Higher Chemistry Unit 1

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Highers Highers (Chemistry) Mind Map on Higher Chemistry Unit 1, created by alltimeclara on 01/26/2015.

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Higher Chemistry Unit 1
1 Reaction rates
1.1 Affected by
1.1.1 temperature the average kinetic energy of the particles in a substance an energy distribution diagram could be useful only the particles in the shaded area have enough energy to collide successfully
1.1.2 concentration increasing concentration causes more collisions
1.1.3 particle size the bigger the surface area the faster the reaction
1.1.4 catalysts speed up chemical reactions two types homogeneous same physical state as reactants heterogeneous different physical state to reactants ADSORB reactant molecules onto active site surfaces catalytic poisons bond so firmly to active sites there are no sites for reactant molecules form an alternative activated complex therefore provides an alternative reaction pathway catalysts in industry iron used in haber process to make ammonia platinum used in oxidation of ammonia to nitrogen dioxide in ostwald process nickle used to hydrogenate vegetable oils in the manufacture of margarine zymase used in brewing vanadium oxide used in contact process to make sulphuric acid aluminium oxide used in catalytic cracking turns long chain alkanes into more useful short chain alkenes enzymes biological catalysts large soluble protein molecules works best at an optimum temperature and pH optimum temperature in human bodies is 37*C can be denatured (changes shape) at too high temperatures or pH's
1.1.5 collisions may form an activated complex for a chemical reaction to occur, reactant particles must collide to increase reaction rate, number of successful collisions must be increased Activation energy is the minimum kinetic energy needed for a successful collisiom the difference between the energies of the activated complex and the reactants. highly unstable dpecies formed during a reaction as old bonds break and new bonds form the minimum energy needed to form an activated complex Activated complex is the minimum energy needed to cause a chemical reaction
1.2 1/t
1.2.1 measured in s-1
2 Enthalpy
2.1 combustion
2.1.1 the energy released during complete combustion of one mole of oxygen
2.2 neutralisation
2.2.1 the energy change when an acid an alkali cancel each other out to form one mole of water
2.3 solution
2.3.1 the energy change when one mole of substance dissolves in water
2.4 heat intake/outtake
2.4.1 exothermic gives out heat energy because energy is released to surroundings products have LESS energy than reactants ΔH is negative
2.4.2 endothermic takes in heat energy because energy is absorbed from surroundings products have MORE energy than reactants ΔH is posit8ive
2.5 cannot be measured directly
2.5.1 we can only measure the change in energy ΔH
2.6 potential energy diagrams
3 Patterns in the Periodic Table
3.1 elements in the same group have same number of outer electrons; similar chemical properties
3.2 as you go along a period , an additional proton is added to the nucleus
3.3 covalent radii
3.3.1 along a period atoms get SMALLER increasing number of protons in nucleus attract the electrons more strongly pulling them inwards
3.3.2 down a group atoms get BIGGER there are more electron shells
3.4 ionisation energy
3.4.1 INCREASE along a period increasing nuclear charge makes it harder to remove electrons
3.4.2 DECREASES down a group outer electrons are further from the nucleus and are screened from nuclear attraction y more inner electron shells
3.4.3 the energy needed to remove one mole of electrons from one mole of atoms
3.5 electronegativity
3.5.1 a measure of the ability to attract electrons in a bond
3.5.2 elements at the top right of the table are very electronegative
3.5.3 elements a the bottom left are the least electronegative
4 Bonding, Structure, properties
4.1 we can use electronegativity to make predictions about the bonding between atoms
4.1.1 IONIC = very different electronegativites
4.1.2 COVALENT = very similar electronegativities
4.1.3 POLAR = fairly different electronegativites
4.2 metalls
4.2.1 delocalised outer electrons are free to leave the metal atoms
4.3 elements
4.3.1 nonmetals pure covalent bonding, atoms have identical electronegativities
4.4 structure
4.4.1 discrete covalent sulphur S8 phosphorus P4
4.5 intermolecular/intramolecular
4.5.1 intermolecular hydrogen bonding polar bonds formed between hydrogen and fluorine, nitrogen or oxygen strongest of the intermolecular bonds raises the boiling point bonding BETWEEN molecules van der waals forces molecules become temporarily polar due to a momentary lack of symmetry in the electrons
4.5.2 intramolecular bonding WITHIN molecules covalent bonding
5 The Mole
5.1 relative formula mass
5.2 contains 6.02x10^23 particles
5.3 volume
5.3.1 approximately 24L unless stated otherwise
6 Calculations
6.1 enthalpies
6.1.1 combustion ethanol
6.1.2 solution
6.1.3 nutralisation
7 PPA's
7.1 PPA 1
7.1.1 effect of concentration on reaction rate H2O2 + 2H(+) + 2I(-) ----> 2H2O + I2 sharp colour change starch turns black to blue
7.2 PPA 2
7.2.1 effect of temperature on reaction rate 5(COOH)2 + 6H(+) + 2MnO4(-) ----> 2Mn(2+) + 10CO2 + 8H2O acidified potassium permanganate = purple to colourless all experiments identical except for twmperature evaluation dry beakers to avoid affecting concentrations accurate experiment not possible at room temperature due to gradual colour change
7.3 PPA 3
7.3.1 enthalpy of combustion of ethanol
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