Elements of life

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  • Created on: 20-04-21 18:00

Elements of life

Atomic model

  • John Dalton = solid spheres
  • JJ Thompson = plum pudding model
  • Rutherford = nuclear model
  • Henry Mosely = nuclear charge increases by 1
  • James Chadwick = neutron
  • Bohr = Bohr model

Isotope - atoms of the same element with the same number of proteons but a different number of neutrons 

Releative atomic mass - average of relative isotopic masses taking into account abundances, releative to 1/12 of a carbon 12 atom

Nuclear fusion 

  •  two light atomic nuclei fuse to form a new single heavier nucleus of a new element. 
  • only at a high temperature and pressure
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Mass spectroscopy

  • measures atomic mass and abundance of isotopes
  • separates on mass:charge ratio

Absorption spectrum - visible light spectrum with black lines

Emission spectrum - black spectrum with coloured lines

Wave theory

  • light behaves like a wave
  • speed of light = wavelength x frequency 

Particle theory

  • light is a beam of photons
  • energy = plank constant x frequency
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Bohr's theory

  • electrons in indefinite shells
  • electrons move up shells by absorbing frequencies, emit frequencies when they move back down
  • energy levels closer at higher frequencies
  • energy of photon equal to difference between two energy levels

Shell sizes

  • n=1 - 2 electrons
  • n=2 -  8 electrons 
  • n=3 - 18 electrons 
  • n=4 - 32 electrons

Subshells

  • s subshell = 2 electrons
  • p subshell = 6 electrons
  • d subshell = 10 electrons
  • f subshell = 14 electrons
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Orbitals 

  • holds two electrons each with opposite spin
  • s orbital = spherical
  • p orbital = dumbell shaped
  • half shell stability - electron promoted so all half shells are full

Periodic table

  • Mendeleev - atomic mass order, grouped with same properties
  • today - atomic number order

Melting points periodicity

  • Groups 1-4 increasing, due to giant metallic structure
  • Groups 4-5 sharp increase, due to giant covalent structure
  • Groups 5-0 low, due to simple covalent structure
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Water of crystallisation

  • water inside a crystal lattice 
  • pentahydrate = 5 moles of water

Ideal gas equation

  • molar volume - 1 mole of any gas at RTP has a volume of 24 dm-3

Yields

  • percentage yield = actual yield/theoretical yield
  • never 100% due to unwanted side reactions, impurities and equilibrium system

Covalent bond - shared pair of electrons attracted to both nuclei

Dative covalent bond - shared pair of electrons provided by 1 atom only, normally occurs with ions

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Shapes of molecules

  • Valence shell repulsion theory - electron pairs repel each other as far apart as possible
  • lone pairs have a stronger repulsion so reduce angle by 2.5 degrees

Molecule shapes

- linear = 2 areas of electron density,180 degrees

- trigonal planar = 3 areas of electron density , 120 degrees

- tetrahedral = 4 areas of electron density, 109.5 degrees

- pyramidal = 4 areas of electron density, 1 lone pair, 107 degrees 

- non-linear = 4 areas of electron density, 2 lone pairs, 104.5 degrees 

- trigonal bipyramidal = 5 areas of electron density, 120 degrees and 90 degrees

- octahedral = 6 areas of electron density, 90 degrees 

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Ionic structure

  • loss/gain of electrons creating an electrostatic atraction between positive and negative ions
  • giant regular lattice 
  • high melting point/boiling point
  • only conducts electricity as molten or aqueous

Metallic structure

  • metal cations regularly arranged in a sea of delocalised electrons
  • giant metallic lattice
  • high melting point/boiling point
  • conduct electricity 

Simple covalent

  • small groups of atoms covalently bonded with weak intermolecular forces
  • low melting point/boiling point
  • don't conduct electricty 
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Giant covalent structure

  • high melting points/boiling points
  • good thermal conductors
  • can't conduct electricity 
  • insoluble in polar substances (no ions)

Polymers

  • long chain molecules
  • low melting points/boiling points
  • don't conduct electricity 

Group 8 - monatomic, weak forces of attraction

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Ionisation enthalpyenergy required to remove one electron from every atom in one mole of a gaseous element

Factors affecting enthalpy 

  • nucleus charge - more protons = stronger attraction
  • disrance - smaller distance = weaker attraction
  • electron shielding - electrons between electron and nucleus so nuclear attraction is weaker
  • half shell stability - electrons not repelled by each toher so more energy required

Standard solutions

  • C1V1 = C2V2
  • stock solution = original standard solution

Acid and bases

  • amphoteric - acts as an acid and base
  • Bronsted - lowery theory - acid proton donor, base proton acceptor.
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Group 2 theory

  • alkaline earth metals 
  • reactivity increases down the grop
  • thermal decomposition - MCO3 --> MO + CO2
  • hydroxide solubility increases down the group
  • carbonate solubility decreases down group

Carbonates of group 2

  • thermal stability increases down group
  • metal ion polarises negative cloud around carbonate making it less stable so it thermally decomposes
  • charge density (meaaure of concentration of a charge on an ion) decreases down group 2 
  • higher charge density = higher thermal decomposition

Oxonium ion

  • hydrogen from dissolved acid attracted to negative region of water molecule forming dative covalent bond
  • proton can then be donated
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Solubility rules

  • group 1 and ammonia soluble
  • all nitrates soluble 
  • Halides soluble except lead and silver
  • sulfates soluble except lead, silver, barium and calcium)
  • hydroxides and carbonates insoluble except group 1 or ammonia 

Precipitation reaction

  • two soluble salts mixed and then filtered to makes insoluble salt

Soluble salt production with metal/insoluble base 

  • add metal/base in excess to acid. Filter and carry out crystallisation

Soluble salt production 

  • Titration with indicator
  • calculate average of concordant results (between 0.1 of each other)
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Sodium hydroxide test

  • silver = brown precipitate 
  • calcium = white precipitate
  • copper = blue precipitate 
  • lead = white precipitate 
  • Iron II = green precipitate 
  • Iron III = red/brown precipitate 
  • Zinc/aluminium = white and redissolved in excess NaOH

Carbonate test

  • add HCL
  • bubblr through limewater, cloudy means carbon dioxide present

Sulphate test

  • add dilute HCL to remove carbonate
  • Add barium chloride - white precipitate means sulphate present
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Test for hydroxide

  • Dip red litmus paper into solution. Turns blue.

Test for Halides

  • add dilute nitric acid
  • add silver nitrate 
  • chlorine = white precipitate 
  • bromine = cream precipitate 
  • iodine = yellow precipitate

Adding dilute ammonia to halide test

  • silver chloride redissolves
  • silver bromide redissolves in concentrated ammonia
  • silver iodide won't redissolve
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