C1 EXAM BRIEFINGThursday 17th May 2018, 9.00AM

ATOMIC STRUCTURE &

THE PERIODIC TABLETopic 1

Atomic Structure• ~ 100 different elements.

• Radius of around 0.1 nanometers(1x10-10m)

• Nucleus has a radius of around 1x10-14m (1/10,000th of atom)

When Who What was discovered?

Beginning 19th

CJohn Dalton Atoms=solid spheres, different atoms=different

elements

1897 J J Thomson Plum Pudding model

1909 Rutherford / Marsden

Alpha particle scattering– nuclear model of atom

Around 1913 Niels Bohr Electrons contained in shells (calcs agreed w. obs)

1920s Rutherford &others

Showed nucleus can be divided into protons

1932 James Chadwick Showed the existence of Neutrons

• Different no. protons = different element• Different no. electrons = ion• Different no. neutrons = isotope

• Atomic no. = p• Atomic mass (Relative atomic mass (Ar)) = p+n

• Electronic structure = 2,8,8,2

Type of Equations

• Balanced equation:

• Half equations:

• Ionic equations:

Mixtures

• Filtration - separate insol. solid from liquid

• Crystallisation - separate sol. solid from liquid

• Simple distillation - separate mixture of liquids

• Fractional distillation - as above

• Chromatography - separate mixture of liquids

Required Practical

Periodic Table Group no. = no. electrons in outer shell (Group 0 = full)Period no. = no. electrons shells

1st attempt- strict order of atomic weightMendeleev- left gaps for undiscovered elements, changed order from weight to suit properties

Metal = + ions Non-metals = - ions

Group 0Full outer shell = stable & unreactive (inert).Bp inc. down group.

Group 1- alkali metals1 electron outer shell = v. reactive.Fizz, bubble with water = metal hydroxide pH 1-6.More reactive down group.MP & bp increases down group.Ar increases down group.+ oxygen metal oxide.1+ ions.

Group 7- halogens7 electrons outer shell = v. reactive. Less reactive down group.Mp & bp increases down group. Ar increases down group. 1- ions.Displacement reactions.

BONDING, STRUCTURE, AND

THE PROPERTIES OF MATTER

Topic 2

Bonding

Types of covalent diagrams: Polymers:

Graphene & Fullerenes• Graphene - single layer graphite (carbon).

Delocalised electrons, conduct electricity.

• Fullerenes – hollow molecules (tubes / balls) carbon, heptagons, hexagons, pentagons.• First = buckminsterfullerene (C60, 20 hexagons + 12

pentagons)• Nanotubes = tiny carbon cylinders. Length : diameter =

v. high. Conduct heat & electricity.• Uses fullerenes:

• Medicine- cage other molecules• Catalysts- large surface area• Lubricants- machine parts & artificial joints• Strengthening materials- high tensile strength• Electronics- microchips

Nanoparticles• Course particles: dust, PM10 (particulate matter up to 10 micrometers)• Fine particles: PM25

• Nanoparticles: diameter between 1nm & 100 nm. Few hundred atoms. • Atomic diameter ~ 0.1 nm• Small molecules diameter ~<1 nm – 10 – 10,000 x larger than an atom.

• High surface area : volume = good catalysts. • As particles decrease in size, SA:V increases.• Platinum used in fuel cells

• Gold nanoparticles = lower mp than bulk material• Nanomedicine- easily absorbed• Electronics- conduct electricity- thin, light displays, small memory chips• Deodorants- silver nanoparticles antibacterial• Sun cream- more effective, better skin coverage• Cosmetics- non-oily moisturisers, deliver active ingredients to lower layers skin.

States of Matter

As energy is added to particles, kinetic energy increases, particles spread out = less dense.

• Limitations of kinetic particle theory model:• Forces of attraction between particles are not seen in

a static image.• Movement of particles in terms of direction and speed may not be accurately captured

in a static image.• Distance between particles in model may not be scaled accurately to the actual

distance between particles.

(A on graph) (C on graph)

(B on graph) (D on graph)

Solid = (s), liquid = (l), gas = (g), aqueous = (aq)

QUANTITATIVE CHEMISTRY

Topic 3

Mass (g and Mr) & MolesMass reactants = mass products Charge reactants = charge products

• Open system:• Product lost as gas See ionic equations & half equations• Reactant = atmospheric gas • Raw materials not pure• Products left behind• Reaction not finished• Unexpected reactions

1 mole (mol) = 6.02 x 1023 = the Avogadro constant

1 mole C = 12g1 mole Mg = 24g

Balancing no.s can be determined from ratio of no. moles

mass

n (Mol) Mr

Uses ± symbol

Balancing equations using reacting masses

• 4.6g of sodium reacted with 1.6g of oxygen to form 6.2g of sodium oxide (Na2O, Mr = 62). Write a balanced equation, using the reacting masses.

1. Divide the mass of each substance by its relative formula mass (Mr) to find the number of moles.4.6g Na / 23 = 0.2 moles 1.6g O / 16 = 0.1 moles 6.2g Na2O / 62 = 0.1 moles

2. Divide the number of moles of each substance by the smallest number of moles in the reaction0.2 moles / 0.1 moles = 2 0.1 moles / 0.1 moles = 1 0.1 moles / 0.1 moles = 1

3. If the numbers aren’t whole numbers, multiply all the numbers by the same amount so that they become whole numbers

They’re whole numbers this time

4. Write the balanced symbol equation for the reaction by putting these numbers in front of the chemical formulae.

____ Na + ____ O Na2O2 Na + O Na2O

Example Calculation

Example,:A compound contains 75% C and 25% H. What is its empirical formula?

1 C H

2 Amount 75 25

3 Convert to moles ( /Mr)

/12 = 6.25 /1 = 25

4 Calculate mole ratio (divide by smallest number)

6.25/ 6.25 25/6.25

5 = 1 = 4

5 Empirical formula C H4

1. list the elements

2. underneath put mass or %

3. divide by Mr to get mole ratio

4. simplest ratio of moles

5. formula

Example Calculation

Concentration(g/dm3). 1 mol = 6.02 x 1023 / dm3

1 dm3 = 1 litre = 1000 cm3

‘Conc. acid’ = acid has a very large mass per volume of H+ ions in it.

Example Calculation

Calculate the number of grams of NaOH needed to make 50cm3 of NaOHwith a concentration of 2 mol/dm3.

n = ?c = 2 mol/dm3

v = 0.05dm3

n = c x Vn = 2mol/dm3 x 0.05dm3

n = 0.1 mole NaOH

First Triangle;

Second Triangle;Mass = n x MrMass = 0.1 moles x 40Mass = 4g

Mass = ?n = 0.1 molesMr = 40

mass

n (Mol) Mr

Limiting Reactants, % Yield, Atom Economy, Moles of Gases• Limiting reactant- reactant that is fully used up, limits amount of product.

• Other reactant is ‘in excess’. • Can determine mass / no. moles of reactant/ product used from balanced

equation.

• % Yield

• Atom Economy

• Moles of gases• One mole of any gas at room temperature and pressure (20oC and 1

atmosphere pressure) is 24 dm3.

CHEMICAL CHANGESTopic 4

Reactivity• Metals form positive ions

• Potassium more reactive than lithium• Electrons attracted to positive protons in the nucleus.• Potassium has more electron shells than lithium.• Outer shell electrons are further away from nucleus.• Therefore attraction between outer shell electrons and nucleus

is weaker so the electrons are lost more easily.• When potassium loses electrons it is oxidised and becomes a K+/

1+ ion.• Because potassium loses its outer shell electron easily, it forms

ions quickly (and strongly ionises water) and reacts quickly and more vigorously.

• The more easily an atom can become ionised, the more quickly that ion can react and therefore the metal is more reactive.

More ions = more reactive

• Increasing concentration makes an acid more reactive because it contains more ions!

Also-displacement

reactions

Redox

• Oxidation • = bonding with oxygen

• Metal + oxygen metal oxide• Loss of electrons

• Metals = positive ions

• Reduction• = losing oxygen

• Metal oxide + carbon metal + carbon dioxide• Gain of electrons

• Non-metals = negative ions

Half equations:2 Fe3+ + 6e- 2 Fe3 O2- 3 O + 6e-

• Unreactive metals e.g. = native metals

• Metals < reactive than C extracted from oxides by reduction with carbon

• Metals > reactive than C extracted from oxides by electrolysis

Acid + reactions• acid + metal ? + ?

• acid + alkali (metal hydroxide) ? + ?

• acid + metal oxide ? + ?

• acid + metal carbonate ? + ? + ?

• Hydrochloric acid will always produce ? salts

• Sulfuric acid will always produce ? salts

• Nitric acid will always produce ? salts

Making Salts

Required Practical See also Separation Techniques Topic 1

Titration

A ‘strong acid’ is one that is completely ionised in water-

More H+ ions are released.

A ‘weak acid’ is one that only partially ionises in water

Fewer H+ ions are released

Required Practical

See also Quantitative Chemistry

Required Practical

TitrationSee also Quantitative Chemistry

Electrolysis of Molten Electrolyte

• Ions must be free to move

• Charge of ions can be determined by group no. (except transition metals- some have 2+ oxidation states)

• Electrolysis = £

• Cryolite used to lower mp of aluminium oxide

Can also be a solution (dissolved solute in solvent)

Copper sulfate: Cu2+ + 2e-CuSO4

2- SO4 + 2e-

The electrons leave the sulfate and attach to the copper.

Lead bromide Pb2+ + 2e- Pb2Br- Br2 + 2e-

Together; Pb2+(aq) + 2Br-

(aq) Pb(s) + Br2 (aq)

Aluminium oxide 4Al3+ + 12e- 4Al6O2- 3O2 + 12e-

Electrolysis of Solution of Ions•Sodium chloride solution (brine)•At the negative cathode: Na+

(aq) + 2e-Na(s)

•At the positive anode: 2Cl-(aq)Cl2 (g)+ 2e-

•What is the overall equation?• 2Na+

(aq) + 2Cl-(aq) + 2H+

(aq) + 2OH-(aq) 2NaOH + Cl2 (g) + H2(g)

• Hydrogen is used as a fuel and for making ammonia.

• Chlorine is used to kill bacteria in water, and to make bleach and plastics.

• Sodium hydroxide is used to make soap and bleach.

Electrolysis of Solution of Ions• Negative electrode:

• The metal will be produced if < reactive than hydrogen• Hydrogen will be produced if the metal > reactive than hydrogen

• Positive electrode:• The halide will be released if there is one in the solution• Oxygen will be produced if not

SolutionPositive electrode

(anode)Negative electrode

(cathode)

Copper (II) chloride chlorine copper

Copper (II) sulfate oxygen copper

Sodium chloride chlorine hydrogen

Sodium sulfate oxygen hydrogen

Required Practical

ENERGY CHANGESTopic 5

Energy Change

Exothermic• Transfers energy from chemicals to

surroundings, T surroundings increase.

• Making new bonds is exothermic.

• E.g. combustion, many oxidation reactions, neutralisation, self-heating cans, hand warmers.

Endothermic• Transfers energy from surroundings

to chemicals, T surroundings increase.

• Breaking bonds is endothermic.

• E.g. thermal decomposition, citric acid + hydrogencarbonate, sports injury packs, photosynthesis.

Overall energy change can be calculated by energy taken in breaking bonds – energy given out making new bonds.+ = endothermic, - = exothermic.

Energy Changesa Stand the polystyrene cup in the beaker.

b Use the measuring cylinder to measure out 5 cm3 of hydrochloric acid and pour it into the polystyrene cup.

c Measure the initial temperature of the hydrochloric acid and record it in a suitable table.

d Add 5 cm3 of sodium hydroxide solution. Stir with the thermometer and record the maximum or minimum temperature reached.

e Work out the temperature change and decide if the reaction is exothermic or endothermic.

Required Practical IV: volume sodium hydroxide (cm3 )DV: Maximum temperature (°C)CV: volume hydrochloric acid (cm3), temperature of surroundings (°C)

Insulated cup & lid- could be a gapDigital thermometer = more precise.

Cells & Batteries• Cell made of

• two different metals• metals separated from each other by electrolyte• metals connected by wires through which electrons can flow

• Batteries = 2+ cells = greater voltage.

• Non-rechargeable = chemical reactions stop when one of the reactants has been used up; alkaline batteries

• Rechargeable = chemical reactions are reversed when external electrical current supplied.

Determining Reactivity from Voltage Results• Bigger difference in reactivity between electrodes= bigger voltage of cell

• Conclusions;• Copper is always the least reactive metal• Zinc – iron – chromium – tin – copper

• Negative voltages happen when more reactive metal on other electrode

• Iron must be ‘middle’ reactivity• Aluminium – zinc – iron – tin – lead

• Voltages of multiple cells add together to make battery’s voltage.

Battery voltage = 1.71V

Fuel Cells• hydrogen + oxygen water

• 2H2(g) + O2(g) 2H2O(l)

• Fuel (hydrogen) enters cell on one side, becomes oxidised (reacts with oxygen)- sets up potential difference.

Fuel CellsAdvantages

• No greenhouse gases, nitrogen oxides, sulfur dioxide or carbon monoxide• Only by-products are water + heat.

• Electric cars• Don’t produce many pollutants

either, but batteries are more polluting to dispose of (highly toxic metal compounds).

• Batteries are rechargeable but limit to no. times can be done.

• Batteries more expensive than to make than fuel cells

• Batteries store less energy than fuel cells- need to be recharged more often- takes a long time.

Disadvantages

• Hydrogen = gas- large volume

• Hydrogen = explosive when mixed with air- difficult to store safely.

• Hydrogen made from hydrocarbons (non-renewable) or by electrolysis (electricity often generated by burning fossil fuels).