OCR Chemistry Module 3

first ionization energy

the energy required to remove the first electron from an atom

atomic radius

the greater the distance between the nucleus and the outer shell electrons the less the nuclear attraction

nuclear charge

the more protons there are, the greater the attraction between the nucleus and outer shell electrons

electron shielding

Inner electrons repel outer electrons, reducing the nuclear attraction

large jumps in successive ionization energies for a single element are observed when passing from..

one shell to another which is closer to the nucleus

Trend in first ionisation energy down a group

atomic radius increases, more shielding electrons, nuclear attraction decreases, first ionisation energy decreases

Trend in first ionisation energy across a period

nuclear charge increases, similar shielding electrons, nuclear attraction increases, atomic radius decreases, first ionisation energy increases

sub shell trends across a period

slight fall when moving to fill a different sub shell

metallic bonding

Strong electrostatic attraction between cations and delocalised electrons

electrical conductivity of metallic bonding

they conduct as the delocalised electrons can move and carry charge

melting and boiling points of metallic bonding

most have high, due to the large amount of energy required to break the strong electrostatic attraction

solubility of metallic bonding

don't dissolve, maybe some interaction between polar solvents and the charges in the lattices, but these would lead to reactions not dissolving

Giant covalent structure

billions of atoms held together by a network of strong covalent bonds

melting and boiling points of giant covalent structure

high due to strong covalent bonds

Solubility of giant covalent structures

insoluble

electrical conductivity of giant covalent structures

non-conductors, except for graphene and graphite, where they have 3 / 4 outer shell electrons bonded to the last one is free to carry charge

Group 2 elements reactions with oxygen

metal oxide

group 2 elements reaction with water

alkaline hydroxide and hydrogen

group 2 elements reaction with dilute acid

salt and hydrogen

Group 2 reactivity down the group

ionisation energies decrease as the nuclear attraction decreases as a result of increasing atomic radius and shielding electrons

group 2 oxide compounds with water

metal hydroxide. reaction releases OH-, these hydroxides are only slightly soluble so a ppt will form

Solubility of hydroxides

Increases down the group

Group 2 compounds in agriculture

calcium hydroxide added to fields to increase their pH of acidic fields. so it neutralises forming water.

Group 2 compounds in medicine

often used as anti acids for treating acid indigestion.

halogen trend in boiling point down the group

more electrons, stronger london forces, more energy required to break these forces, boiling point increases

Halogen-halide displacement reactions

a solution of each halogen is added the the aqueous solution of another, if the halogen added is more reactive than the halogen in the aqueous solution, the solution changed colour due to a reaction that has taken place.

Halogen trend in reactivity down the group

atomic radius increases, more electron shielding, nuclear attraction decreases, reactivity decreases

chlorine and water

Cl2 + H2O -> HClO + HCl

chlorine and cold dilute aqueous NaOH

NaClO(bleach) + NaCl + H2O

Benefits of Chlorine use

purification, kills bacteria

Risks of Chlorine use

respiratory irritant in small conc. fatal in large conc., cancer causing

test for halide ions

Ag+ (aq) + X- (aq) = AgX (s)

Carbonate test

Add dilute nitric acid, effervescence of CO2 (can then bubble through lime water, turning milky)

Sulfate test

add barium nitrate (or chloride- but you cannot test for halides after) white ppt of BaSO4

Halide Test

Nitric acid and silver nitrate, then white/cream/yellow precipitate. (Chloride, bromide, iodide)

Add dilute ammonia (NH3) chloride white ppt dissolves. Add conc. ammonia bromide ppt dissolves, iodide ppt doesn't dissolve.

Ammonium test

add NaOH, when heated, ammonia gas forms, test with moist pH paper turning blue

Enthalpy

energy stored within bonds, it cannot be measure, but its changes can

activation energy

the minimum amount of energy required to start a chemical reaction

exothermic

endothermic

standard conditions

100kPa

298K (25 degrees C)

1mol dm^-3

Enthalpy change of formation

the enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions

enthalpy change of combustion

The energy change that takes place when 1 mole of a substance is completely combusted with oxygen in their standard states under standard conditions

enthalpy change of neutralisation

enthalpy change when one mole of water molecules are formed when an acid reacts with an alkali under standard conditions in their standard states

kelvin scale

0K =-273

calculating energy change

q=mc delta T

determine energy change of combustion

1 - calculate q in kJ

2- calculate amount of substance burnt

3- divide answer 1 by answer 2 (in kJmol^-1)

how accurate is the determination of energy change of combustion

heat lost to surroundings

incomplete combustion

evaporation of methanol from the wick

non-standard conditions

all but the last one would lead to a less exothermic result

spirit burner experiment (combustion energy change)

1 - measure out water and record initial temp.

2- weigh before and after adding methanol to spirit burner

3- place burner under beaker and light

4- after a bit stop experiment and immediatly record the temp. of the water

5- reweigh spirit burner

determining the enthalpy change of reaction

1- calculate q in the solution in kJ

2- calulate the amount in mol of the substance not in excess

3- divide 1 by 2 in kJmol^-1

Average bond enthalpy

The energy required to break one mole of a specified type of bond in a gaseous molecule

delta H +

endothermic, breaking bonds

delta H -

exothermic, making bonds

Why are bond enthalpies always endothermic?

Because energy has to be supplied in order to break the bonds

How can you calculate an average bond enthalpy?

Take An average of actual bond enthalpies in different environments

Why is bond breaking endothermic?

Energy must be supplied to break existing bonds

Why is bond making exothermic?

Energy is released when new bonds are formed

What are the limitations of using average bond enthalpies?

The actual energy involved in breaking and making individual bonds would be slightly different

Why would it be difficult to measure the enthalpy change of this reaction:

4C + 5 H2 = C4H10

Because carbon and hydrogen react to produce many products, so C4H10 would only be one of the many products formed.

Suggest reasons why standard enthalpy changes of combustion determined experimentally are less exothermic than calculated theoretical values?

- heat released to surroundings

- incomplete combustion

- substances may not be present in standard states

It's difficult to determine the standard enthalpy change of formation of hexane directly - suggest why?

- Many different hydrocarbons would form

- activation energy too high

- reaction too slow

- They don't react together

list important processes where hydrogen is used?

production of margarine, ammonia, Haber process

The student looked in a textbook and found that the actual value for the standard enthalpy change of combustion of propan-1-ol was more exothermic than the experimental value - suggest why?

-elements might not be in their standard states

-water evaporates

-thermal capacity of beaker ignored

-combustion could be incomplete

-propan-1-ol evaporates

hess' law

If a reaction can take place by more than one route and the initial and final conditions are the same, the total enthalpy change is the same for each route.

working out enthalpy changes

1 - construct enthalpy cycle, elements at the bottom, reactants on the left, products on the right

2 - add values of known enthalpys and calculate unknown enthalpy

buildings form up, and burn down

rate of chemical reaction

measures how fast a reactant is being used up or how fast a product is being formed

rate of reaction

change in conc. over time

The rate of a chemical reaction depends on

temperature

concentration

catalyst

surface area of solid reactants

reason for ineffective collision of particles

not the correct orientation

not sufficient energy to overcome activation energy barrier

increasing the conc. affect on the rate

increases the number of particles in the same volume, they are closer together and collide more frequently in a given period of time, therefore more successful collisions per unit of time

increasing pressure of a gas affect on rate

conc. of gas molecules increase as the same number of molecules fill a smaller volume, they are closer together and collide more frequently in a given period of time, therefore more successful collisions per unit of time

monitor gas production

or measuring cylinder, measured at timed intervals

catalyst

changes the rate of reaction without undergoing any permanent change

homogeneous catalyst

In the same state as the reactants, it reacts to form an intermediate

heterogeneous catalyst

Different physical state to reactants, they are absorbed onto the surface of the catalyst where the reaction takes place- then deabsorbed

dynamic equalibrium

rate of the forward reaction is equal to the rate of the backward reaction

Le Chatelier's Principle

States that if a stress is applied to a system at equilibrium, the system shifts in the direction that minimise the stress.

Effect of catalyst on equilibrium

no effect, merely speeds up the rates of the forward and reverse equally