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the atom

  • contains protons +1, neutrons 0 and electrons -1 move in orbitals around nucleus
  • top number - protons plus neutrons, bottom number - just protons
  • ions have different numbers of electrons
  • isotopes have different numbers of neutrons
  • relative atomic mass - average mass compared to carbon 12
  • molar mass - mass of one mole
  • number of moles = mass / relative formula mass
  • number of moles = concentration x volume
  • number of moles = volume (dm) / 24
  • ppm = mass of solute / mass of soluble x 1,000,000
  • empirical - smallest whole number ratio
  • molecular - actual number in molecule
  • ionic equation - only ions no spectators with state symbols
  • aq - aqueous, s - solid, g - gas, l - liquid
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  • hydrated - have water of cyrstalisation
  • MASH metal + acid = salt + hydrogen
  • OAWS oxide + acid = water + salt
  • HAWS hydroxide + acid = water + salt
  • CAWCS carbonate + acid = water + CO2 + salt
  • percentage yield = actual yield / theoretical yield x100
  • atom economy = molecular mass of desired / molecular masses of all x100
  • addition reactions are 100% atom economy, substitution less
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mass spectrometry

  • using a mass spectrometer to measure relative masses
  • vaporisation - gas
  • ionisation - electrons knocked off 
  • acceleration - electric field
  • deflection - magnetic field
  • detection - graph
  • relative atomic mass on bottom, % abundance on side
  • average atomic mass of all isotopes = % abundance x relative atomic mass (plus all other isotopes) / 100
  • mass of molecules shown by peak furthest right
  • used for drugs testing, carbon dating, testing in sport
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electron structure

  • s orbital is circular
  • p orbital is dumbbell
  • 1s1 2s2 2p6 3s2 3p6 4s2 3d10 
  • fill singularly before sharing in p and d orbitals
  • first ionisation energy is the energy required to remove one electron from each atom in a mole of gaseous atoms to form one mole of gaseous 1+ ions
  • nuclear charge is increased by increasing number of protons
  • distance from nucleus increases as you fill shells, those in closest shells held tightest
  • shielding is the afffect the electrons have to reduce the effect of nuclear charge
  • first ionisation energies across period increase
  • first ionisation energies down group decrease
  • succesive ionisation energies in the same atom increase - increasing nuclear charge, decreasing shielding and distance, less repulsion
  • big rise in energy mean new shell has been entered
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periodic properties

  • group number is the number of electrons in elements outer shell
  • group one and two are s block
  • group two centre are d block
  • group three four five six seven and zero are p block
  • atomic radius decreases across period
  • ionisation energy increases across period
  • drops between group two and three (Be to B) because the extra electron is in next subshell so the extra shielding outweighs the nuclear charge increase
  • drop between group five and six (N to O) because nitrogen has filled all its orbitals singularly meaning the extra electron O has must share and the repulsion between electrons make removing one easier
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ionic bonding

  • ions joined by electrostatic attraction - holds anions and cations together
  • electrons transferred from one atom to another
  • cations are positive and have lost electrons, anions are negative and have gained electrons
  • elements in the same group form ions with the same charge
  • soluble in water, not in non-polar solvents
  • conduct when molten or dissolved - ions fixed when solid, free to move and carry current as liquid or in solution
  • positive ions smaller than negative as the effect of the nuclear charge is stronger and less repulsion
  • isoelectronic means different ions of atoms with the same number of electrons
  • green copper chromate turns damp filter paper blue at cathode yellow at anode, positive copper ions move to cathode, negative chromate ions move to anode under current
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covalent bonding

  • two or more atoms share electrons so they have full outer shells
  • sigma bonds in s orbitals overlap in straight line giving highest electron density
  • pi bonds formed when two p orbitals cross are weaker than sigma, so more reacitve
  • dative covalent - both electrons given by the same atom
  • form giant structures
  • insoluble in polar solvents
  • form hard cystals with high melting points
  • cant conduct
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metallic bonding

  • metals form giant lattice structures with many metallic bonds
  • outermost shell of electrons delocalised so can move around leaving positively charged ions which attract the electrons holding everything together
  • more delocalised electrons means higher boiling point as bonds held stronger
  • metals are malleable, can be shaped as the layers can slide over each other
  • good thermal and electrical conductors as electrons free to move and pass on current or kinetic energy
  • insoluble except in liquid metals due to strength of bonds
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organic compounds

  • general formula described any member of a family of compounds
  • structural shows structure carbon by carbon
  • displayed shows lines and bonds
  • skeletal shows bonds of carbon represented as lines
  • 1 meth, 2 eth, 2 prop, 4 but, 5 pent, 6 hex, 7 hept, 8 oct, 9 non, 10 dec
  • find and count longest chain of carbons
  • alkanes - ane
  • alkenes - ene
  • halogenalkanes - chloro bromo iodo
  • alcohols - ol
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  • general formula CnH2n+2
  • hydrocarbons as they only contain hydrogen and carbon
  • saturated, cant form any more bonds, four bonds per carbon
  • cycloalkanes CnH2n
  • structural isomers are different arrangements of the same atoms
  • chain isomers have the same structural formula but are branched in different ways
  • burn completely in oxygen making carbon dioxide and water
  • free radical substitution is where a hydrogen is swapped with a chlorine or bromine, the reaction is started by UV light and has three stages
  • initiation, sunlight used to break chlorine bonds to make two free radicals (atom with an unpaired electron) which are highly reactive
  • propogation, free radicals are used in a chain reaction, 1) the chlorine free radical attacks methane to make HCl and CH3'. 2) the CH3' attacks a Cl2 molecule making CH3Cl and Cl'. 3) repeats until all Cl2 and CH4 gone
  • termination, there are three possible reactions: Cl' + Cl' to make Cl2, CH3' + CH3' to make C2H6, Cl' + CH3' to make CH3Cl
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  • crude oil is a mixture of different length hydrocarbons
  • can be separated by fractional distilation, heated until vaporised at 350 degrees, put into fractioning column, largest condense immediately, collecting according to boiling point
  • smallest fractions most useful, can be cracked to make shorter chains from long
  • some alkanes in petrol are straight chain, others are branched which burn better
  • can be reformed into cycloalkanes and branched alkanes using a catalyst
  • useful for fuel but harm the enviroment as they release CO2
  • sulfur oxides are poisonous, cause breathing problems and can be converted to sulfuric acid which causes acid rain leading to habitat and building damage
  • carbon monoxide is released during incomplete combustion and can cause health problems
  • nitrogen oxides add to the problem of acid rain and turns to ozone in sunlight
  • greenhouse gases absorb and re-emit infrared radiation to earth which keeps the planet warm however there is too much so the planet is becoming too warm which could lead to melting of ice caps and sea levels rising
  • using fossil fuels is unsustainable as they will run out and harm the environment
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  • general formula is CnH2n
  • have at least one carbon double bond so unsaturated
  • more reactive, has a sigma and pi bond due to double bond so higher electron density
  • pi bond is above and below rest of molecule so prone to attack
  • alkenes react with hydrogen to form alkanes
  • electrophillic addition with halogens: 1) double bond repels electrons in Br2 which polarises it. 2) heterolytic fission takes place - this is unequal so the Br closest to the double bond gives up bonding electrons to become positive and joins a carbon. 3) carbocation means a organic ion containing a positively charged carbon atom, the double bond breaks and the other Br joins the carbon to form 1,2-dibromoethane
  • bromine water used to test for double bonds, will decolourise in their presence
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E/Z isomerism

  • carbon atoms and the atoms they're bonded to all lie on the same plane and cant rotate
  • the double bond causes the rigid structure and E/Z isomerism
  • stereoisomerism means alkenes have the same structural formula but different arrangements in space
  • alkenes have this due to their inability to rotate
  • occur when two double bonded carbon atoms have different groups attached
  • when the same groups are on the same side (both above/ below) the double bond then its z or cis
  • when the same groups are on opposite side of the double bonds its E or trans
  • cis - trans is only for two different groups, if more then E Z must be used and one group takes priority
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  • alkenes can open up double bonds and join to form addition polymers
  • small individual alkenes are monomers
  • polyalkenes are unreactive as they have lost their double bond
  • the monomer is found by working out what the repeat unit is and putting in a double bond
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shapes of molecules

  • all electrons are negatively charged therefore repel each other moving to the point of minimum repulsion
  • lone pair lone pair repulsion > lone pair bond pair repulsion > bonding pair bonding pair repulsion
  • linear - 2 bonding = 180
  • trigonal plannar - 3 bonding = 120
  • tetrahedral - 4 bonding = 109.5
  • trigonal bipyramidal - 5 bonding = 120 and 90
  • octrahedral - 6 bonding = 90
  • pyramidal - 3 bonding 1 lone = 107
  • angular - 2 bonding 2 lone = 104.5
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carbon structures

  • allotypes are different forms of the same element in the same state
  • carbon forms three allotypes
  • diamond is where each carbon is covalently bonded to four other carbon atoms arranged in a tetrahedral shape in a crystal lattice
  • has a very high melting point - sublimes at 3800K and is very hard, cant conduct electricity and doesnt dissolve in solvents but is a good thermal conductor as it can pass kinetic energy
  • graphite is the second allotype, the carbon atoms are arrange in flat hexagons covalently bonded to three carbons and the third atom is delocalised
  • there are weak bonds between layers which are easier broken making it a good lubricant as layers can slide over each other
  • delocalised electrons are free to move so its a good conductor, and makes lightweight sports equipment that is strong as the layers are spaced quite far apart, its insoluble and the layers themselves are very hard to break
  • fullerenes are the final allotype, they are hollow carbon tubes or balls, each carbon forming three covalent bonds with free electrons so it can conduct, they're nanoparticles, buckminsterfullerene C60 was the first to be discovered - naturally occurs in soot, soluble in organic solvents, form bright coloured solutions
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  • graduation from ionic to covalent
  • electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond
  • fluorine is most electronegative, oxygen nitrogen and chlorine are also fairly strongly
  • covalent bonds can become polar covelent due to differences in electronegativity
  • bonding electrons are pulled towards the more electronegative element
  • elements with similar electronegativities will be non polar, in polar bonds a difference induces a dipole
  • polar bonds dont always mean polar particles
  • if the polar bonds all pull in the same direction the molecule will be polar, if they are in opposite directions the polarity will cancel out
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intermolecular forces

  • london forces: cause all atoms and electrons to be attracted to each other, charge clouds always moving and inducing temporary dipoles, which induces another dipole in the opposite direction on a neighboring atom which causes a domino affect. Larger molecules means more electrons so stronger london forces so higher boiling points
  • permanent dipole - dipole force: polar molecules with permanent dipoles lead to weak electrostatic forces of attraction between molecules. Putting a charged rod next to a jet of liquid will move it towards the rod if its polar, Chlorine, nitrogen, oxygen and fluorine make permanent dipoles
  • hydrogen bonding: hydrogen is covalently bonded to fluorine, oxygen and nitrogen, all very electronegative so draw electrons away from hydrogen so its polarised and has a high charge density so can form weak bonds with lone pairs on nitrogen fluorine or oxygen, hydrogen bonds have a big effect on melting and boiling points
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  • for a substance to dissolve the bonds in the substance and the solvent must break and those formed must be stronger or the same as those broken
  • polar solvents such as water where the bonds are hydrogen
  • non-polar solvents such as hexane where there are london forces
  • ionic substances dissolve in polar solvents as the ions are attracted to the oppositely charged water parts pulling them away from the lattic and they become surrounded
  • alcohols dissolve in polar solvents such as water even though they are covalent because they can form hydrogen bonds - shorter chains are more soluble, only the OH bonds so the carbon part isnt attracted
  • not all molecules with polar bonds dissolve in water as not all of them have bonds stronger than those already present in water such a halogenalkanes
  • non-polar substances dissolve best in non-polar solvents as they form similar bonds
  • non-polar substances dont dissolve well in polar solvents as the bonds already present in the solvents tend to be too strong
  • like dissolves like, substances dissolve best in solvents with similar bonds
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