Chemistry AS

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• Created by: jp3louis
• Created on: 07-08-17 22:03

Atomic structure and isotopes (C2)

Isotopes - same element with different amounts of neutrons.

Mass number (top) = protons + neutrons

Atomic number (bottom) = Proton number

Cations have fewer electrons than protons so cause a positive charge. Whereas anions has more electrons than protons so cause a negative charge.

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Relative mass (C2)

Relative isotopic mass - mass of a isotope relative to 1/12th of the mass of an atom of carbon-12.

14N has a relative istopic mass of 14 (like 16O has a mass of 16)

Relative atomic mass -   the weighted mean meass of an atom of an element relative to 1/12th the mass of an atom of carbon 12.

it takes into account the percentage abudance and relative isotopic mass of each isotope.

Percetage abudances are found using a mass spectrometr.

To work out the relative atomic mass of clorine, you take the mass number of 35 and times it by the percentage abudance and then add that to the other isotopes such as cl 37. then you divide by 100.

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Formulae and equations (C2)

Binary compound = contains 2 elements only and they end in -ide such as sodium oxide.

Polyatomic ions = contains more than 2 elements such as:

• Ammonium NH4 +
• Hydroxide OH-
• Nitrate NO3 -
• Carbonate CO3 2-
• Sulfate SO4 2-
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Amount of substance and the mole (C3)

1 mole is the amount of substance in 6.02 x 10^23 particles.

Avogadros constant = 6.02 x 10^23

Molar mass is where you add up all of the masses of the elements in a compound eg Na2CO3 = 23 x2 + 12 + 16 x3 = 106.0g mol-1 or the molar mass of carbon is 12 g mol-1

AMOUNT (mol) = MASS (m)(grams) / MOLAR MASS (M)(g mol -1)

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Determination of formulae (C3)

Molecular formula = Number of atoms of each element in a molecule. A moleculeis 2 or more atoms held together by covalent bonds. Eg H2, N2, O2, F2, Cl2, P4, S8

Empirical formula = Simplist whole number ratio of atoms of each element in a compound.

Relative molecular mass = compares the mass of a molecule with the mass of an atom of carbon-12. You add together the relative atomic masses of the elements in a molecule ( H2O = 1 x 2 + 16 = 18) It tends to be for simple molecules such as water and carbon dioxide.

Relative formula mass = Compares the mass of a formula unit with the mass of an atom of carbon-12. It is calculated by adding together the relative atomic mases of the elements in the empirical formula ( NaCl = 23 + 35.5 = 58.5. Ca(NO3)2 = 40.1 + (14 + 16 x3) x2 = 164.1)

Hydrated salts/ Crystalisation (look at page 24 for a run through)

How acurate is an experimental formula when using hydrated salts? If the hydrate salt and anhydrous has different colours then it is easier to see all of the water is lost but you cant see inside of the salt so some water may be left. To solve this, heat at constant mass until no more mass change.

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Moles and volumes (C3) - loads of

1 dm3 = 1000cm3 = 1000ml

n (mol) = c(mol dm-3) x V(dm3)

IF ITS IN DM3 n = c x V(dm3)/1000

Standard solutions = Solutions of known concs

The molar gas volume = volume per mole of gas molecules at a stated temp and pressure.

• RTP is about 20 degreese C and 101 kPa pressure
• At RTP, 1 mole of gas has a volume of 24 dm3 or 24000 cm3
• Therefore, at RTP, the molar gas volume = 24 dm3mol-1

The ideal gas equation = pV =nRT

• p = pressure in Pa,    V = volume in m3,       n = amount of gas molecules (mol)
• R = ideal gas constant (8.31 J mol-1K-1)       T = temp in K
• cm3 to m3 = x10-6
• dm3 to m3 = x10-3 same as with kPa to Pa = x10-3
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Reacting quantities (C3) loads of examples!!!

Stoichiometry = the balanced numbers in an equation give the ratio of the amount.

Percentage yield = % yield = actual/theoretical x100%

Atom economy = how well the chemicals have been utilised. Reactions with higher atom economies produce a large proportion of desired products and few unwanted products.

atom economy = sum of molar mass of desired / sum of molar of all x100

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Acids, bases and neutralisation (C4)

When dissolved in water, an acid will release hydrogen ions as protons.

HCl (g) + aq --> H+ (aq) + Cl- (aq)

A strong acid completely dissociates (like HCl)  HCl (aq) --> H+(aq) + Cl-(aq)

A weak acid partially dissociates (Ethanoic acid) CH3COOH(aq) = H+ (aq) + CH3COO-(aq)

Bases = metal oxides, hydroxides and carbonates along with ammonia (NH3). A base neutralises an acid to form a salt.

An alkali is a base that dissolves in water, releasing its OH- into solution. NaOH(s) + aq --> Na+ (aq) + OH-(aq)

Table of common bases

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acids, bases and neutralisation (2) C4

Neutralisation

h+ ions react to form a salt and water. The H+ ions are replaced by metal or ammonium ions from the base.

Neutralisation of acids with metal oxides and hydroxides

It forms a salt and water ONLY. CuO(s) + H2SO4 (aq) ---> CuSO4 (aq) + H20 (l)

Alkalis

All reactants are in solution and with metal oxides, the overall reaction forms a salt and water only. Acid + alkali --> salt and water.           HCl(aq) + Naoh(aq) --> NaCl(aq) + H2O (l)

Neutralisation of acids with carbonates

Carbonates neutralise acids to form a salt and water and carbon dioxide.

ZnCo3 (s) + H2SO4 (aq) --> ZnSO4 (aq) + H20 (l) + Co2(g)

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Electron structure(C5)

Shells are regarded as energy levels and the energy increases as the shell number increases. The shell number or energy level is called the principal quantum number (n)

An atomic orbital is a region around the nucleus that can hold up to 2 electrons with opposite spins

• 2 e- in shell one (n=1)        S          Number of orbitals = 1 so electrons is 2
• 8 e- in shell two (n=2)        SP        Number of orbitals = 3 so electrons is 6
• 18 e- in shell three (n=3)   SPD     Number of orbitals = 5 so electrons is 10
• 32 e- in shell four (n=4)     SPDF   Number of orbitals = 7 so electrons is 14

Packets of energy are called quanta and a single packet is a quantum. Cr AND Cu both go into the 4S1 subshell instead so they are more stable.

The 3D subshell has more energy than the 4s so the order is 3p,4s,3d

There is a drop between group 2 and 3 as boron has a 2p orbital whereas beryllium has a 2s. This shows boron has a lower energy level and the increased distance results in reduced attraction so reduces ionisation energy. The 2p orbital is shielded by the 1s2 and 2s2 so lowered Ea.

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Ionic bonding and structure (c5)

Ionic bonding - Electrostattic attraction between positive and negative ions. It holds together cations and anions in ionic compounds. A giant ionic lattic contains billions of ions.

• Common cations include - Metal ions (Na+, Ca2+, Al3+) and ammonium ions NH4+
• Common anions include - Non metal ions (cl-, O2-) and polyatomic ions (NO3-, SO42-)

Melting and Boiling points - HIGH

• Solid at RTP as there is insufficent energy to overcome the strong electrostatic forces. The lattices containing ions with greater ionic charges as there is strong attraction (like 2+)

Solubility - SOLUBLE

• Ionic compounds dissolve in polar solvents as the solvent surrounds each ion. With larger charges, the ionic attraction may be too strong so then it wont be soluble.

Electrical conductivity

• Solid - WONT. melted/dissolved will.
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Covalent bonding (C5) look at pg 65

Covalent bonding - Strong electrostatic attraction between a shared pair of e- and the nuclei of the bonded atoms. It occurs in non-metallic elements (H2 and O2), compounds of non-metallic elements for example H2O and CO2. Polyatomic ions like NH4+.

A covalent bond is the overlap of atomic orbitals, each contain 1 e- to give a shared pair of e-

The bond is localised as it only acts between the shared pair of electrons and nucei of the 2 bonded atoms.This results in a molecule.

BORON - it is BF3 because it has the electronic configuration of 1s22s22p1 so only 3 outer-electrons can be paired so therefore it makes boron trifluoride with 6e- in outer shell.

Double covalent bonds - electrostatic attraction between 2 shared pairs of e- and the nuclei of bonded atoms such as CO2

Triple covalent bonds- electrostatic attraction between 3 shared pairs of e- and the nuclei of bonded atoms such as N2 or HCN.

Dative covalent bond - Shared pair of e- have been supplied by one of the bonded atoms. NH4+

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Shapes of molecules and ions (C6) DRAW

A lone pair of e- is slightly closer to the central atom and occupies more space than a bonded pair.

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Electronegativity and polarity (C6)

Electronegativity - The attraction of a bonded atom for the pair of electrons in a covalent bond.

It changes when nuclear charges are different, atoms are different sizes and the shared pair of e- may be close to one nucleus than another.  The pauling electronegativity scale - Across the table : nuclear charge increases, atomic radius decreases. The score is 0 is if the bond is covalent. 0<1.8 is polar covalent and bigger than 1.8 is ionic.

Non - polar bonds - bonded e- pair is shared equally

• The bonded atoms are the same or the atoms have the same/similar electronegativity
• Pure covalent bond is when both atoms come from same element like O2 or H2

Polar bonds - Bonded e- pair is not shared equally

• When bonded atoms are different and different electronegativity values - polar covalent bond
• Such as HCL and the opposite charges are dipoles - permanent dipole.
• H-Cl only has 1 bond but those with more than 2 atoms may have more bonds acting in diff directions.
• Examples on page 76
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Intermolecular forces (C6) pg 80

Intermolecular forces - weak interactions between dipoles of diff molecules. They fall in 3 places:

• Induced dipole-dipole interactions (London forces)
• Permanent dipole-dipole interactions
• Hydrogen bonding

Induced dipole-dipole interactions (London forces)

• They exist between ALL molecules whether non/polar.
• They are instantaneous and then induce another dipole on a neighbouring molecule
• The more e- in each molecule, the larger the instantaneous/induced dipoles, the greater the induced dipole-dipole interactions and the stronger the attractive forces between molecules.

Simple molecular substances - Low BP, non-polar are insoluble, no electrical conductivity

• Made of simple molecules containing a definite number of atoms.
• Simple molecular lattice - atoms held by weak intermolecular forces, atoms are bonded strongly by covalent bonds.
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Hydrogen bonding (C6)

Hydrogen - Special type of permanent dipole-dipole interaction found between molecules with:

• an electronegative atom with a lone pair of e- (oxygen, nitrogen or fluorine)
• A hydrogen atom attached to an electronegative atom like H-O, H-N, H-F
• It acts between a lone pair of e- on an electronegative atom in 1 molecule and a H atom in a different molecule.

Ice is less dense than water

• Hydrogen bonds hold water molecules apart in an open lattice structure
• These are further apart in ice than water
• Solid ice is less dense than liquid water so floats.

Water has a relatively high MP/BP

• Hydrogen bonds are over london forces and more energy is needed to break the bonds
• When ice lattice breaks, the rigid arangement of H bonds is broken.
• It also has a high surface tension and viscocity
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Ionisation energies (C7)

First ionisation energies - Energy required to remove 1 e- from each atom in 1 mole of gaseous atoms of an element to form one mole of gaseous 1+ ions               Na(g) --> Na+(g) + e-

Second ionisation energies - Energy required to remove one electron from each atom in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions        He+(g) --> He2+ + e-

Factors affecting ionisation energy - Nuclear charge - Electron shielding - Atomic radius

Atomic radius - decrease across period due to - sheilding staying the same, no protons increasing and the nuclear att on the outer electorn increases so reduced radius. It increases down a group as theres more shells, more shielding electrons so outweigh the nuclear attraction.

Theres a general increase of ionisation across the period and a sharp decrease between periods.

Trends down a group - atomic radius increa, shielding increa, nuclear att dec, DECREASES

Trends across a period - Atomic radius dec, Nuclear att + charge Inc, INCREASE

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Ionisation energies 2 (C7)

In period 2, the first ionisation energy rises to beryllium before falling to boron. Then it rises again to nitrogen before falling to oxygen and then rises again. This is due to the addition of orbitals from 2s to 2p and then pairing the 2p electrons.

Comparing beryllium and boron

This marks the filling of the 2p sub shell. The 2p subshell of boron has a higher energy than the 2s subshell in beryllium and therefore in boron, the 2p electron is easier to remove than one of the 2s subshells in beryllium. This means the first ionisation of boron is less than beryllium.

Comparing nitrogen and oxygen

This marks the filling of the paring of electrons in the 2p orbitals. In oxygen, the paired electron is easier to remove as it repels the other one. Therefore the first ionisation energy of oxygen is less than nitrogen

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Periodic trends in bonding and structure (C7)

Metallic bonding - The strong electrostatic attraction between cations and delocalise e-

Each atom donates its negative outer shell e- to a shared pool of e- which are delocalised when it is solid. The cations left behind are in the nucleus and inner shell e-. It can also form a giant metallic lattice.

The strength of a metallic bond depends on:

• Charge on metal ions and the size of the ion. A higher charge will mean theres a greater attraction and will be stronger. The smaller the ion, the closer the positive nucleus is to the electrons so there is a stronger bond.
• Down the group of metals, the size of the ion increases and this causes charge density to decrease. This results in a weaker electrostatic force of attraction between metal ions and the delocalised sea of e-.

Properites

• High m/bp but low solubility but it can conduct electricity due to delocalised e-
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Group 2 (C8) write redox reactions

Each group has 2 outer shell e- and thus redox reactions are most common. Each metal is oxidised, loosing 2 e- to form a 2+ ion. It is also called a reducing agent.

The reactivity increases down group 2 when redox reactions occur. This is due to the loss of +2 ions and this requires two ionisation energies. The ionastion energies decrease down the group because the attraction between the nucleus and the outer e- decreases due to an increasing atomic radius + atomic shielding.

Reactions of group 2 compounds

Solubility of hydroxides

Increases down a group so the solutions contain more OH- ions and are more alkaline. It can easily be seen by adding a group 2 oxide to water and a saturated solid will be in the tube. The ph should then be measured.

Using a group 2 compound as a base can neutralise acids. Calcium hydroxide Ca(OH)2 is added as lime to fields to lower the pH of acidic soils.

It is also used as anatacids to treat indegestion.

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The halogens (C8)

At RTP halogen exist as diatomic molecules. As you go down group 7, there is more electrons, more london forces and this means more energy is needed to break bonds so BP is higher. The halogens form lattices with simple molecular structures.

Redox reactions

• Each halogen has 7 outer shell e-. Redox is the most common and they are reduced gaining one electron to form a 1- halide ion with the e- configuration of the nearest noble gas. It is an oxidising agent

Halogen - halide displacement reaction

• The reactivity of halogens decrease down the group. A solution of each halogen is added to aq solution of other halide. Solutions of iodine and bromine can appear the same colour so adding a non-polar solvent like cyclohexane can be added and this allows the halogens to dissolve more and a different colour forms
• cl2 reacts with both Br- and I- whereas Br2 only reacts with I- and I2 doesnt react at all.
• Cl2 is green, Br2 is orange and I2 is purple.
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Qualitative analysis (C8)

Tests for anions

tests based on gases - The carbonate test - Co3 2-

• Carbonates react with acids to form Co2 gas. Add dilute nitric acid to solid and if you see bubbles, it could be a carbonate - bubble through calcium hydroxide and it should go cloudy.

Tests based on precipitates - sulfate test

• Sulfate is soluble in water but barium sulfate BaSO4 is very insoluble. A white precipitate would form is the basis for the sulfate test in which barium ions are added to an unknown compount. Ba2+ ions are added as aq barium chloride/nitrate. If you want to do a halide test, use barium nitrate.

Halide test

• most halides are soluble in water but silver halides are insoluble. aq silver ions react with halide ions to form precipitates. Add aq silver nitrate (AgNO3) to an aq solution of halide. Cl is white, Br is cream and I is yellow.
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Enthalpy changes (C9)

Enthalpy (H) is the measure of heat energy in a chemical system. A chemical system refers to the atyoms, molecules of ions that make up the chemicals

Enthalpy change (delta H) = H(products) - H(reactants)

• If the energy goes from the system to the surroundings, its exothermic (delta H is negative)
• If the energy goes from the surroundings to the system, its endothermic (delta H is positive)

Exothermic - The temperature of the surroundings increase as they gain energy.

Endothermic - The temperature of the surroundings decrease as they lose energy

Activation energy (Ea) - the minimum amount of energy needed to break bonds.

EQUATIONS FOR STANDARD ENTHALPY CHANGES.

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Measuring enthalpy changes (C9)

Kelvin - 0 K = -273 de.C. Ice melts at 273K (0) and water boils at 373K (100)

Calculating an energy change = q = mcT

• Mass of surroundings (m) - this is measured by weighing in g
• Specific heat capacity of surroundings (c) - Energy required to raise the temperature of 1g of a substance by 1K. Good conductors has small values of c - JK-1g-1
• Temperature change of the surroundings (delta T) - K

% error for apparatus

• (error/value used X no times it has been used) X 100

How accurate is the enthalpy change of combustion value?

• there may be heat loss to the surroundings other than water
• Incomplete combustion - carbon monoxide instead of dioxide
• Evaporation of a substance
• No standard conditions.
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Bond enthalpies (C9)

Average bond enthalpy is the energy required to break 1 mole of a specified type of bond in a gaseous molecule. Bond enthalpies are ALWAYS endothermic so are positive.

Bond breaking and making

• Energy required to break bonds - endothermic - deltaH is positive
• Energy required to make bonds - exothermic - deltaH is negative

Limitations

The makin/breaking of bonds may be in different environments so the results differ

Calculating enthalpy changes from average bond enthalpies

The enthalpy change of reaction (delta rH) can be found by calculating the bond enthalpies of the bonds in the reactants and the products.

• For a reaction involving gaseous moloecules of a covalent substance :
• delta rH = E (bond enthalpies in the reactants) - E (bind enthalpies in the products)
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Hess' law and enthalpy cycles (C9)

Hess law states that if a reaction can take place by 2 routes and the starting and finishing conditions are the same, the total enthalpy change is the same for each route.

LOOK IN BOOK FOR EXAMPLES

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Reaction rates (C10)

rate = change in conc/ time (units are mol dm-3/ s = mol dm-3 s-1)

If the rate increases, theres more successful collesions per unit of time.

The ROR is fastest at start of reaction as they are at the highest conc.

Factors affecting rate of reaction

• Conc
• Temp
• Use of catalust
• Surface area of solid reactants.

Some collisions are not effective as the particles have no colided with the correct orientation or there is insufficent energy to overcome the activation energy.

If there is a decrease, the graph will go left and higher and if its increased, it will be right and lower.

You can monitor by balance or collecting gas.

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Catalysts (C10) (equations)

Homogeneous catalyst

Same physical state as the reactants. It reacts with the reactants to form an intermediate and this intermediate breaks down to give the product and regenerate the catalyst.

Examples include : making esters with sulfuric acid as a catalyst or ozone depletion

Heterogeneous catalysts

Different physical states than reactants. They are usually solids and come in contact with gas/liquid reactants. Reactant molecules are adsorbed onto the catalyst where the reaction takes place. Then they leave the surface by desoprtion

Autocatalysts are when a product acts as a catalyst.

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The boltzmann distribution (C10)

No molecules have 0 energy and that the area under the curve is equal to the total number of molecules. There is no maximum energy for a molecule.

Temperature - look in book for the diagram

• At higher temperatures - More molecues have an energy greater than or equal to the activation energy. Therefore, a greater proportion of collisions will lead to a rection so it increases the ROR. Collisions will also be more frequent as the molecules are moving faster.

Catalysts

• They lower Ea as they provide an alternative route with a lower energy. This then allows more molecules to exceed the new lower Ea
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Dynamic equilibrium and le chatelier's principle (

Dynamic equilibrium

The rate of forward reaction is equal to the rate of the reverse reaction. The concentrations of reactants and products do not change. The system must be a closed system so it is isolated from its surroundings. The ROR is fastest at the start.

le Chatelier's principle

The equillibrium will alter to oppose the effect of change.

The effect of concentration/temperature of equilibrium - ONLY AFFECTS GAS SYSTEMS

Changing the concentration will make it go forwards or backwards. If its an exothermic reaction, its releasing energy so the temp will increase compared to endothermic. Increase in temp will move it in the endothermic way.

The effect of catalysts on equilibrium - doesnt affect it but will speed the reaction up.

The effect of pressure - based on how many gaseous moles on each side. High - low

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The equilibrium constant Kc (part 1)(C10) example

Equilibrium law

The exact position of th equilibrium is calculated using this law. aA + bB = cC + dD

Kc = [C]c [D]d / [A]a [B]b          or see it as [products] / [reactants]. The brackets means for concentration of.

For example :

Kc tells us the relative proportions of reactants and products in the equilibrium system. A Kc value of 1 indicates the position of equilibrium is halfway between products and reactants.

If it is greater than (>) 1 when the equilibrium is towards the products.

If it is less than (<) 1 then it is towards the reactants.

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Organic chemistry (C11)

Carbon can have 4 covalent bonds that can be single, double or triple.

Hydrocarbon = Contains carbon and hydrogen only

Saturatec = Single bonds only.  Unsaturated = Double/triple bonds

Homologous series = family of compounds with similar chemical properties whose members differ by -CH2-. The simplist is alkanes.

Functional group = Responsible for the chemical properties.

Hydrocarbons - naming them

• Aliphatic - carbon atoms joined in un/branched chains or non-aromatic rings
• Alicyclic - Carbon atoms are joined in a ring with/out brances
• Aromatic - Some are found in benzene ring.
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alkene

Alcohol

Haloalkane

Aldehyde

Ketone

Carboxylic acid

Ester

Acyl chloride

Amine

Nitrile

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Representing the formulae of organic compounds (C1

• Molecular formula
• Shows the number and type of atoms of each element. It doesn't show how the molecules are joined together so for example ethanol is C2H6O
• Empirical formula
• Simplist whole number ratio of atoms.  Glucose is CH2O (molecular is C6H12O6)
• General formula
• Simplist algebraic formula of any memeber of a homologous series. alkanes is CnH2n+2
• Displayed formula
• Shows the relative positioning of all the atoms (basically drawing it out )
• Structural formulaSmallest amount of detail necessary
• Skeletal formula
• Simplified formula ( just lines)
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Isomerism (C11)

Structural isomers = Compounds with the same molecular formula but different structural formula.

Butane and 2-methylpropane are structural isomers that have the same molecular of C4H10

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Introduction to reaction mechanisms (C11) (look in

Covalent bond = A shared pair of electrons and can be broken by homolytic or heterolytic fission.

Homolytic fission

• Each of the bonded atoms take one of each electrons in the bond.
• Each atom now has a single unpaired electron
• This is now called a radical

Heterolytic fission

• both electrons go to one of the bonded atoms
• The atom that doesn't take electrons is a positive ion
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Introduction to reaction mechanisms (2) (C11) (loo

Types of reactions

Addition - Two reactants joing together to form 1 product. A molecule is added to the unsaturated alkene, breaking the double bond.

Substitution - An atom or group of atoms is replaced by a different atom.

Elimination - The removal of a small molecule from a larger one.

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Properties of the alkanes (C12)

General formula = CnH2n+2

The bonding

• Each carbon bond is joined by 4 other atoms by single covalent bonds called sigma bonds
• The sigma bond is the result of an overlap of 2 orbitals. It is directly between bonded atoms
• A sigma bond has 2 electrons

The shape of alkanes

• Tetrahedarl at 109.5. The sigma bond means it can move in any direction.

BP in alkanes

• BP increases as the chain increases because of london forces.No branches means more surface area of contact and stronger london forces. Branches means fewer points of contact.
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Chemical reactions of the alkanes (C12)

Alkanes do not react with common reagents due to C-C and C-H sigma bonds are strong. The C-C bonds are non polar and the electronegatvity of carbon and hydrgoen is so similar that the C-H bond can be considered non-polar.

Combustion of alkanes

They react with lots of oxygen to prevent carbon dioxide and water.

Reactions of alkanes with halogens

Alkanes react with halogens. UV radiation gives the initial energy

Mechanism for bromination of alkanes - RADICAL SUBSTITUTION - initiation, propagation, termination ( all on back of que card)

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The properties of alkenes and stereoisomerism (C13

Alkenes are unsaturated hydrocarbons with at least one double bond. The gen form is CnH2n

For each carbon atom in the double bond, 3 of the 4 electrons are used in 3 sigma bonds. 1 to another carbon atom of the double bond and the other 2 electrons to 2 other atoms either carbon or hydrogen. This leaves 1 electron on each carbon atom to form a pi bond sideways over the 2 p -orbitals. The pi bond electron density is concentrated above and bellow the bond.

The shape is a trigonal planar so 120 degreese.

Stereoisomers = same structural formula but diff arangement in space.

E/Z is only with a double bond and there must be different groups attached to each carbon in the double bond whereas optical can be with a lot more. E is oposite sides whereas Z is the same side

Cis - trans occurs when one of the attached groups on the carbon must be hydrogen. Cis is the  isomer and trans is the E isomer.

Cahn-Ingold-Prelog rules state the priority is based on atomic number.

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Reactions of alkenes (C13)

Alkenes are reactive due to the pi bond. Being on the outside of the bond, the pi electrons are more exposed than sigma so are easier to react by addition reactions. The pi bond has a smaller enthalpy too so is weaker to break.

ADDITION REACTIONS OF ALKENES

Hydrogenation with a nickel catalyst - added hydrogen and passed on catalyst at 423K

Halogenation with cl or br. - Bromine test will go orange to colourless.

Hydrogen halides - Gaseous hydrogen halides at RTP will form haloalkanes.

Hydration - Add steam over H3PO4 catalyst.

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Electrophillic addition in alkenes (C13) DRAWINGS

Alkenes undergo addition reactions to form saturated compounds - ELECTROHILLIC ADDITION

The pi bond means a high electron density is above and below the double bond meaning it attracts electrophiles.

Electrophile = Atom or groups of atoms that is attracted to an electron rich centre and accepts an electron pair. It is usually a positive ion or a partial positive charge.

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Polymerisation in alkenes (C13)

ADDITION POLYMERISATION produces long saturated chains. Additional polymers have high molecular masses.

Poly(ethene) - Made from heating polymers at high pressure. - bags, kids toys, shampoo bottles.

Poly (chloroethene) also known as PVC can make a polymer flexible or rigid.

Poly(propene) used to make kids toys, fibres for ropes and packing crates.

Poly(tetrafluroethene) is made to coat non-stick pans and cable instulation.

The lack of reactivity means they can store foods but affects the environement. THey can be recycled to stop fossil fuels. They have to be sorted by type, washed and chopped. PVC recycling is hazardous due to the Cl inside and cant be burnt or else it will be HCl. new technology dissolves it in a solvent and then recovered via precipiation from the solvent. Some polymers can be recylced as fuel. Feedstock recycling describes the chemical and thermal processes to reclaim monomers and gases from waste. These can be stored as raw materials.

Biodegradable and photodegradable polymers also now exisist.

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Properties of alcohols (C14)

The hydroxyl group is responsible for both the physical and chemical properties in alcohols. Methanol (Ch3OH) is the simplist and is converted into polymers paints and solvents.

Physical properties

Compared to alkanes, alcohols are less volitile, have higher mp and greater water solubility. These differences become smaller as the length of the carbon chain increases. These differences can be explained by considering the polarity of the bonds in both alcohols and alkanes.

• Alkanes are and have non-polar bonds (H and O are similar electronegativity.) Theres weak london forces.
• Alcohols are polar and have polar bonds (O and H are different electronegativity) so theres weak london forces but also hydrogen bons
• Intermolecular hydrogen bonds holds alcohol together so this means more energy is needed to turn it into a gas.
• Hydrogen bonds form bonds with water so alcohol is water soluble.

Primary, secondary and tertiary alcohols.

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Reactions of alcohols (C14)

Combustion of alcohols - burn in oxygen to produce co2 and water. It is exothermic. As the number of carbons atoms increases, the quanitiy of heat per mole also increases. Primary and secondary alcohols can be oxidised with an oxidising agent. It is potassium dichromate (k2Cr2O7) and diluted sulfuric acid (H2SO4). it will go ORANGE TO GREEN

Oxidation of primary

• Aldehydes - Gently heated and distilled
• Carboxylic acids - Heated under reflux with excess potassium dichromate

Secondary

• Ketones - heated under reflux

Dehydration - Heated under reflux in the presence of an acid catalyst (sulfuric). Its an elimination reaction.

Substitution reactions - React with h halides to form haloalkanes. Heated under reflux with sulfruic acid and a sodium halide and the hydrogen bromide is fomed in situ.

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The chemistry of the haloalkanes (C15)

Reactivity - types of nucleophiles in clude hydroxide ions, water and ammonia (NH3)

C-halogen bonds means they are more electronegative. The elctron pair is closer to the halogen atom so it is polar. In haloalkanes, the carbon is slightly positive and can attract speices with a lone pair of electrons called NUCLEOPHILES. They are attracted to an electron deficient carbon. When a nucleophile replaces the halogen in the subsittuiton reaction, a new compound is formed containing a diff functional group - NUCLEOPHILLIC SUBSTITUTION

Primary haloalkanes undergo nucleophillic sub with a variety of nucleophiles. Hydrolysis occurs where the halogen is replaced by an -oh group.

Idoalkanes react faster than bromoalkanes but they react faster than chlroalkanes. Fluroalkanes are unreactive. To work out the rate of reaction, adding aq silver nitrate to the water and were furoalkanes are unreacive in water, the presence of an ethanol solvent is added. It allows there to be 1 layer rather than 2.

Chloro  - white precip forms slowly, bromo - cream precip forms faster, ido - yellow forms quickest,

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Organohalogen compounds in the environment (C15) D

The ozone layer

10-40 Km above the earths surface and it absorbs the UV. In the stratosphere, ozone is continusly broken down by uv radiation. o2 -> 2O. A steady state is set up involving O2and oxygen radicals o2 + o = O3

CFC and the ozone layer

Used to be used in fridges, air craft and aerosols. CFC are stable because of the strength of the carbon-halogen bonds. They remain stable until the stratosphere where they break down formind cl raicals which break down the ozone layer. The UV lights provide energy so the carbon- halogen bond breaks by homolytic fission.

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Mass spectrometry (C17) draw and write

In a mass spectrometer, an element looses and electron to form a positive ion, the molecular ion. The mass spectrometer detects the mass-to-charge ratio (m/z) of the molecular ion which gives the molecualr mass of the compound.

To find the molecular mass, the molecular ion peak (M+ peak) has to be located and its the highest peak on the right hand side. After this peak is the m+1 peak which is because of the small amount of carbon-13 isotope. Some molecular ions break down during fragmentation. The other peaks are caused by fragment ions.

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Infrared spectroscopy (C17) write and draw

Covalent bonds posses energy and can vibrate around a central point. A stretch is a rhythmic movement along the line between 2 atoms so the distance increases and decreases. A bend is a change in the bond angle. The amount it stretches or bends depends on the mass of the atoms as heavier atoms vibrate more slowly than lighter atoms and also the strength of the bond - stronger bonds vibrate faster than weaker bonds.

Any bond can only absorb radiation that has the same frequency as the natural frequency of the bond.

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Ionisation energies (C7) (2)

Second ionisation energy - Energy require to remove 1 e- from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions

Ionisation energies down a group

• Atomic radius increases
• Shielding increases
• Nuclear attraction on outer shells decrease
• First ionisation energy DECREASES

Ionisation energies across a period

• Nuclear charge increases
• Nuclear attraction increases
• Atomic radius decreases
• First ionisation energy INCREASES

first ion energy is less for boron than berylium as theres a pair of electrons in boron so repulsion

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Periodic trends in bonding and structure(7 pt2)

Giant covalent structures

• Boron, carbon and silicon form giant covalent lattices - all 109.5 (tetrahedral)
• M/BP is very high
• Solubility - none as the covalent bonds are too strong
• Conductivity - none except carbon (graphene and graphite) as one e- is avaliable to conduct

Peridic trends in melting points

• Across period 2 and 3
• it increases from group 1 to 4
• Theres a sharp decrease between 4 and 5
• then its low from 5 to 0
• The sharp decrease shows the change from giant to simple molecular structure
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The halogens (C8) 2 (wiite equations

Trends  - reacitvity decreases down the group

• Increase in size down the group as theres more protons and electrons. This leads to more london forces and more energy is needed to break them THEREFORE BP INCREASES.
• IONISATION ENERGIES DECREASE DOWN THE GROUP as the number of shels increase - more shielding electrons and this outweighs the nuclear charge
• ELECTRONEGATIVITY DECREASES DOWN GROUP

Disproportion - redox reaction where the same element is both oxidised and reduced. The reaction of chlorine with water and with cold, dilute NaOH are two examples of reactions. Look for an element that appears on one side of the equation but twice on the other side.

The reaction of Cl2 with H2O :

The reaction of Cl2 with cold, dilute NaOH:

Benefits of cl : kills bacteria and saves lives but is carcinogens and toxic.

Tests for halide ions - add aq silve ions to form a precipitate : Ag+(aq) + X-(aq) --> AgX(s)

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Qualitative analysis (C8) (2)

Sequence of tests

• Carbonate - sulfate - halides this is because:
• Carbonates - looking for effervescene of co2 gas and neither a sulfate or halide will produce bubbles with dilute acid. No bubbles means you move on
• Sulfate - looking for a white insoluble precipitate and if you do this on a carbonate, you will get the same precipitate so its important to do carbonates first.
• Halide - lookinfg for precipitate that could form in the sulfate  and carbonate test.

Mixture of ions

• Carbonate test - add dilute nitric acid until the bubbling stops.
• Sulfate test - Add excess Ba(NO3)2 then filtre to remove barium sulfate.
• Halide test - Add AgNO3 an any carbonate/sulfate ions will be removed. Add NH3 to confirm which halide you have.

Tests for cations - Ammonium ion NH4+

• When heated, aq NH4+ and aq OH- form ammonia gas NH3. add aq Naoh to NH4+ then use damp red lithmus paper and it should go blue
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