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ATOMIC STRUCTURE
Atomic model changes over time – due to discovery of subatomic particles
John Dalton 1803: solid spheres makes up different elements
Thomson 1897: electrons
Plum Pudding Model: ball of positive charge with electrons in it
Rutherford 1909: gold foil experiment – tiny positive nucleus with electrons around
Bohr 1913: fixed electron shells
Nucleus
protons & neutrons
electrons orbit in shells – energy levels
Very small compared to overall size of atom
Holds most of atom’s mass
Proton
1 amu
Positive

Neutron
1 amu
Neutral

Electron
1/1840
Negative

mass number = mass of nucleus = protons + neutrons
Atomic number = number of protons = number of electrons – atom has no charge
Number of protons unique to element
Isotopes: same element / different number of neutrons
Decimal: if element has 1+ isotope – average of mass numbers & taking into account how much there
is of each one
Calculate: relative atomic mass from relative masses and abundances of element’s isotopes
Relative abundance: proportion of one isotope in an element
Relative atomic mass Ar: average mass of all atoms of that element – usually most common isotope
Ar = (mass of isotope x % of isotope) + (mass of isotope x % of isotope)
100

PERIODIC TABLE
Atomic number of elements increases through periodic table
Rows/periods: atomic number
Columns/groups: similar properties
Metals/non-metals divided in table – explain division in terms of atomic structures of elements
Mendeleev
arranged elements in periodic table:
In order of atomic mass – pattern appeared
Put elements with similar chemical properties into groups
used table to predict existence/properties of undiscovered elements
Left gaps to keep elements with similar properties together
Used properties of other elements in column to predict properties of undiscovered elements
He sometimes swapped positions of elements he thought fitted better with their properties
Atomic mass he had was wrong – other isotopes were more abundant

BONDING
non-metals
low boiling point
Don’t conduct electricity

metals
Shiny solids
High melting points
High density
Conduct electricity

Limitations of models
dot and cross: presents electrons from both elements as if different & no size/arrangement
ball & stick models: don’t show correct scale for atoms/ions & appear to have big gaps between atoms
2D: no shape/size
3D: only show outer layer of substance
Ionic Bonding
‘–ide’ ending: negative ions of only one element (except hydroxide)
Oxide: O2Hydroxide: OHHalide:
‘–ate’ ending: negative ions of oxygen + another element
Nitrate: NO3Carbonate: CO32Sulphate: SO42ion: atom with charge – lost/gained electron
cation: + electron lost
anions: – electron gained
Positive metal + negative non-metal
Metal ion: group number 1/2 = electrons lost = positive charge
Non-metal ion: 8 – group number 6/7 = electrons gained = negative charge
Ionic equations
Magnesium + copper sulphate → magnesium sulphate + copper
Mg(s) + Cu2+(aq) + SO42-(aq) → Mg2+(aq) + SO42-(aq) + Cu(s)
Remove ‘Sulphate’ – spectator ions, doesn’t change in reaction
Mg(s) + Cu2+(aq) → Mg2+(aq) + Cu(s)
Oxidation: Mg → Mg2+ – magnesium lost electrons, positive
Reduction: Cu2+ → Cu – copper gained electrons, negative

redox reaction

Giant ionic lattice
High melting point: strong electrostatic forces of attraction between oppositely charged ions in all directions
Don’t conduct as solid: charged ions in fixed position
Conduct when melted/dissolved in water: charged ions free to move
Some soluble
Covalent bonding
Non-metal molecule
Electrostatic forces of attraction between 2 nuclei & shared pair of electrons – very strong
Number of covalent bonds formed = number of electrons needed for full outer-shell
Molecule size increase = intermolecular force increase – melting/boiling point increase

Simple molecular
Low boiling point: weak intermolecular forces (between covalent molecules) – covalent bonds don’t need to
be broken
Don’t conduct: not charged – no free electrons
Some soluble
Giant covalent
All atoms bonded
High boiling point: lots of strong covalent bonds
Graphite conducts electricity: 1 delocalised electron between layers
Not soluble
Graphite
Layers: strong bonds within / weak bonds between
Each atom: 3 covalent bonds – 1 electron not
shared & moves between layers
Uses
Lubricant: slip over eachother
– weak bonds between layers
Electrode: conducts electricity

Diamond
Each atom: 4 covalent bonds – no delocalised
electrons
Strong bonds held in rigid lattice structure
Hardest natural-occurring substance
Use: Cutting tools – very hard

Fullerenes – allotrope of carbon
Made of carbon atoms shaped as hexagons/pentagons/heptagons
Used to cage other molecules: deliver drug directly to cells in body
Huge surface area – help make industrial catalyst
C60 – Buckminsterfullerene
Forms hollow sphere
20 hexagons / 12 pentagons
Stable molecule
Forms soft brown-black crystals
Graphene: Sheets of graphite hexagons
Conducts electricity
High tensile strength: doesn’t break when stretched – used to strengthen material without adding weight
Simple polymers
covalently bonded carbon chains
poly(ethene)
Metallic Bonding
Electrostatic forces of attraction between metal ions & delocalised electrons
Alloys: stronger – pattern altered – layers don’t slide over each-other so easily – ions are different sizes
Group 1: weaker electrostatic attraction – only one electron on outer shell = less delocalised
Group 2: stronger electrostatic attraction – more positively charged than Group 1 cations
Giant Structures: metal ions regularly arranged
High boiling point: lots of strong electrostatic forces of attraction
Conduct:
electricity: delocalised electrons move through structure
Heat: ions transfer energy between them
Malleable/Ductile: layers can slide over each-other
Not soluble
Dense: ions packed closely together

CALCULATIONS involving mass
Relative formula mass Mr = sum of relative atomic masses Ar
Empirical formula: simplest ratio of atoms in compound
Percentage or mass
Mr
Find ratio
Magnesium oxide experiment to determine empirical formula
Weigh crucible + lid
Add 5cm clean magnesium ribbon
Weigh crucible + lid + magnesium
Place crucible on pipe clay triangle
Heat crucible strongly
Lid on: prevent solid escaping
Lid slightly off: allow oxygen to enter
Remove lid: when all magnesium burnt
Crush solid to grey powder
Weigh crucible + lid + magnesium oxide
Add masses to equation and calculate
Law of conservation of mass:
closed system: precipitation reaction in closed flask
non-enclosed system: reaction in open flask that takes in / gives out gas
one mole =

Avogadro number 6

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