C1 EXAM BRIEFING Thursday 17 th May 2018, 9.00AM
ATOMIC STRUCTURE & THE PERIODIC TABLE Topic 1
Atomic Structure ~ 100 different elements. Radius of around 0.1 nanometers (1x10-10 m) Nucleus has a radius of around 1x10-14 m (1/10,000 th of atom) When Who What was discovered? Beginning 19 th C John Dalton 1897 J J Thomson Plum Pudding model 1909 Rutherford / Marsden Atoms=solid spheres, different atoms=different elements Alpha particle scattering nuclear model of atom Around 1913 Niels Bohr Electrons contained in shells (calcs agreed w. obs) 1920s Rutherford & others 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 Showed nucleus can be divided into protons 1932 James Chadwick Showed the existence of Neutrons
Type of Equations Balanced equation: Half equations: Ionic equations:
Required Practical 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
Periodic Table Metal = + ions Group no. = no. electrons in outer shell (Group 0 = full) Period no. = no. electrons shells Non-metals = - ions Group 1- alkali metals 1 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- halogens 7 electrons outer shell = v. reactive. Less reactive down group. Mp & bp increases down group. Ar increases down group. 1- ions. Displacement reactions. 1 st attempt- strict order of atomic weight Mendeleev- left gaps for undiscovered elements, changed order from weight to suit properties Group 0 Full outer shell = stable & unreactive (inert). Bp inc. down group.
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 (C 60, 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, PM 10 (particulate matter up to 10 micrometers) Fine particles: PM 25 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.
Solid = (s), liquid = (l), gas = (g), aqueous = (aq) States of Matter As energy is added to particles, kinetic energy increases, particles spread out = less dense. (A on graph) (C on graph) (B on graph) (D on graph) 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.
QUANTITATIVE CHEMISTRY Topic 3
Mass (g and M r ) & Moles Mass reactants = mass products Open system: Product lost as gas Reactant = atmospheric gas Raw materials not pure Products left behind Reaction not finished Unexpected reactions Uses ± symbol Charge reactants = charge products See ionic equations & half equations Balancing no.s can be determined from ratio of no. moles 1 mole (mol) = 6.02 x 10 23 = the Avogadro constant mass n (Mol) M r 1 mole C = 12g 1 mole Mg = 24g
Example Calculation Balancing equations using reacting masses 4.6g of sodium reacted with 1.6g of oxygen to form 6.2g of sodium oxide (Na 2 O, 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 Na 2 O / 62 = 0.1 moles 2. Divide the number of moles of each substance by the smallest number of moles in the reaction 0.2 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 Na 2 O 2 Na + O Na 2 O
Example Calculation Example,: A compound contains 75% C and 25% H. What is its empirical formula? 1. list the elements 2. underneath put mass or % 3. divide by Mr to get mole ratio 4. simplest ratio of moles 5. formula 1 C H 2 Amount 75 25 3 Convert to moles ( /Mr) 4 Calculate mole ratio (divide by smallest number) /12 = 6.25 /1 = 25 6.25/ 6.25 25/6.2 5 5 = 1 = 4 5 Empirical formula C H 4
Example Calculation Concentration (g/dm 3 ). 1 mol = 6.02 x 10 23 / dm 3 1 dm 3 = 1 litre = 1000 cm 3 Calculate the number of grams of NaOH needed to make 50cm 3 of NaOH with a concentration of 2 mol/dm 3. First Triangle; n =? c = 2 mol/dm 3 v = 0.05dm 3 Second Triangle; Mass =? n = 0.1 moles M r = 40 n = c x V n = 2mol/dm 3 x 0.05dm 3 n = 0.1 mole NaOH Mass = n x M r Mass = 0.1 moles x 40 Mass = 4g mass n (Mol) M r Conc. acid = acid has a very large mass per volume of H + ions in it.
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 (20 o C and 1 atmosphere pressure) is 24 dm 3.
CHEMICAL CHANGES Topic 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! Alsodisplacement reactions
Redox 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 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 Fe 3+ + 6e - 2 Fe 3 O 2-3 O + 6e -
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
Required Practical See also Separation Techniques Topic 1 Making Salts
Required Practical See also Quantitative Chemistry 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 Titration
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 = Can also be a solution (dissolved solute in solvent) Copper sulfate: Cu 2+ + 2e - Cu SO 2-4 SO 4 + 2e - The electrons leave the sulfate and attach to the copper. Lead bromide Together; Pb 2+ + 2e - Pb 2Br - Br 2 + 2e - Pb 2+ (aq) + 2Br - (aq) Pb (s) + Br 2 (aq) Cryolite used to lower mp of aluminium oxide Aluminium oxide 4Al 3+ + 12e - 4Al 6O 2-3O 2 + 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) Cl 2 (g) + 2e - What is the overall equation? 2Na + (aq) + 2Cl - (aq) + 2H + (aq) + 2OH - (aq) 2NaOH + Cl 2 (g) + H 2(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.
Required Practical 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 Solution Positive electrode (anode) Copper (II) chloride chlorine copper Copper (II) sulfate oxygen copper Sodium chloride chlorine hydrogen Sodium sulfate oxygen hydrogen Negative electrode (cathode)
ENERGY CHANGES Topic 5
Energy Change 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. Exothermic Transfers energy from chemicals to surroundings, T surroundings increase. Making new bonds is exothermic. E.g. combustion, many oxidation reactions, neutralisation, selfheating cans, hand warmers.
Required Practical Energy Changes IV: volume sodium hydroxide (cm 3 ) DV: Maximum temperature ( C) CV: volume hydrochloric acid (cm 3 ), temperature of surroundings ( C) a Stand the polystyrene cup in the beaker. b Use the measuring cylinder to measure out 5 cm 3 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 cm 3 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. Insulated cup & lid- could be a gap Digital 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 Fuel (hydrogen) enters cell on one side, becomes oxidised (reacts with oxygen)- sets up potential difference. hydrogen + oxygen water 2H 2(g) + O 2(g) 2H 2 O (l)
Fuel Cells Advantages 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).