⊖.
define mean bond enthalpy
The ∆H for a reaction X − Y → X + Y averaged over a large selection of compounds
with X − Y bonds is the mean bond enthalpy for X − Y bonds
define enthalpy of combustion
Enthalpy change when 1 mole of substance is combusted to liquid water and gaseous CO2
and N2 (if any N-atoms) under standard conditions. Symbol: ∆comH⊖
define activity
A dimensionless (no units) quantity related to concentration or pressure that is used to
express the equilibrium constant. The activity of species A may be denoted {A} or aA
for pure solids or liquids the activity is 1
define gibbs energy
Gibbs energy, also called free energy, G, is defined as:
G = H − T S
or equivalently
G = U + pV − T S
define useful work
Work done that corresponds to expansion of the system against an external pressure is
called pV -work or expansion work. All other types of work are called ’useful’ work.
define equilibrium constant
A ratio of activities of products and reactants, raised to the powers of the stoichiometric
coefficients in the balanced chemical equation. The equilibrium constant for a particular
choice of reaction depends only on the temperature
define intensive variable
A thermodynamic variable that does not scale with the size of the system
define extensive variable
A thermodynamic variable that does scale with the size of the system
define degrees of freedom
A degree of freedom is an intensive variable whose value can be chosen by the experimenter. p, T , mole fractions xi.
define phases
a phase is a uniform state of matter
define components
A component is an independent chemical species in the system, whose quantity can be
chosen by the experimenter. The number of components is the number of species minus the number of equilibria amongst them and minus the number of electroneutrality constraints.
define mole fraction
The fraction of a species in a mixture calculated on the basis of numbers of moles of each
species.
define phase rule
F = C - P+2
define normal melting point
The temperature at which the solid and liquid phases are in equilibrium at a pressure of
1 atm.
define normal boiling point
The temperature at which the liquid and gaseous phases are in equilibrium at a pressure of 1 atm
define critical point
The highest temperature at which the liquid phase exists. The ’end’ of the liquid-gas
boundary line on the phase diagram
define triple point
The unique temperature and pressure at which the solid, liquid and gaseous phases of a
pure substance are in equilibrium
define half cell reaction
The reaction occurring at a single electrode in which electrons are transferred to/from
solution
define cell rection
The overall reaction taking place in the cell. Conventionally the difference between the
right and left half cell reactions
define faraday’s law
Relationship between number of moles of an electrode reaction and the electrical charge
passed.
* Q = nF
* Charge, Q in Coulombs, C. Number of moles, n and the Faraday constant F = 96485
C mol−1
define the standard hydrogen electrode
The conventional zero for the electrode potential scale.
* The electrode reaction for the SHE is:
H+ + e− = 1/2H2
* When all species are in their standard states E (SHE) = 0 V.
* It is implemented by dipping an inert Pt electrode in 1M HCl and bubbling H2
through the solution at 1 bar
define cell potential
The potential difference between the two electrodes (half cells) of a cell.
E = ER − EL
define electrode potential
the cell potential for a cell in which the left hand electrode is the she
what is the acid dissociation constant
The equilibrium constant for the dissociation equilibrium of a weak acid.
HA = H+ + A−
* Usual symbol Ka
* Usually reported as a p-value: pKa = − log10 Ka
what is the base dissociation constant
The equilibrium constant for the dissociation equilibrium of a weak base.
B + H2O ⇀ BH+ + OH−
* Usual symbol Kb
* Usually reported as a p-value: pKa = − log10 Kb
define pH
pH = − log10{H+}
{H+} is the hydrogen ion activity
what is the ionic product of water
The equilibrium constant for the autoionization (acid dissociation) reaction of water.
H2O ⇀↽ H+ + OH−
* Usual symbol KW
* At 298 K KW ≃ 10−14
* pKa + pKb = pKW = 14
define solubility product
The equilibrium constant for the dissolution of a salt.
* Usual symbol Ksp
define azeotrope
A mixture of two or more liquids which has a constant boiling point and composition
throughout distillation. In particular, the vapour has the same composition as the liquid
define fractional distillation
Separation of a mixture by repeated evaporation-condensation-evaporation
define eutectic
A homogeneous mixture of solids that melts at a temperature that is lower than any of its
components. The eutectic mixture melts and freezes as if it were a pure compound rather
than a mixture - it does not separate into different components upon melting/freezing.
* The eutectic temperature (aka eutectic point) is the lowest possible melting tem-
perature amongst all of the mixing ratios for the components involved
define colligative properties
A property of a solution/mixture that depends on the mole fraction of solute species and
not on their chemical identity
define osmotic property
The difference in pressure between a solvent and a solution that are separated by a semi-
permeable membrane that only allows solvent molecules to cross
what is Dalton’s law
The partial pressure of a gas in a mixture is the mole fraction times the total pressure.
pA = xAp
* Each gas contributes a partial pressure to the overall pressure that is proportional
to its mole fraction.
* Dalton’s law is accurate in so far as the gases are ideal
what is Raoult’s law
The partial vapour pressure of a solvent in a mixture is proportional to its mole fraction
in the solution and its vapour pressure when pure.
pA = xAp∗
* p∗ is the vapour pressure of pure solvent.
* Solutions that obey Raoult’s law are called ideal solutions
what is Henry’s law
The partial vapour pressure of a solute in a solution is proportional to its mole fraction.
pB = xB KB
* KB is NOT the vapour pressure of pure solvent unless the solution is ideal for all
mole fractions.
* Solutions that obey Henry’s law are called ideal-dilute solutions
define entropy of mixing
The change in entropy that occurs when two or more components are mixed in an ideal
mixture (or solution).
∆mixS = −nR (xA ln xA + xB ln xB + xC ln xC …)
* The entropy of mixing is always positive.
* The formula above is accurate in so far as the mixture is ideal
define tie line
A horizontal line on the temperature-mole fraction diagram for a mixture. The ends of
the tie line indicate the compositions of the liquid and vapour phases (or solid and liquid
phases). The amounts of liquid and vapour are given by the lever rule.
what is the 0th law of thermodyamics
if 2 systems are in equilibrium with a 3rd system then they are in equilibrium with each other
what is the 1st law of thermodynamics
Δ U =0 for an isolated system and Δsys U + Δsurr U =0 for open/closed systems
Meaning conservation of energy. If a system is not isolated, then energy lost by the system is
gained by the surroundings, i.e., is transferred from system to surroundings
what is the second law of thermodynamics
Δ S≥0 for an isolated system and Δ Ssys+ ΔSsurr ≥0 for open/closed systems
Meaning in an isolated system, entropy increases until equilibrium is reached when it stays constant; entropy does not decrease in an isolated system. If a system is not isolated, then entropy lost by the system must be at least balanced by that gained by the surroundings; the total entropy of
system + surroundings never decreases.
what is the 3rd law of thermodynamics
S=0 at 0 K or a perfect crystal.
. The entropy of a perfect crystalline substance is zero at
absolute zero.
What is troutons rule
states that the (molar) entropy of vaporization is almost the same value, about 85–88 J/(K·mol), for various kinds of liquids at their boiling points.
What is Kirchhoffs law
At constant pressure, the heat capacity is equal to change in enthalpy divided by the change in temperature
What is coulombs law
the force of attraction or repulsion between two charged bodies is directly proportional to the product of their charges and inversely proportional to the square of the distance between them.
