AS CHEMISTRY UNIT 1
- Created by: Charlotte
- Created on: 07-04-13 16:31
Atoms
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.
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
Acids
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
Redox
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.
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
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
Bonding
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
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
Periodicity
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
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
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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