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  • Created by: Charlotte
  • Created on: 07-04-13 16:31


Atomic mass - concentrated in positive nucleus = protons and neutrons

Proton number = number of protons (and electrons) in an atom

Isotope = atoms of the same element but different number of neutrons

Relative isotopic mass: mass of an atom of an isotope of an element on a scale where an atom of Carbon-12 is 12.

Relative atomic mass: the average mass of an atom of an element on a scale where an atom of Carbon-12 is 12.

Relative molecular mass: average mass of a molecule on a scale where an atom of Carbon-12 is 12.

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Moles and Equations

Avogadro constant, Na = number of particles per mole (6.02x1023 mol-1)

Mr (molar mass) = relative molecular mass (gmol -1)

Empirical fromula: smallest whole number ratio of atoms in a compound

Molecular formula: actual number of atoms in a molecule


                 Mass                             Vol                            Moles

              Mr    Mole                  Mole   24(000)                 Vol    Conc.


                                           It is 24000 for cm 3

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Acids - release H+ ions in aqueous solutions (proton donor)

Bases - remove H+ ions from aqueous solutions (proton acceptor)

Common bases: metal oxides, metal hydroxides, ammonia

Alkali = soluble base (releases OH- ions in aqueous solutions)

When acids release H+ ions - you get a salt if it is replaced by metal or ammonium ions.

metal oxide + acid --> salt + water                 metal hydroxide + acid --> salt + water

metal + acid --> metal salt + hydrogen

metal carbonate + acid --> metal salt + carbon dioxide + water

Anhydrous - doesn't contain water of crystalisation

Hydrated - contains water of crystalisation

Water of crystalisation - water molecules in a lattice

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Oxidation numbers: uncombined and identical atom bonds=0, ions=charge, neutral compunds=0, compounds overall=ion charge, combined oxygen=-2, combined hydrogen=+1

Oxidation = loss of electrons - reducing agents are oxidised (donate electrons)

Reduction = gain of electrons - oxidising agents are reduced (accept electrons)

Fe2O3 + 3CO --> 2Fe + 3CO2               Fe reduces = +3 to 0   C oxidised = +2 to +4

Metals generally form ions by losing electrons - forms +ions - higher oxidation number

Non metals generally gain electrons - forms -ions - lower oxidation number

If you react metals with dilute HCL or dilute H2SO4 metal atoms are oxidised (electron loss), hydrogen ions are reduced (electron gain) - forms hydrogen molecules.

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Electron Structure

First ionisation energy: energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms, to form 1 mole of gaseous 1+ ions

Succesive ionisation energies/second: energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions

Ionisation energies are affected by: nuclear charge (attraction), distance from nucleus and electron sheilding

Period = number of shells                        Group = number of electrons in outer shell

1st shell = 2 electrons, 2nd = 8, 3rd = 18, 4th = 32

Orbials in sub-shells can hold up to 2 electrons of opposite spins

Sub-shell s has 1 orbital, p has 3, d has 5 and f has 7 (4s fills before 3d)

Groups 1, 2 and H, He = s block, groups 3 to 8 = p block, transition metals = d block

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Bonding and Structure

Ionic bonding: electrostatic attraction between oppositely charged ions

Group 1 = 1+ ions, group 2 = 2+ ions, group 7 = 1- ions, group 6 = 2- ions

Covalent bonding: shared pair of electrons, dative covalent bonding: donated pair

Electronegativity: ability of an atom to attract the bonding electrons in a covalent bond

Permanent dipoles - covalently bonded, bonded atoms have different electronegativities

Permanent dipole-dipole interactions = weak electrostatic forces

Hydrogen bonding = strongest intermolecular force

Temporary dipole-induced dipole or van der Waals forces = weakest type

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Hydrogen bonding: hydrogen must be covalently bonded to F, N or O. The hydrogen of one molecule forms a weak bond with a lone pair (of F, N or O) from another molecule. This effects the substance properties - higher boiling point, ice = less dense than water

Metallic Bonding: attraction of positive ions to delcocalised electrons

Giant ionic lattice = very strong ionic bonds, high melting point, conducts electricity when molten or dissolved, soluble

Giant covalent lattice = high melting point, good thermal conductor, insoluble

Giant metallic lattice = malleable and ductile, good conductors, insoluble

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Bond Shapes

The shape of a simple molecule is determined by repulsion between electron pairs and lone pairs. Lone pairs repel more strongly!

2 electron pairs = linear molecules, 180

3 electron pairs = trigonal planar, 120

4 electron pairs (no lone pairs) = tetrahedral, 109.5

                        (1 lone pair) = trigonal pyramidal, 107

                        (2 lone pairs) = non-linear, 104.5

6 electron pairs = octahedral, 90

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Atomic radius: decreases as you go accross a period (ionisation energy increases)                                   increases as you go down a group

Periodic table: arranged by increasing atomic number.

Period - the chemical properties show a repeating trend

Groups - have similar chemical properties

There is a decrease in first ionisation energies because the atomic radius is increasing.

The electron sheilding outweighs the nuclear charge

Reactivity - increases down group 2, periods 2 and 3 show similar trends in boiling and melting point

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Group 2

Elements (from group 2) react with water to produce hydroxides, and burn in oxygen to form oxides

Reactivity - increases going down the group                                                                                       ionisation energies decrease because increasing atomic radius and sheilding                     effect

When they lose electrons they form cations (positive ions)

The oxides formed react with water to form metal hydroxides, they dissolve = strongly alkaline solutions

Group 2 carbonates decompose to form the oxide and carbon dioxide

Thermal stability - increases down the group

Ca(OH)2 = used to neutralise acid soils

Mg(OH)2 = used in some indigestion tablets

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Group 7

Boiling point of halogens (e.g. Cl2, Br2 and I2) - increases down the group = van der Waals forces increase, volatility decreases and reactivity decreases

                            Br2(aq) + 2I-(aq) --> 2Br-(aq) + I2(aq)

Oxidation No.:          0           -1   -->     -1          0

Disproportionation: reaction where an element is simultaneously oxidised and reduced  Cl2(g) + H2O(l) <==> HCl(aq) + HClO(aq)          2NaOH + Cl2 --> NaClO + NaCl + H2O                 0                          -1           +1                               0          +1         -1

Chlorine: kills bacteria but harmful if breathed in as a gas, ethical issue = no choice  - Fluoridated water: prevents tooth decay but can cause bone cancers

Test for halides: add dilute nitric acid, then add silver nitrate solution                           Chloride = white precipitate, dissolves in dilute NH3 

 Bromide = cream precipitate, dissolves in concentrated NH3

 Iodide = yellow precipitate, insoluble in concentrated NH3 

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