National Chemistry

 
This page contains information which is essential to success in the exam.

1.   Define and recognise examples of :-
isotopes        isomers       catalyst     electrolysis neutralisation        precipitation     redox        unsaturated
hydrolysis        condensation      addition homologous series       

2. Learn about the following
Bonding in:- (a) metals (b) ionic compounds (c) covalent substances
Conductivity of:- (a) metals ionic (b) compounds (c) covalent substances

low melting point = covalent molecule

 tetrahedral shape eg CCl4  SiH4 CH4  (beware NH3  )

3. Electronic configuration of atom and ion.
What caused the change ?

4. Write ionic formula zinc nitrate iron(II) oxide calcium hydroxide
Write ionic formulae for aluminium sulphide, potassium oxide, magnesium chloride and iron(III) bromide
See how to write formula using the complex (polyatomic) ions found on page 4

5. Copy out and/or recognise oxidation and reduction - redox - use page 7 of the data book
for reactions, cells and electrolysis

6. Calculate given two figures the third using

  mole = mass in question   concentration = mole
  mass of 1 mole litres

7. Non-metal oxide = acidic = Hydrogen ion = low pH (high acidity low pH !!) ; turns indicator red
  metal oxide = alkali = hydroxide ion = high pH (NB alkali metals); turns indicator blue
  pH 7 is neutral where hydrogen ions = hydroxide ions

8.   Acid + metal oxide ----> water + salt
  Acid + metal hydroxide ----> water + salt
  Acid + metal carbonate ----> water + salt + carbon dioxide
  Acid +(reactive) metal ----> hydrogen + salt (also redox)
pH moves to 7 (increases); H+ decreases = neutralisation

9. Electrochemical series - always happens to most reactive metal
  most reactive oxidises faster
  most reactive loses electrons to less active;
  most reactive displaces less active
 
  in a cell electrons flow from more reactive to less reactive
  further aprt the metals in the electrochemical series (page 7) the greater the voltage

  react fast break down slow
  the morer reactive a metal the harder it is to extract from a compound with the most reactive metals needing electrolysis to obtain the metal
  A mid reactive metal such as iron is obtained by heating the oxide with carbon (a reducing agent - steals the oxygen) in a blast furnace

10. Draw alkane and alkene with two to six carbons
bromine water decolourises by alkenes
fractional distillation based on melting points
properties- viscosity, Boil pt, ignition temp., density, increase with size/mass
cracking produces smaller molecules one of which is an alkene

11. Corrosion requires water and oxygen
Iron (Fe) rusts to iron (II) ions (Fe2+) which turn ferroxyl indicator blue
appearance of the blue colour means oxidation (loss of electrons) of iron>
Iron (II) ions (Fe2+) oxidise further to iron (III) ions (Fe3+)

12. More reactive metal (eg Mg, Zn) sacrificially protects iron
Coating with zinc called galvanising
Iron corrodes to protect less reactive metal (eg Sn)

13. Calculate empirical formula
Calculate percentage composition

14. Alkene monomers react by addition (breaking one of the bonds in the C=C)
to polymerise (watch)
  monomer polymer repeating unit

15. Plastics
advantages - light, durable, insulator,
disadvantages - do not rot (durable), burn to give poisonous fumes
Themosetting- not softened by heat;
thermoplastic</STRONG> softened by heat

16. Fertilisers required to produce sufficient food.
Natural Bacteria methods cheaper
nitrogen is an unreactive gas - needs lot of energy to react with oxygen
Haber (Iron catalyst) process reacts nitrogen and hydrogen to make ammonia - high pressure (expensive) moderately high temperature
moderately high temperature = low yield of ammonia but faster reaction
Oswald (platinum catalyst) process ammonia to produce nitrogen dioxide to make nitric acid - high temperature to start with then as exothermic supplies own heat
HI=Haber & Iron; OP=Oswald & platinium

17. Substances with Carbon burn to give CO2
Substances with Hydrogen burn to give H2O
Substances - Hydrocarbons and carbohydrates - that burn to give CO2 and H2O must contain H and C (but could contain other elements).
>Carbohydrates produced by photosynthesis also produces oxygen
respiration in animals use carbohydrates

18. Carbohydrates
  Fructose C6H12O6 monosaccharide turns hot blue Benedict's brick red
  Glucose C6H12O6 monosaccharide turns hot blue Benedict's brick red

  Maltose C12H22O11 disaccharide turns hot blue Benedict's brick red
  Sucrose C12H22O11 disaccharide no test
 
  Starch (C6H10O5)n poloysaccharidde turns iodine blue-black

Isomers have same molecular formula but different structural formula

19. Glucose molecules joining together to form starch and water is condensation polymerisation
Starch or sucrose molecules being broken down by water is hydrolysis reaction.
Small glucose molecules can pass through the gut wall.
Hydrolysis occurs in body using biological catalysts called enzymes
Enzymes used in fermentation to produce (ethanol) a member of the alcohol homologous series.

20. Gases
Noble gases (He, Ne, Ar etc) do not react - monatomic - full outer orbital
Diatomic gases - elements O2, H2, Cl2, F2, N2
  - compounds CO, HCl
Air - 80% nitrogen and 20% oxygen
Metal oxides (eg alkali metals) produce alkalis and
ammonia NH3 is only alkaline gas
Non-metal oxides C02, N02, SO2 produce acids

CO2 produced by heating carbonate or reacting an acid + carbonate;
turns lime water milky; used in photosynthesis

H2 released when metals (not copper or unreactive metal) react with acid
or when very reactive metals react with water;
produced when acid electrolysed;
 burns with a pop

O2 relights glowing splint; produced by photosynthesis; used in combustion

CH4 natural/North Sea gas; burns to give C02 and H2O; has tetrahedral shape;

NH3 alkaline gas; produced when sodium hydroxide reacts with ammonium compound

Making gases
acid + metal (not Cu. Ag etc) ---> H2
acid + carbonate ---> CO2
alkali + ammonium ---> NH3

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