Physical Chemistry Facts

Oxidation: Fe2+ forms
Fe3+ + e-
Reduction: MnO4- + 8H+ +5e- forms
Mn2+ + 4H2O
Redox: MnO4- + 8H+ +5Fe2+ forms
Mn2+ + 5Fe3+ + 4H2O
Redox titrations using MnO4-
Solution ACIDIFIED with sulphuric acid, not hydrochloric acid which REACTS with MnO4- and produces CHLORINE GAS. MnO4- is in BURETTE and iron is in the CONICAL flask. Stop when there is a PERMANENT PINK colour.
Redox titration
Transfer of electrons not protons as in normal titrations. Titrate an oxidising agent against a reducing agent.
Why use MnO4- (potassium manganate)?
Self-indicating; Mn2+ ions pale pink, MnO4- ions deep purple. When in excess, MnO4- ions form deep purple solution.
Oxidation change from MnO4- to Mn2+
Oxidation change from Fe2+ to Fe3+
Ratio of Fe2+ to MnO4-
Large Kstab value
Equilibrium lies to the RIGHT. Indicates complex ion is MORE STABLE. Ion is formed MORE EASILY.
Small Kstab value
Equilibrium lies to LEFT. Indicates complex ion is LESS STABLE. Ion is MORE DIFFICULT to form.
Tables of Kstab values
Compare stability of complexes with those containing H2O ligands
Why is water left out of Kstab equation
All species are dissolved in water which is in large excess and concentration is constant
Why is carbon monoxide fatal
Binds to haemoglobin more strongly than oxygen at the same binding site; irreversible ligand substitution. Tissues do not receive oxygen so cannot respire.
What is carbon monoxide
Formed in incomplete combustion of carbon compounds/ tobacco. Colourless/ odourless.
Protein complex composed of 4 polypeptide chains and 4 haem groups in red blood cells.
Haem group
Complex ion with Fe2+ ion at centre; 4 coordinate bonds to nitrogen atoms of the haem structure; coordinate bond to protein globin; coordinate bond for oxygen.
Why there are more water ligands than chloride ligands in a complex
Chloride ligands are larger than water ligands and have larger forces of repulsions so fewer can fit around a metal ion
Stages of ammonia addition to [Cu(H2O)6]2+ solution
Stage one; small amount of ammonia added; formation of Cu(OH)2; pale blue precipitate. Stage two; excess ammonia added; formation of [Cu(NH3)4(H2O)2]2+; dark blue precipitate
Requirements for optical isomers of complex ions
OCTAHEDRAL. Three bidentate ligands. Two bidentate ligands and two monodentate ligands. One hexadentate ligand.
(Ethylenediaminetetraacetic acid) Hexadentate ligand with a 4- charge
Uses of EDTA
Chelating agent; binds to metal ions so dereases their concentration in a solution. DETERGENT; binds to calcium and magnesium to reduce hardness of water. STABILISER; removes metal ions catalysing oxidation in some foods. MEDICAL; prevents clotting and relieves lead and mercury poisoning.
Cis isomerism in complex ions
Monodentate ligands are adjacent; 90 degrees apart
Trans isomerism in complex ions
Monodentate ligands are opposite; 180 degrees apart
Examples of bidentate ligands
ethane-1,2-diamine and ethanedioate
Criteria for cis-trans isomerism in complex ions
Two types of ligand. 4 of one monodentate and 2 of another monodentate. 2 bidentate and 2 monodentate. 2 of one monodentate and two of another.
Shape of complex ions with 4 monodentate ligands of 2 types
Square planar
Bond angles in an octahedral
90 degrees
Haber process
Forms ammonia from nitrogen and hydrogen for fertilisers
Haber process equation
N2+3H2 forms 2NH3
Contact process
COnverts sulphur dioxide to sulphur trioxide in the manufacture of sulphuric acid.
Contact process equation
2SO2+O2 forms 2SO3
Hydrogenation of alkenes
Forms saturated compounds from unsaturated compounds
Decomposition of hydrogen peroxide
Forms water and oxygen
Decomposition of hydrogen peroxide equation
2H2O2 forms 2H2O+O2
Ways transition metals are catalysts
Provide a surface for adsorption and desorption. Can change oxidation states by gaining or losing electrons; bind to reagents to form intermediates in chemical pathway with lower activation energy.
Physical properties of transition metals
Shiny, high density, high melting point, high boiling point, can conduct electricity as form giant metallic lattices with delocalised electrons when solid.
Alloy of nickel and copper
Makes silver coins
Titanium is a component of
Joint replacement parts
Iron is used in
Telephone and post boxes, construction of buildings and bridges
Why do transition metals form coloured ions
Partially filled D orbitals
Oxidation state formed by every transition metal
+2; electrons lost from the 4s orbital
Highest oxidation state of a transition metal is often found in
Strong oxidising agents
Why scandium is not a transition metal
Forms a 3+ ion in which d orbitals are empty
Why zinc is not a transition metal
Forms a 2+ ion in which d orbitals are full
Transition elements that do not follow the Aufbau principle
Chromium and copper; electron repulsion between outer electrons is minimised; increased stability of atoms
Electron configuration of chromium
Each orbital contains one electron
Electron configuration of copper
Each 3d orbital is completely full, only one electron in 4s
Transition elements lose electrons from orbitals
4s and 3d, 4s first
Limitations of hydrogen economy
Must be publicly and politically accepted. Logistical problems in handling and maintenance of hydrogen systems. More energy may be used in making hydrogen, by electrolysis, than is saved by its use.
Limitations of hydrogen fuel cells
Large scale storage and transportation of hydrogen is not cost or energy efficient. Storing pressurised liquid is not feasible. Adsorbers and absorbers of hydrogen have a limited lifetime. Fuel cells have a limited life time and cost to replace and dispose of. Fuel cells are produced with toxic chemicals.
Storage of hydrogen
Liquid under pressure at a low temperature; logistical problems. Can be adsorbed onto surfaces of solids. Can be absorbed within some solids.
Efficiency of fuel cell vehicles
Petrol engines are 20% efficient at converting chemical energy, hydrogen fuel cell vehicles are 40-60% efficient
Hydrogen rich fuels
Methanol, natural gas, petrol; mixed with water and converted into hydrogen gas by on board reformer, which operates at 250-300 degrees
Pollution of hydrogen fuel cells
Only produce a small amount of carbon dioxide and air pollutants, do not produce carbon monoxide in incomplete combustion.
Advantages of methanol fuel cells
Liquid, so is easier to store than hydrogen gas. Methanol can be generated from biomass.
Hydrogen oxygen fuel cell
Hydrogen and oxygen flow in while water flows out. Continue to operate as long as oxygen and hydrogen are input.
Negative terminal of hydrogen oxygen fuel cell
2H2O+2e- forms H2+2OH-
Positive terminal of hydrogen oxygen fuel cell
1/2O2+H2O+2e- forms 2OH-
Overall cell reaction of hydrogen oxygen fuel cell
H2+1/2O2 forms H2O
In electrochemical cells, electrons flow from
Negative anode to positive cathode
Main types of electrochemical cells
Non-rechargeable batteries. Rechargeable batteries. Fuel cells.
Non-rechargeable batteries
Provide electrical energy until chemicals have reacted to such an extent that the voltage falls
Rechargeable batteries
Recharging reverses the cell reaction so the chemicals in the cells are regenerated and can be reused
Fuel cells
Use external supplies of fuels as oxidants which are consumed and must be continuously supplied to continue to provide electrical energy
In electrochemical cells, electrons are provided by
The anode; the negative terminal where oxidation occurs
In electrochemical cells, electrons are donated to
The cathode; the positive terminal where reduction occurs
EMF of the cell
Standard electrode potential of POSITIVE TERMINAL-Standard electrode potential of NEGATIVE TERMINAL
Conventional redox equilibria
Show electrons on left
Feasible electrochemical reactions take place between
Species on the right of the equilibrium with the more negative EMF and on the left of the equilibrium with the more positive EMF
The larger the difference between EMFs
The more likely it is that a reaction will take place
If the difference between EMF values is less than 0.4V
A reaction is unlikely to be feasible
Why an electrochemical reaction may not take place
Too slow due to a high activation energy. Reactions may take place under conditions that are not standard. Species may not be aqueous.
Measuring standard electrode potential
Standard half cell is connected to a standard hydrogen half cell. Standard half cell; metal strip in metal ion solution at 1 moldm-3. Connected by a wire and salt bridge to hydrogen half cell which contains a platinum electrode, acid at 1 moldm-3 and hydrogen gas entering at 298K and 1 atm.
Metal ion/ metal ion half cell
Equimolar solution containing ions of the same element in different oxidation states and an inert platinum electrode connected to a wire.
Hydrogen half cell
Solution of 1moldm-3 acid, hydrogen gas entering at 298K and 1atm, an inert platinum electrode coated in platinum black which is linked to a wire.
Cells from metal/ metal ions
Two solutions of metal ions at 1moldm-3, each containing their metal electrode connected by a wire to a high resistance voltmeter. A salt bridge connects the two solutions
Salt bridge
Filter paper soaked in aqueous solutions of ionic compounds that do not react with either half cell solution. Allows transfer of ions between solutions
Ionic compounds for salt bridges
KNO3 or NH4NO3
Negative enthalpy, positive entropy
Always negative G, feasible reaction
Positive enthalpy, negative entropy
Always positive G, never feasible reaction
Negative enthalpy, negative entropy
Negative G at low temperatures, feasible at low temperatures
Positive enthalpy, positive entropy
Negative G at high temperatures, feasible at high temperatures
Most exothermic reactions are
Spontaneous processes happen when
G is negative
Spontaneous processes lead to
Lower energy and increased stability
Entropy change is positive when
Change makes a system more random
Entropy change is negative when
Change makes a system less random
Entropy increases when
There is an increase in gas molecules, a solid lattice dissolves or when the temperature increases
Energy tends to change from
Being localised to more dilute
Factors affecting hydration enthalpy
IONIC SIZE; smaller, makes hydration enthalpy more negative; exerts more forces of attraction on water so releases more energy. IONIC CHARGE; small highly charged ions; makes hydration enthalpy more negative; greater force of attraction on water molecules
Factors affecting lattice enthalpy
IONIC SIZE; smaller; pack more closely together; attract each other more strongly; lattice enthalpy becomes more negative. IONIC CHARGE; small highly charged ions; attract eachother strongly; more negative lattice enthalpy
Sum of clockwise enthalpies
Sum of clockwise enthalpies
Lattice enthalpy
Formation-(Atomisation+ionisation+electron affinity)
Enthalpy changes with upward arrows
Enthalpy changes with downward arrows
Features of lattice enthalpy
EXOTHERMIC; bonds formed. Measure of IONIC BOND STRENGTH. Covalent structures do not have lattice enthalpy. Cannot be measured DIRECTLY; can’t make exactly one mole of an ionic lattice
Enthalpy of formation
Usually exothermic; bonds formed
Enthalpy of atomisation
Usually endothermic; bonds broken
Ionisation energy
Endothermic; electron must overcome attraction to nucleus
Second electron affinity
Endothermic; electron repulsion by negative 1- ion must be overcome
First electron affinity
Exothermic; electron attracted by nucleus
Enthalpy of neutralisation for strong acids
Very similar values because they all dissociate to the same extent
Titration process
Acid in conical flask, base in burette
Titration curve
Add base, pH increases slightly and acid is in great excess. When near the equivalence point the pH increases more rapidly. A vertical section caused by one drop of base is the equivalence point. Further base added causes little change in pH as base is in excess
Have one colour in their acid form and one colour in conjugate base form. When at end point, equal amounts of both exist forming a mixed colour.
Choosing an indicator
End point must be as close as possible to equivalence point
Strong acid strong base titration
Start low and end high
Strong acid weak base titration
Start low and end low
Weak acid strong base titration
Start high and end high
Weak acid weak base titration
Start high and end low
Carbonic acid
H2CO3, weak acid
HCO3-, conjugate base
Healthy blood pH
Condition if blood pH falls below 7.35
Condition if blood pH rises above 7.45
Addition of acid to blood
Reacts with conjugate base HCO3- and is removed from blood, shifting equilibrium to the left
Addition of alkali to buffer in the blood
Reacts with H+ ions to form water, causing H2CO3 equilibrium to shift to the right
How carbonic acid is removed
Converted by an enzyme to aqueous carbon dioxide, which is converted to gas in the lungs and exhaled
Addition of acid to a buffer
Reacts with conjugate base and shifts equilibrium to the left, removing the H+ ions
Addition of alkali to a buffer
Reacts with H+ ions to form water, shifting the equilibrium to the right to restore the H+ ions that have reacted
Buffer solution ingredients
Weak acid and its salt or weak acid and aqueous alkali
Strength of base
measure of its dissociation to form OH- ions
Indices of H+ and OH-
Add up to -14
pH of neutrality
Weak acid calculation assumptions
H+ is less than HA. H+ is equal to A-, so equal H+2. As few HA molecules have dissociates, HA concentration is the same as it was before dissociation.
Low pH
Large H+ concentration
High pH
Small H+ concentration
Each whole number pH value
is a factor of 10
Hydronium ions
H3O+; dative covalent bond between proton and lone pair on oxygen
Ionic equations
Solid ionic compounds are shown undissociated, Cancel out species that do not change.
Acid and carbonate
Salt, carbon dioxide and water
Acid and base/ alkali
Salt and water
Acid and metal
Salt and hydrogen
Changes in concentration during equilibrium
Concentrations in Kc equation gives a different ratio. Equilibrium position must shift to restore ratio back to original Kc value.
Increase on bottom of Kc and decrease on top
Shifts from right to left
Increase on top of Kc and decrease on bottom
Shifts from left to right
Changes in pressure during equilibrium on Kc
Pressure increase will increase concentrations of both reactants and products. It mostly increases the side with greater number of molecules. This changes the ratio of concentrations on the Kc equation so changes the value of Kc. Equilibrium shifts position to restore ratio back to original Kc value.
How catalysts affect Kc
THEY DON’T, but they increase the rate of the forward and backward reaction.
Equilibrium constant Kc
Large value; reaction is product favoured; equilibrium lies to the right. Endothermic forward reaction; increase in temperature increases Kc. Exothermic forward reaction; increase in temperature decreases Kc. Kc equation can be written from balanced equation.
Rate constant k
Large value; fast rate of reaction. Increases with an increase in temperature because this increases the reaction rate. Can only be found from experimental data that shows the rate determining step.
Compromising Kc and k
k must be large enough to have an efficient reaction but Kc must be large enough to be product favoured.
Temperature increase on an exothermic reaction
Kc decreases, equilibrium shifts left and becomes reactant favoured
Temperature increase on an endothermic reaction
Kc increases, equilibrium shifts right and becomes product favoured
Kc value of 1
equilibrium is half way between reactants and products
Kc value of more than 1
Reactant is product favoured and products on the right hand side of the equation are predominant at equilibrium
Kc value of less than 1
Reactant is reactant favoured and reactants on the left hand side of the equation are predominant at equilibrium
Kc equation
Concentration of products/ concentration of reactants
Dynamic equilibrium
Forward and reverse reaction take place at the same rate. Concentration of reactants and products are constant.
Units of Kc
Cancel out moldm-3
Powers in the rate equation show
Number of molecules of this species reacting in the rate determining step
When the temperature is increased by 10 degrees
The rate doubles; more particles are able to exceed activation energy
Zero order rate concentration graph
Horizontal line
First order rate concentration graph
Upward diagonal line
Second order rate concentration graph
Upward curve line
Determining initial rates from concentration time graphs
Draw tangent at t=0 and find the gradient of this tangent
Clock reactions
Show the time for a certain amount of a product to be formed, show visual change such as formation of a PRECIPITATE, disappearance of a SOLID or a change in COLOUR.
Initial rate is proportional to
Zero order concentration time graphs
Downward diagonal line
Zero order half life
Decreases with time
First order concentration time graphs
Gentle downward curve
First order half life
Constant; exponential decay
Second order concentration time graphs
Steep downward curve
First half life
measured from 0
Second half life
Measured from 0-first half life
Units of rate constant
Cancel down from moldm-3 for reactants, moldm-3s-1 for rate
Overall order
Sum of the individual orders
Affect of concentration on rate
Greater concentration; more collisions per second; faster reaction
Rate as reaction proceeds
Decreases as concentrations of reactants decrease; fewer collisions per second
Calculating rate from concentration time graphs
Draw a tangent to the curve and measure its gradient
Buffers are most effective when
Hydration enthalpies are always

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