CHEMISTRY UNIT 1

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  • Created by: charlie
  • Created on: 16-02-14 21:48

triangles

DIAGRAM 

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empirical formula from percentage

-

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working out chemical formulae

  • find valency then use the 'cross method' 

DIAGRAM 

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strong + weak acids

STRONG 

H2SO4 --> 2H + + SO4 2-

0%              100%

strong acids disassociate fully 

WEAK

CH3COOH <--> H + +CH3COO- 

40%                      60%

weak acids partial ionization in water 

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neutralisation

ACID + ALKALI --> SALT + WATER 

e.g. HCL (aq) + KOH (aq) --> KCL (aq) + H2O (l) 

IONIC EQUATION = H+(aq) + OH-(aq) --> H2O (l) 

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salt making

ACID + BASE --> SALT + WATER 

ACID + METAL --> SALT + HYDROGEN 

ACID + CARBONATE --> SALT + WATER + CARBON DIOXIDE 

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water of crystallisation

.

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calculating R.A.M

DIAGRAM 

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3 factors affecting nucleus attraction

1. ATOMIC RADIUS 

  • GREATER DISTANCE = WEAKER ATTRACTION

2. ELECTRON SHIELDING 

  • 'repulsion between electrons in different inner shells which reduces net attractive force from the positive nucleus on the outer shell electron' 
  • MORE INNER SHELLS = LARGER SHEILDING EFFECT 

3. NUCLEAR CHARGE 

  • GREATER NUCLEAR CHARGE = GREATE ATTRACTIVE FORCE ON OUTER ELECTRON 
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electron arrangement

Electrons arranged in shells (energy levels) --> sub shells (sub-levels) --> orbitals (hold 2 e-)  ORBITALS  S orbital 

  • spherical 
  • 1 section X 2e- 

P orbital 

  • 8 shape 
  • 3 sections X 2e- 

D orbital 

  • 5 sections X 2e- 

F orbital 

  • 7 sections X 2e- 
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ions

electrons gained or lost from highest energy orbitals 

BUT... transition ions lose from 4S first 

e.g. 

  • Fe =   [Ar] 4s2 3d6 
  • Fe 2+ =   [Ar] 3d6
  • Fe 3+ =   [Ar] 3d5 
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filling orbital rules

DIAGRAM

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shapes of molecules

'depend on total number of electron pairs around the central atom' 

  • electron pairs repel each other as fair apart as possible 
  • lone pairs repel more apart than electron pairs 

3D MOLECULE 

diagram... 

  • every lone pair reduces bonding angle by 2.5 d.c 
  • N.B (treat double bonds as single bonds when working out shape)
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3 types of inter molecular bonding

1. VAN DER WAALS 

  • temporary dipole to dipole attraction induced 
  • uneven distribution of electrons (constant movement) 
  • BIGGER MOLECULE = GREATER V.D.W 

2. PERMANENT DIPOLE - DIPOLE ATTRACTION 

  • molecules with overall dipole 
  • attractions between peramanent dipoles of neighbouring molecules 

3. HYDROGEN BONDING 

  • permanent dipole - dipole attraction where H atom bonded to very elctronegative atom (F,O,N)
  • polar bond leaves H nucleus expose (1 electron) 
  • strong attraction between lone pair on F,O,N molecule to exposed nucleus 
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hydrogen bonding properties

H2O 

  • anomalously high BP 
  • high surface tension 
  • FROZEN WATER - less dense than liquide water - h-bonds hold water in open lattice (relatively long)
  • good solvent for solutes that can h-bond with water - ionic + covalent molecules 
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ionic, covalent + metallic bonding

IONIC 

  • bond isnt directional 

COVALENT 

  • bond is directional - acting soley between 2 atoms in bond 

METALLIC 

  • form GIANT METALLIC LATTICE - strong metallic bonds 
  • INC ATTRACTION (+3 ions) = INC BP + MP 
  • E.G. metals in group 1 form 1+ ions provide 1 electron to mobile sea
  • metals in gorup 3 form 3+ ions provide 3 electrons to moblie sea (STRONGER) 
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metallic bonding properties

malleable/ductile 

  • bendy - made into wire 
  • DUE TO... free electron still attract metal ions even if they are shifted 

high MP/BP 

CONDUCT electricity

  • delocalised electron cane move throughout 3D lattice + carry charge 
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giant covalent lattices

  • huge networks of covlently bonded atoms 
  • CARBON can form structure - 4 strong covlent bonds 

DIAMOND (allotrope of carbon) - diff. form of same element in same state 

  • each carbon attached to another 4 in tetrahedral shape 
  • high MP, extrememly hard, vibrations can travel (thermal conductor)
  • CAN'T CONDUCT (all electrons held in localised bonds), wont dissolve in any solvent 

GRAPHITE (allotrope of carbon) 

  • carbon atoms in flat hexagon sheets - 3 bonds each 
  • 4th outer electrons delocalised between sheets 
  • sheets of hexagons bonded together by VDW 
  • slippery (Weak bonded sheets), CAN CONDUCT (delocalised electron) 
  • less dense (far apart layers) high MP (strong covalent bonds in hexagons)
  • insoluble in any solvent (covalent bonds too hard to break)
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group 2 element properties

down the group 

  • inc. raidus = more shells of electrons 
  • dec. I.E - bigger radius / more shielding / weaker attraction despite inc. protons
  • inc. reactivity  

MP higher than group 1 

  • stronger attraction between 2+ ions and more electrons 

medium density 

reactions with O2, water + acid 

  • vigourous - REDOX 
  • 2Sr(s) + O2(g) --> 2SrO(s) 
  • Sr(s) + 2H2O(l) --> Sr(OH)(aq) + H2(g)
  • Sr(s) + 2HCl(aq) --> SrCl2(aq) + H2(g) 
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group 2 compounds

OXIDES basic + neutralises acids 

  • NOT REDOX instead is an ACID - BASE REACTION 
  • SrO(s) + 2HCl(aq) --> SrCl2(aq) + H2O(l)

forms HYDROXIDES with water 

  • down group hydroxides more SOLUBLE so realease more OH- ions to water (INC. PH)
  • SrO(s) + H2O(l) --> Sr(OH)2(aq)

CARBONATES all undergo thermal decomposition 

  • down group become more THERMALLY STABLE 
  • SrCO3(s) --> SrO(s) + CO2(g) 
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group 2 elements - limestone cycle

DIAGRAM 

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group 2 hydroxides - uses

Mg(OH)2  pH 8.5 

  • non-toxic = used to NEUTRALISE excess stomach acid 
  • e.g. INDIGESTION TABLETS 

Ca(OH)2  pH 10

  • NEUTRALISE soil acidity = CROPS can grow 
  • NOT CaO as with water is EXOTHERMIC - producing HEAT - kills humus + dangerous 
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halogens

  • non-metals all exist as separate diatomic moleucles e.g. Cl2 

BP INC. down group 

  • inc. no of electrons = inc. VDW 

ELECTRONEGATIVITY DEC. down group 

  • despite more PROTONS, INC. SHELLS (shielding), INC. RADIUS (less attraction) 

OXIDISING power DEC. down group 

  • inc. sheilding + atmoic radius 
  • CHLORINE - (high electronegativity+small size) can bring highest oxidation state in elements

DISPLACEMENT REACTIONS 

  • Cl2 + 2Br- --> Br2 + 2Cl-
  • Cl2 + 2I- --> I2 + 2Cl-
  • Br2 + 2I- --> I2 + 2Br-
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halogens (2)

DISPROPORTIONATE REACTION of chlorine 

  • chlorine is oxidised + reduced 
  • e.g.
  • 2NaOH(aq) + Cl2(g) ---> NaCl(aq) + NaOCl(aq) + H2O(l)

USES OF HALOGENS 

  • Cl2: bleach, solvents, polymers, water purifiying 
  • F2: CFC's, polymers 
  • F-: prevent tooth decay 
  • HF: etch glass 
  • AgBr: photographic film 

TESTING FOR HALIDE IONS (dilute nitric acid + silver nitrate + dilute ammonia+concen. amm.)

  • Cl - : WHITE (soluble in both ammonia)
  • Br - : CREAM (soluble in concen. ammonia)
  • I - : YELLOW (insoluble in both) 
  • Ag+(aq) + X-(aq) ----> Ag+X-(s)
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oxidation numbers

  • UNCOMBINED ELEMENT :     0
  • COMBINED OXYGEN        :    -2 
  • COMBINED HYDROGEN  :     +1
  • SIMPLE ION                      : (charge on ion)
  • COMBINED FLUORINE    :     -1

EXCEPTIONS 

  • OXYGEN + FLUORINE:     oxygen has +2 
  • PEROXIDES                :     oxygen has -1 (H202)
  • HYDRIDES + METAL :   hydrogen has -1 

most electronegative is dominant element 

e.g. CuO(s) + H2(g) -----> Cu(s) + H2O(l)

       +2 -2        0                 0          +1 -2 

Copper reduced + Hydrogen oxidised 

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