gcse science chemistry unit C4 cards

the theory of the atomic structure of the atom has changed throughout history:

• 1800s- John Dalton- solid spheres+diff. spheres make up diff. elelments 
• 1897- J J thompson- measurements of charge+mass show must be made of smaller particles- electrons- instead of a solid sphere it was now a plum pudding 
• 1909- Rutherford- gold foil experiment- firing postivive particles at sheet of gold- expeected most to be defelected from positive plum pudding- but passed straight through- concluded a tiny postive nucleus surrounded by electrons with empty space 
• Bohr model- electrons contained in shells- fixed orbits with shells of fixed energy- supported by many experiments 

nucleus- protons+neutrons, almost whole mass concentrated in nucleus 

electrons- move around in electron shells, negatively charged, virtually no mass 

mass no- total no. of protons+neutrons 

atomic no.- no. of protons 

• elements in order of ascending atomic number 
• groups corresponds to number of electrons in outer shell 
• rows called periods- each new period represents another full shell of electrons 
• period it relates to corresponds to amount of shells 

isotopes- different forms of same element- same no. protons diff. no. neutrons 

history: 

• 1800's only thing that could be measured was atomic mass- measured in order of mass
• 1828- Dobereiner- groups based on chemical properties- placed in triads (threes) middle eleemnt had average mass of three 
• 1864- Newland-every eigth element had similar properties- listed in rows of seven- sets were called newlands octaves- but broke down on thrid row- left no gaps - critised- groups contained elements that didnt have same property- mixed up metals+non- didnt leave gaps for others 
• 1869- Mendeleev- arranged 50 known elements in order of atmoic mass but left gaps for undiscovered where properties didnt match- vertical groups had similar properties -more sense when later discoveries of atomic structure- atomic no. one more than previous- matches way electrons arranged in atom's shells 

• shells sometimes known as energy levels- lowest filled first 
• shells go- 2,8,8...
• use electron configuration to work out, period, group, atomic no. 

period= same as no. of shells 

group no.= how many electons occupy outer shell 

atomic no.= adding up electrons 

• atoms lose/gain electrons forming charged particles(ions) stongly attracted to eacother 
• atoms with 1 outer electron react stongly with atoms with 7- create full outer shell 

ionic compounds form giant ionic lattices:

• ions closely packed, do not conduct when solid, very stong chemical bonds 
• MgO+NaCl: high MP+BP due to stong attraction to non-metal, MgO higher MP due to double charge+O2 smaller (compact), melt+conduct electricity (ions separate+free to move)

ions= charge particels of single atoms or group of atoms 

metals= ions when lose electron       non-metals= ions when gain electron 

working out the formula you have to balance negative+postive charges 

all atoms end up with full outer shell of electrons- as a result of ionic bonding 

• sharing electrons between non-metals to make full outer shell 
• each covalent bond provides one extra electron

e.g. hydrogen, chlorine gas, methane, water, carbon dioxide

• substances formed are simple molecular structures 
• atoms held together by strong covalent bonds- forces of attraction between molecules is weak
• feeble intermolecular forces- low MP+BP- dont conduct electricity as no free electrons/ions

group 1- alkali metals 

• go down+become more reactive- outer electron easily lost- less energy further away from nucleus 
• low MP+BP, low density, very soft- always form ionic compunds 
• forming stable structure- lose electron- OXIDATION 
• react with water produce H2- inc. in reactivity as go down= alkali hydroxide forms 
• burn with characteristic colours- dip wire into hydrochloric acid- put powdered metal on- burn in blue bunsen flame- LITHIUM=RED    SODIUM=YELLOW/ORANGE    POTASSIUM=LILAC 

group 7- halogens 

• go down+become less reactive- less incline to gain extra electron- further away from nucleus 
• go down- MP+BP inc.
• forming stable structure- gain electron- REDUCTION 
• react with alkali metals=salts (metal halides)
• more reactive displace less reactive 
• CL2=fairly reactive,poisonous,dense green gas(low BP)BR2=dense,poisonous,orange liquid I2= dark grey crystalline solid (high BP)

crystal structures 

• same basic properites- metallic bonding- outer e move freely- sea of free e giving properties 
• most hard, dense,shiny, strong attraction between free e+positive ions- strong metallic bond 
• high MP+BP- dec. as atomic radium inc. 
• high tensile strength yet malleable 
• good conductors of heat+electricity- free e carries current+heat energy 

metal properties 

• stainless steel--conductor, doesnt rust--saucepan
• copper--best conductor, easily bent---electrical wiring 
• aluminium---low density,strong,doesnt corrode---aeroplanes 
• steel---brittle(iron+carbon)---bridges 

superconductors- metals at very low temps

• resistance disappears completely- no electrical energy turned into heat+wasted 
• start current via battery-take battery out-current carry on flowing forever 

can create

• loss-free power transmission cables, strong electromagnets, faster electronic circuits 

problems 

• expensive to make so cold(-265d.c), scientists are getting there- metal oxide at (-135d.c) want to create one working at room temperature 

metals in middle of the table- transistion elements 

• typical metallic properties- make good catalysts(iron-haber process/nickel-hydrogenation of alkenes)
• colourful compounds due to transition metal contained-
• IRON(II)=LIGHT GREEN          IRON(III)=ORANGE/BROWN          COPPER=BLUE 

thermal decompostion- breaking down with heat into other substances  

• transition metal carbonates break down when heated into metal oxide+carbon dioxide result in colour change

precipitation- solid forms in solution 

• two solutions react forming insoluble solide- said to precipitate out+called precipitate 
• soluble transition metals+sodium hydroxide=insoluble hydorxide 
• can write in form of ionic equations 

use precipitation to test for transition metal ions  

• insoluble transition metals have distinct colours- e.g iron(III) hydorxide forms (orange/brown) so you know you have Fe3+ ions in solution 

2 sources in UK 

• surface water: lakes, rivers, reservoirs 
• ground water: aquifer 
• limited as depenent on annual rainfall- have to conserve water 

purified in water treatment plants- groundwater fairly pure but surface water:

• 1) filtration: wire mesh screens out large twigs 
• 2) sedimentation: iron/aluminium sulphate added making fine particles clump together+settle at bottom 
• 3) chlorination: gas bubble through killing harmful bacteria+microbes 

tap water impurities 

• nitrate residues: from excess fertiliser run off- prevent blood oxygen carrying capacity 
• lead compounds: from old pipes- poisonous particularly for children 
• pesticide residues: spraying too close to lakes/rivers 

fresh water created by distilling it (sea) in dry countries: needs loads of energy so expensive+not practical for producing large amounts of fresh water 

you can test water purity for various dissolved ions- precipitate reactions:

test for sulphate ions using barium chloride:

• add dilute hydrochloric acid- then10 drops of barium chrloride:
• white precip. (sulphate ions present)

test for halide ions using silver nitrate 

• add dilute nitric acid- then 10 drops of silver nitrate 
• halide ions present:
• white precip. (chloride ions)
• cream precip. (bromide ions)
• iodide ions. (pale yellow precip)

• silver nitrate+sodium chloride=silver chloride+sodium nitrate 
• silver nitrate+sodium bromide=silver bromide+sodium nitrate 
• silver nitrate+sodium iodide=silver iodide+sodium nitrate 

you can test water purity for various dissolved ions- precipitate reactions:

test for sulphate ions using barium chloride:

• add dilute hydrochloric acid- then10 drops of barium chrloride:
• white precip. (sulphate ions present)

test for halide ions using silver nitrate 

• add dilute nitric acid- then 10 drops of silver nitrate 
• halide ions present:
• white precip. (chloride ions)
• cream precip. (bromide ions)
• iodide ions. (pale yellow precip)

• silver nitrate+sodium chloride=silver chloride+sodium nitrate 
• silver nitrate+sodium bromide=silver bromide+sodium nitrate 
• silver nitrate+sodium iodide=silver iodide+sodium nitrate