Equilibria Flashcards
equilibria intro
reactions do not go to completion
forwards and backward reactions
reactions are dynamic
how do we understand equilibria?
- concentrations
- partial pressures in gas phase reactions
- thermodynamics
at equilibrium
- a reaction has not stopped
- forwards and back are equal
- concentration of species is constant
- partial pressures of gases are constant
- no net change in mixture composition
activity
a measure of concentration that takes into account the interactions in the solution
- concentration can affect physical and chemical properties
- extent and nature of solute-solute and solute-solvent interactions
- structure of solution species
ionic strength
summative measure of the charge environment
- a characteristic of solution
I = 1/2 sum mi x zi^2
mi is molality or molarity
zi is the charge on the ion
have happens at high ionic strengths
greater than 0.1M
effective and nominal concentrations are not the same
what does activity represent?
- the non-ideal behaviour brought about by intermolecular forces
- can represent a decrease in ion concentration because of counter-ions in the immediate vicinity
in a dilute solution there are electrostatic forces between ions, in a high concentration of salts, there is a weakened attraction because of interference with other ions
equation for activity
ai = (yi mi)/standard mi
it is dimensionless (no units)
yi is the activity coefficient which corrects for non-ideal behaviour
standard as the solution has been indefinitely diluted at standard molality and P = 1 bar where solute interactions are negligible. standard mi = 1
other activity equations
ai = yimi (because standard mi is 1)
ai = mi (approx) in dilute solutions as yi = 1
ai = ci (approx) as molarity is a common measure of concentration so activity = concentration (mol/L)
are there some more activity equations to learn?
yes
what are the important points of activity?
- activity of a dilute solution is basically equal to the numerical value of its concentration (mol/L)
- the activity of an ideal gas is basically equal to the value of its partial pressure (1 bar)
- the activity of any pure substance in its standard state is 1
what is the equilibrium constant?
concentrations when equilibrium is established
notes on Kc
- temperature dependent (measure temp)
- pure solids and liquids are not included (activity = 1)
- dimensionless
Kc values
less than 10^-3 reactants favoured
greater than 10^3 products favoured
between these points, at equilibrium
Kp
equilibrium constant in terms of pressure
Ka
the acid dissociation constant (Kb for a base)
Kw
water self-ionisation constant (involved with Ka and Kb)
Ksp
solubility product constant for when salts hardly dissolve
Q > k in supersaturated and precipitation
Kstab
or beta, the overall stability constant
Van’t Hoff equation
ln k = -(standard enthalpy)/RT + c
this is the linear form
can be modified for two temperatures (exothermic if ln k increases as 1/T increases)
Le Chatelier’s Principle
when a change is made to the conditions of a dynamic equilibrium the system moves to counteract the change, causing changes in quantities of reactants and products
concentration
add reactant or remove product - moves right with no effect on Kc
add product or remove reactant - moves back (no effect on Kc)
pressure
increase pressure or decrease volume - moves towards fewer gas molecules
decrease pressure or increase volume - moves towards more gas
no effect on Kc
temperature
increase - endothermic direction (right for endo, left for exo) Kc increases for endo and decreases for exo
decrease - exothermic direction (left for endo, right for exo) Kc decreases for endo and increases for exo
catalyst
add or remove - no effect on direction or KC
an acid is
a proton donor
a base is
a proton acceptor
pH
how acidic or basic a solution is
pH = -log[H3O+]
[H3O+] is the concentration of an acid
pOH
for OH-
pOH = -log[OH-]
[OH-] is concentration of a base
conversions between pH and pOH
[H3O+] x [OH-] = 1 x 10^-14
pH + pOH = 14
strong acids
- completely dissociate when dissolved in water
- hydrohalic: HCl, HBr (not HF)
- oxoacids: HNO3, H2SO4, HClO4 (O atoms outnumber H by at least 2)
strong bases
- soluble compounds containing O2- or OH-
- cations of usually most active metals
- M2O or MOH (M = group 1A)
- MO or M(OH)2 (M = 2A metals)
weak acids
- do not completely dissociate
- hydrohalic acid HF
- H is not bound to O or halogen
- oxoacids where more O than H
- organic acids such as RCOOH
weak bases
- electron rich nitrogen (lone pairs) such as NH3
- amines (RNH2)
polyphonic acids
- acids that are capable of donating more than one proton
- oxalic and carbonic acid
- pKa (2) and (1)
pKa is
the pH at half the volume to equilibrium
pKa and Ka formula
pKa = -log(subscript 10) Ka
conjugate pairs
acid 1 pairs with base 1
- strong acid gives a weak base
- weak base gives a strong acid
ka and kb formulae
ka x kb = [H3O+][OH-] = kw
pKa + pKb = pKw = 14
can you use ICE to work out pH?
yes
adding a base to an acid
adding acid to base
gets more basic
gets more acidic
equivalence point
acid has completely reacted with the base
strong base with strong acid
- high pH to low pH
- equivalence at 7
strong acid with strong base
- low pH to high pH
- eqivelecne at 7
weak acid and strong base
- starts less acidic than strong acid
- equivalence is basic
- halfway to equivalence point, the concentrations are equal and this is pKa
weak base and strong acid
- Strats less basic than strong base
- equivalence is acidic
- halfway to equivalence is point is base and conjugate acid concentrations equal pKa
buffer
solutions that maintain an approximately constant pH even when small amounts of acid or base are added
what are bases usually made from? how do they work?
- weak acid and its conjugate base
- weak base and its conjugate acid
- when small amounts of H3O+ or OH- are added, a small amount of one buffer is converted to the other
- ions have little effect on pH
Henderson-Hasselbach equation
- rearranges ka equation
- designs buffers with certain pH ranges
pH = pKa + log ([A-]/[HA])
this helps determine the buffer range
pH = pKa + log [10]/[1] to pKa + log [1]/[10] (acid over base)
pH range = pKa +/- 1
but this does not tell us about how resistant a buffer is to pH change
buffer capacity
beta = (dCb/dpH) = - (dCa/dpH)
the amount of acid dCa or base dCb that must be titrated to give a defined pH change
- big beta values indicate a resistant buffer
max buffering occurs when the pH = pKa
ocean acidification
- more acidification in colder water
- partial pressure increases in atmosphere and CO2 moves into oceans
- CO2 becomes carbonic acid which has acidic H+ ions
- shells, coral and algae with calcium break down
buffering in natural waters
the process that has been unbalanced:
- buffering of CO2 and dissolved mineral carbonate control pH
redox reactions
transfer of electrons from one substance to another
oxidation
loss of electrons (OIL)
- electrons are a product
reduction
gain of electrons (RIG)
- electrons are a reactant
balancing in acidic conditions
- identify reactants and products for half equations
- balance all atoms except O and H
- balance O with H2O
- balance H with H+
- balance net charge of reactions by adding e-
- balance e- in half equations so that e- gained = e- lost
- add balanced equations and cancel
- check stoich
balancing in basic conditions
- follow rules for acidic conditions
- add OH- to reactants and products based on H+
- combine OH- and H+ to form water
- cancel H2O on both sides
- check stoich and charge
spontaneous redox equations
- potential difference between half cells, energy is gained (ei. chemical to electrical)
- batteries and galvanic cells
non-spontaneous equations
- potential differences (ie. electrical energy) is required to overcome delta G for ox and red reactions to proceed (ie. electrical to chemical).
- electrolytic cells
potential difference
difference in electrons between two half cells and is measured in volts
1 V = 1J/C
oxidation numbers
hypothetical charge that an individual atom/ion would possess in a molecule if the shared electrons in the covalent bond were assigned to the most electrogentaive element in the bond
ox no rules
- zero for neutral atoms or molecules
- charge on an ion
- charges add up to charge on molecule
- F is always -1
- H is +1 except in metal hydrides
- O is -2
- oxidation: increases ox. no
- reduction: decrease in ox. no
anode
negative electrode, electrons produced, oxidation occurs, cations produced
cathode
positive, electrons gained, reduction occurs, cations are reduced to solids etc.
salt bridge
- inert salts such as KCl
- move and balance change
- required for current
galvanic cell
chemical to electrical energy
electrodes
- metal ion/metal where the metal is an electrode and dissolves
- metal ion/metal ion where an inert electrode such as graphite or platinum is used and both metals are dissolved
- non-metal ion/gaseous nonmetal where glass electrode incorporates platinum electrode and the non-metal ion is dissolved
conditions of the electrochemical series
25 degrees, 1x10^5 Pa, solution concentration of 1
using the electrochemical series
- predicts spontaneity
- looking for a z shape
- strong oxidants become weak reductants at the top left to right (reduced)
- strong reductants become weak oxidants at bottom right to left (oxidised)
cell potential
standard E of cell = E reduction - E oxidation
positive for galvanic and negative for electrolytic
electroplating
not spontaneous, electrons are supplied to make the reaction go the other way, polarities are switched (cathode is negative)
- electroplating
- recharging batteries
