Chemistry summary

Chemistry

Topic 1 States of matter

Brownian motion- random movement particles/ diffusion

Atom: smallest particle can’t be broken by chemical means

Molecules: two or more atoms

Ions: charged particles

Heating curve for water: melting oºC/evaporation/ boiling 100ºC/condensation below 100ºC/freezing or solidifying below 0ºC. At boiling-melting points temperature the

same as used to break bonds .

Solid: particles packed fixed pattern+ strong bonds hold in position+ only move to and fro

Liquid: particles slide past each other+ weaker bonds+ not in fixed pattern

Gas: particles far apart+ move very quickly+ no forces+ collide/bounce off

Melting: solid heated, particles energy vibrate more, solid expands, particles break bonds, no fixed position

Boiling: liquid heated particles energy, move faster, collide more often, bounce further apart, liquid expands, particles use energy to overcome forces.

Evaporating: some particles have more energu and able to escape quicker as a gas.

Condensing and solidifying: gas cooled, looses energy, move slowly, when collide little energy to bounce away, stay close and liquid forms.

Takes more heat to boil than to melt as more energy to separate particles completely than to free from lattice.

Gas pressure- depends on: temperature+ volume.

- Temperature: gain more energy particles move faster, hit walls more often with more force.

- Volume: smaller space, particles hit wall more often.

Difussion- depends on:

- Mass of particles: lower = faster – Cotton wool soaked in ammonia solution ( ammonia gas) Cotton wool soaked hydrochloric acid ( HCL gas) when meet smoke forms closer to HCL cotton.

- Temperature: higher = faster – heat energy, move faster

Mixture: more than one substance, not chemically combined.

Solution = solvent (water) + solute ( sugar)

If solid soluble in water, gets more soluble as temperature rises. Water aqueous solution other solvents : ethanol, propane, white spirit. Volatile= evaporate easily at room temperature

Pure substance: no particles of other substances mixed in it. Check by:

- Melting boiling point: pure has a sharp one/ if impure melting point falls and boiling point rises

Impurity : unwanted substance mixed with substance wanted.

Separation methods:

SOLID FROM LIQUID

1. Filter: eg: chalk from water. Chalk trapped filter paper = residue. Water = filtrate2. Centrifuge: small amounts of solid and liquid. Test tubes spun quickly solid flung

to bottom.

SOLID FROM SOLUTION

1. Evaporate: evaporate solvent by heating eg: obtain salt from aqueous solution.2. Crystallise: eg: Coppert (II) sulphate, heat to get rid some water and make

solution concentrated, leave to cool, crystal appear, remove by filtering.

SOLVENT FROM SOLUTION

1. Distil: Eg: salt water: Solution heated in flask, boils, steam rises to condenser, salt left behind. Condenser cold so steam condenses to water. Water drips into beaker = distilled water.

LIQUIDS FROM EACH OTHER

1. Fractional distillation: Eg: mixture of ethanol and water. Heat mixture, ethanol boils first, some water evaporates all vapors rise up column and condense on glass beads. Glass beads get hot enough to let ethanol by – water still condenses. Ethanol goes into condenser, where condenses and drips into beaker. Used in industry to refine crude oil and producing ethanol.

DIFFERENT SUBSTANCES FROM A SOLUTION

1. Chromatography: ink travels across paper at different rates so separates into rings. Filter paper = chromatogram. Used to identify substances.

- Non- moving stationary phase: paper- Moving, mobile phase: the mixture you want to separate.

When identifying colorless substances ( amino acids) use locating agent: ninhydrin

Use Rf values = distance moved by amino acids

Distance moved by solvent

Topic 2The atom

Element: only contains one kind of atom

Compound: atoms of different elements joined together, chemically combined

Formula: symbol for compound

Atom has: nucleus ( protons & neutrons ) + cloud of electons

Nucleon number = protons + neutrons

Proton number/ Atomic number = protons

- You can identify atom by number of protons.

Isotopes: atoms of same element with different number of neutrons. Eg:Carbon12,13,14

Carbon 14 is Radioactive: nucleus unstable, breaks down/ decays giving out radiation

Half life: how long it takes for half the radioactive substances to decay.

Radiation can harm: large doses give radiation sickness: vomit, tiredness, hair fall, gums bleed, die within weeks, cancer.

Radiation can be used:

- Check for leaks- add radioisotopes to oil/gas, at leak check using Geiger counter. Tracer radioisotopes.

- Cancer treatment: radiotherapy, rays give out kill cancer cells.- Find age old remains: living things have carbon atoms, some c-14.

When dies doesn´t take in no new carbon atoms, but c-14 decay and we measure their radiation. Carbon dating.

Electons circle from nucleus at different energy levels= electron shells. The one closest to nucleus is lower. The further is the higher energy level.

1st shell- 2 electrons

2nd shell and more- 8 electrons

Valency electrons: outer electrons, shows how element reacts.

Group = number of outer electrons

Period = number of shells

Topic 3 atoms combining

Atoms react to obtain full outer shell.

Ion: charged particle- unequal number of protons and electrons.

BETWEEN METAL + NON-METAL IONIC BOND

1. Eg: Sodium + Chloride = Sodium chloride (Na + Cl = NaCl)

Sodium looses one electron, they have opposite charges and attract each other. This force of attraction = ionic bond

2. Eg: Magnesium + oxygen = magnesium oxide ( Mg + O = MgO)

Magnesium loses 2 electrons and gives to oxygen.

3. Eg: Magnesium + Chlorine= magnesium chloride ( mg + 2Cl = MgCl2)

Hydrogen and metals form positive ions Non-metals form negative ions Group 4, 5 don’t react would have to gain/ loose too many electrons Group 0 don´t react atoms have full shells. Transition elements form more than one ion:

- Zinc only Zn2+ - Silver only Ag+- Copper- Cu+ and Cu 2+- Irons – Fe 2+ and Fe 3+

Compound ions- groups of joined atoms

1. NH4+ Ammonium ion2. OH- Hydroxide ion 3. NO3- Nitrate ion 4. SO4 2- Sulphate ion

5. CO3 2- Carbonate ion6. HCO3- Hydrogencarbonate ion

Ionic compounds are hold in strong lattice.

Ionic Bonds are strong and they form a giant structure.

Ionic solids are all crystalline.

Ionic compounds have high melting and boiling points, as bonds are strong and lots of heat needed to break the lattice.

Ionic compounds are usually soluble in water, water attracts ions away from lattice and ions can move freely.

Ionic compounds can conduct electricity when melted or dissolved. Ions are free to move and are charged.

IN NON METALS COVALENT BOND

They share electrons. Positive nuclei of each atom attracts electrons. Force attraction created holds atoms together.

Molecule: group of atoms held together by a covalent bond.

Molecular element: element made of molecules

Diatomic : elements made up of molecules containing two atoms

- Hydrogen H2 (share two , single bond)

- Chlorine Cl2 (share two, single bond)- Oxygen O2 (share four, double bond 0=0)- Nitrogen N2 (share six, triple bond )

Covalent compounds: made of atoms of different elements joined by covalent bonds

- Methane Ch4

Tetrahedral shape, four pairs shared electrons repel each other

- Ammonia NH3

Pyramidal, hydrogen atoms pushed a little closer because two non bonding electrons in nitrogen stronger repelling effect than bonding electrons.

- Water H20

Angular shape, non-bonding electrons repel more strongly.

- Hydrogen chloride HCL

Linear, only two atoms

- Carbon Dioxide CO2

Linear , two groups of bonding electrons repel each other. O=C=O

Molecular solids are held in a lattice, in a regular pattern, with weak intermolecular forces.

Molecular solids make crystals

Molecular solids have a low meting and boiling point, as weak forces

Molecular solids don’t conduct electricity, molecules aren’t charged.

COVALENT STRUCTURES

1. Diamond- carbon atom, forms covalent bonds to four others. - Very hard strong bonds, hardest substance- High melting point- Can’t conduct electricity.

2. Silica- naturally as quartz, main mineral in sand- SiO2

3. Graphite- carbon - Soft slippery, made of sheets slide over each other- Conductor electricity- Dark colour

Diamond & Graphite are allotropes of carbon: two forms of same element

IN METALS METAL BONDING

- Low melting point , weak forces- Giant structures ( sodium chloride & diamond ) higher melting points. - Free electrons move through lattice- Metallic bonds – force attraction between ions and electrons. - Regular arrangements of ions = metals are crystals

- Metals: high melting points, energy to break up lattice : malleable, ductile : good conductors of heat ( free electrons) : good conductors of electricity ( free electrons)

Topic 4. Periodic Table

Columns = groups

Rows = periods = nº shells

Oxides of metals are basic ( react with acids to form salts)

Oxides of non metals are acidic ( react with alkalis to form salts )

- Aluminium oxide in the middles reacts both alkalis/ acids to form salts it is amphoteric oxide

GROUP 1. ALKALINE METALS

- Good conductors heat/ electricity- Softer- Lighter, low density- Low melting, boiling point- Form ions charge 1+

- Ionic compounds white solids dissolve water to give colorless solutions

React violently water- give hydrogen and a hydroxide ( lithium floats, sodium shoots across water, potassium melts and hydrogen catches fire) Hydroxide gives alkaline solution.

React with chlorine: burst into flame React with oxygen: burst into flame ( lithium – red flame, sodium- yellow flame,

potassium- lilac flame)

Downwards group:Softness , density increasesReactivity increases Melting point and boiling point decreases

GROUP 7 HALOGENS- Coloured gases- Fluorine yellow, chlorine green, bromine red, iodine

purple- Poisonous- Brittle and crumbly in solid form+ don’t conduct electricity - Diatomic molecules

Fluorine- pale green solid, Chloride- yellow solid, Bromide- red brown solid, Iodide- black solid.

Downwards in group

Size and mass increaseDensity increasesMelting and boiling point increasesReactivity decreases as smaller atom the easier to attract electron.

GROUP 0 THE NOBLE GASES- Colourless gases, occur naturally in air- Monatomic ( single atoms)- Unreactive, full outer shells, inert

Downwards in group Increase size and massDensity increasesBoiling points increase

UsesAs inert, sageGlow when current pass through them at low pressure

Helium: fill balloons, airships as lighter than air & doesn’t catch fire Argon: provide inert atmosphere in tungsten light bulbs + protect metals being

welded Neon: advertising signs- glows red ( other colours if mixed with other gases) Krypton: lasers ( eye surgery) Xenon: light like bright daylight, blue tinge (car headlamps, lighthouse lamps…)

TRANSITION ELEMENTS

Metals: irons, copper, nickel.

- Hard, tough, strong- High melting points- Malleable + ductile- Good conductors of heat and electricity ( silver best, next copper)- High density

Less reactive group 1- don’t corrode readily water or air ( except iron) No clear trend in reactivity Colored compounds Make ions different charges ( eg copper, iron) Form more than 1 compound with another element ( copper (I) oxide, CuO &

copper (II) oxide CU2O)

Form complex ions.

Uses

- Structures eg: bridges, buildings.. Iron and steel ( alloy) - Alloys eg: chromium + nickel + iron = steel- Conductors of heat, electricity ( copper electricity cables, steel

radiators)- Catalysts eg: Iron in ammonia.

Topic 5. The Mole

Write formula using valencies

Eg: what is formula of aluminium oxide?

Valencies: aluminium 3: oxygen 2

AlO -----balance------ Al2 O3

Relative atomic mass Ar of an element is the average mass of its isotopes compared to an atom of carbon-12

- neutron number ( top nº) Relative molecular mass Mr is mass of its ions

Percentage composition

Eg : Methane CH4 :

Mass of carbon as fraction of total mass= mass of carbon

Total mass

Carbon Ar = 12

Methane Ar = 16

So 12 ÷ 16 = 0.75

Mass of hydrogen as a fraction of total mass mass hydrogen

Total mass

Carbon Ar= 1, BUT 4 CARBON ATOMS SO Ar = 4

So 4 ÷16 = 0.25

Percentage composition of methane is 75 % carbon 25% hydrogen.

Every pure sample of a given compound always has exactly same composition.

Mole is 6.02 x 10 23. Obtained by weighing the Mr or Ar in grams

Number of moles = mass ÷ mass 1 mole ( Ar)

Empirical formula

- Shows simplest ratio in which atoms combine

Eg: 32 grams of suphur combine with 32g of oxygen to form an oxide of sulphur.

Elements that combine sulphur oxygen

Masses that combine 32 g 32g

Relative atomic masses 32/ 32 = 1 32/ 16 = 2

Ratio in which atoms combine 1:2

Empirical formula SO2

Eg: an experiment shows that compound Y is 80% carbon and 20% hydrogen

Elements that combine carbon hydrogen

Masses that combine 80 g 20g

Relative atomic masses 12 1

Moles that combine 89/12 = 6.67 20/1 = 20

Ratio in which atoms combine 6.67:29 or 1:3

Empirical formula CH3

Experiment to find empirical formula: Eg: magnesium oxide.

- Weight a crucible + lid empty. Add coil of magnesium weigh again to find its mass

- Heat crucible. Raise lid carefully at intervals to let oxygen in. Magnesium burns lightly.

- When burning complete let crucible cool. Weight again . Increase in mass = oxygen.

Weighing shows that 2.4 g of magnesium combined with 1.6 g of oxygen. Do table & find empirical formula.

Formula of ionic compound is always the same as its empirical formula. Molecular formula shows the accrual numbers of atoms that combine to form a

molecule.

To find: Mr ÷ empirical mass for substance. This gives N. Multiply numbers in the empirical formula x N.

1.Eg: A molecular compound has the empirical formula HO. It’s relative molecular mass is 34. What is the molecular formula?

( Ar H = 1, O=16)

Empirical formula HO, empirical mass = 17

So 34/ 17 = 2

So molecular formula is 2 x HO or H2 O2

2.Eg: Octane is a hydrocarbon. It is 84.2% carbon and 15.8% hydrogen by mass. It’s Mr is 114. What is molecular formula?

Elements Carbon Hydrogen

Masses 84.2 15.8

Ar 12 1

Moles 84.2/12 = 7.02 15.8/1 = 15.8

Ratio 1: 2.25 4:9

Empirical formula C4H9

Empirical mass = 12x4 + 9 = 57

114/ 57 = 2

Molecular formula = 2 x C4H9= C8H18

Concentration of solution= amount of solute, in grams or moles disolved in 1 dm3 of solution

Concentration( mol/dm3) = amount of solute (mol) / volume of solution (dm3)

1 dm3 = 1 litre

= 1000 cm3

=1000 ml

Topic 6. Chemical equations

Physical change- particles mixed but not bonded. No new product, easy to get original ones.

Chemical change- New chemical substance produced. Energy taken in ( endothermic) or given out (exothermic). Change difficult to reverse.

1 mole of gas occupies the same volume, at same temperature and pressure. At rtp this volume is 24 dm 3

The molar volume of a gas at rtp is 24 dm 3

Volume at rtp (dm3) = number of moles x 24 dm3

% yield = actual mass / calculated mass x 100

Yield = amount of product obtain from a reaction

% purity = mass of pure product/ mass of impure product x100

Topic 7 . Redox reactions

1. Combination or synthesis: substances react to form 1 compound. Needs heat

2. Decomposition: substance breaks down to simpler substances. Eg: Silver chloride light silver+ chlorine 2AgCl light 2Ag + Cl 2Silver chloride is white solid. Ag s grey.

3. Combustion: Substance catches fire and burns on reaction oxygen or other gasEg: magnesium burns oxygen white flame2Mg + O2 = 2MgO

4. Displacement: Element swoops places with another in a compoundEg: copper + silver nitrate = copper (II) nitrate + silver Cu + 2 AgNo3 = Cu ( NO3)2 + 2 AgThis occurs because copper is more reactive than silver. Solution turns blue due to copper ions.

5. Precipitations: Solutions react giving insoluble productEg: silver nitrate + sodium chloride = silver chloride + sodium nitrate AgNO3 + NaCl = AgCl + NaNO3

6. Neutralisation : Acid reacts another substance, destros acidity giving water as product. Eg: NaOH + HCL = NaCl + H2 O

3 groups:Redox reactions ( electrons transfer) 1- 4Precipitation reactions ( no electron transfer) 5Neutralisation reactions ( water as product, no electron transfer) 6

OIL RIG

Oxidation: Gains oxygen

Loss of hydrogen

Loss of electrons

Reduction : Looses oxygen

Gain of hydrogen

Gain electrons

Half – equations:

1. Reaction between magnesium and oxygen

Mg Mg 2+ + 2 e-

O + 2 e- O 2-

2Mg 2 Mg 2+ + 4 e-

O2 + 4e- 2 O2-

2Mg + O2 2MgO

2. Reaction between sodium and chlorine

Na Na+ + e-

Cl + e- Cl-

2Na 2 Na+ + 2 e-

Cl2 + 2e- 2Cl-

2Na + Cl 2 2NaCl

3. Decomposition of lithium chloride

Li Li + + e-

Cl + e- Cl-

2Li 2li + + 2e-

Cl2 + 2e- 2Cl

2LiCl 2Li + Cl 2

Oxidation state = tell how many electrons each atom of an element has gained/ lost.

0 1 2 3 4 5 6 7 = 0 I II III IV VI VII

Iron can have + II and + III Copper can have +I and +II Manganase can have + VII , +IV and +II Chromium can have + VI and +III

If oxidation states changes during reaction it is a redox reaction: oxidation rise in oxidation state/ reduction fall in oxidation state.

Eg: When ammonia and hydrogen chloride gases mix a white smoke of ammonium chloride forms. The equation and oxidation states are:

NH3 + HCL NH4Cl

-III + I +I –I -III +I –I

No change in oxidation state- not redox

Eg: When chlorine is bubbled through a solution of iron (II) chloride iron (III) chloride is formed. The equation and oxidation states are:

2FeCl2 + Cl2 2FeCl3

+II -I 0 +III -I

- Oxidizing agent brings about the oxidation of another subastance ( and itself is reduced)

- Reducing agent brings about the reduction of another substance ( and itself oxidized)

Oxidising agents

Potassium manganate ( VII)

- Transition metal- Purple- KMnO4- Oxidation state= +VII, much more stable in +II. Driven to gain electrons. - Takes electrons in presesnce of little acid- MnO4- reduction Mn2+

Manganate VII purple manganese II colorless - Can be used as test for reducing agent ( if present purple color fades)

Potassium dichromate VI and oxidizing agent

- Transition metal- Has Oxidation state +VI but more stabe in +III - Reducing agent in presence of acid- Wants to gain electrons and reduce oxidation state to +III- Cr2O7 2- 2Cr3+

Dichromate VI orange Chromium green

Reducing agent

Potassium iodideH2O2 + 2Kl + H2 SO 4 I2 + K2SO4 + 2H2O

Hydrogen potassium iodide iodine potassium

Peroxide sulphate

There’s a colour change

2I- oxidation I2

colourless red-brown

Topic 8. Acids and bases

Acids : turn litmus paper red Ph 7-1 Contain H+ ions Proton donors

Eg: sulphuric acid: H2SO4

Nitric acid: HNO3

Hydrochloric acid: HCL Ethanoic acid : CH3COOH Methanoic acid ( ant stings) Citric acid ( fruits) Carbonic acid ( fizzy drinks)

Strong acids: high conductivity+ high ph in water all molecules dissociate into H+ ionsEg: HCL 100% H+ + Cl-

Weak acids: low conductivity + low ph in water only some become H+ ionsEg: CH3COOH less 100% H+ + CHCOO-

Alkalis : turn litmus paper blue Ph 7-14 Contain OH- ions Proton acceptors Eg: sodium hydroxide: NaOh Potassium hydroxide KOH Calcium hydroxide Ca(OH)2 Ammonia NH3

Strong alkalis : high conductivity+ high ph higher concentration OH- ions

Weak alkalis : low conductivity+ low ph lower concentration of OH- ions eg: ammoniaNH3 + H2O NH4 + + OH –

Acid + metal = metal salt + hydrogenH2SO4 + Mg = MgSO4 + H2

Acid + metal oxide = metal salt + waterH2SO4 + CuO = CuSO4 + H2OAcid + metal carbonate = metal salt+ water + carbon dioxide2HCL + CaCO3 = CaCl2 + H2O + CO2

Acid + Alkali = Salt + water

Alkalis react with ammonium compounds driving ammonia out:Calcium + ammonium calcium + water + ammoniaHydroxide chloride chloride

Ca(OH)2 + 2NH4Cl CaCl2 + 2H2O + 2NH3

IONIC EQUATIONS FOR NEUTRALISATION

1. HCL + NaOH NaCl + H2O

H+ + Cl- + Na+ + OH- Cl- + Na+ + H20

So: H+ + OH- H2O

H+ is just a proton , so acid donates protons to hydroxide ions. Hydroxide ions accept these protons and form water.

NEUTRALISATION OF ACID BY INSOLUBLE BASE:eg: magnesium oxide

2HCL + MgO MgCl2 + H2O

Ionic eqt: 2H+ + O2- H2O

Uses

Cooking: lemon, citric acid + vinegar ethanoic acid. Acid breaks down meat tissues make it less tough.

Cleaning: sodium hydroxide or ammonia attack grease

Treating insect stings: stings are acid neutralize by rubbing calamine lotion ( zinc carbonate) or baking soda ( sodium hydrogen carbonate)

Indigestion: neutralize excess stomach acid with baking soda.

Soil: Acidic soil due to rotting vegetation or too much fertilizer. Use limestone (calcium carbonate) quicklime ( calcium oxide) slaked lime ( calclium hydroxide)

Factory waste: slaked lime

Acid rain: rain naturally acidic: carbon dioxide + water carbonic acid CO2 + H2O H2CO3

Sulphur dioxide/ oxides of nitrogen = sulphuric acid and nitric acid Buidings made of limestone eaten away Metal structures attacked

Trees & plants killed, carried to rivers fish & animals affected. Quicklime added to lakes/ factories spay acidic wastes with slaked lime/

catalytic converters

MAKING SALTS

Starting with metal

Eg: Zn + H2SO4 ZnSO4 + H2

Add zinc to acid. Dissolves when bubbles stop reaction finished. Remove excess zinc by filtering. Heat evaporate some water . Crystalise.

Potassium, sodium, calcium too violent. Silver, gold, copper don´t react.

Starting insoluble base

As copper won´t react use copper oxide insoluble.

Eg: CuO + H2SO4 CuSO4 + H2O

Add copper II oxide to dilute sulphuric acid, dissolves & turns blue. Add more till no more will dissolve & filter to remove excess. Heat remove some water+ crystallize.

Starting soluble base ( alkali)

Dangerous add sodium to acid- use sodium hydroxide.

Eg: NaOH + HCL NaCl + H2O

Use indicator Phenolphthalein to see when reaction complete. Pink in alkaline solution colourless in neutral / acid.

Use titration method to find amount acid needed to neutralize alkali. Heat solution evaporate water. Crystallize.

Insoluble salts:

Silver and lead chloride/ Calcium barium lead sulhate ,/ carbonates except sodium potassium and ammonium carbonate

Make them by precipitation.

Eg: make barium sulphate:

BaCl2 + MgSO4 BaSO4 + MgCl2

Ionic equation:

Ba2+ + SO4 - BaSO4

Mix solutions, white precipitate forms. Filter, rinse place warm oven to dry.

To precipitate an insoluble salt must mix a solution contains its positive ions with one contains negative ions

Precipitation used in photographic film:

Silver nitrate + potassium bromide silver bromide + potassium nitrate

Where light strikes silver bromide breaks to silver, turns grey

Topic 9. Electricity & chemical change.

Electricity: stream of moving electrons

Conductors:

- Solids: metals + graphite- Ionic substances when melted or dissolved water

Electrolysis : decomposition caused by electricity

Electrolyte: liquid conducts current

Electrodes : rods cary current into / out of electrolyte

- Positive= anode- Negative = cathode

1. Electrolysis molten ionic compounds

Potassium iodide potassium + iodine

2Kl 2K + I2

At cathode : 2K+ + 2 e- 2K

At anode : 2I- I2 + 2e-

Zinc chloride zinc + chloride

ZnCl2 Zn + Cl2

At cathode: Zn+2 + 2e- Zn

At anode : 2Cl-2 Cl2 + 2e-

2. Electrolysis solutions

H2O H+ + OH-

At cathode: metal or hydrogen. If metal more reactive stays in solution. If less reactive (copper, silver) accepts electrons metal forms at cathode.

At anode: non-metal, or oxygen. If halide ion in high concentration they give electrons. If no halide of not enough concentration OH- gives electrons, oxygen forms.

Concentrated solutions

Potassium bromide KBr.

At cathode: 2H+ +2e- H2

At anode: 2Br- Br2 + 2e-

Dilute solutions

Copper chloride CuCl2

At cathode : Cu2+ + 2e- Cu

At anode : 4OH- 2 H2O + O2 + 4e-

Hydrochloric acid HCL

At cathode :4 H+ + 4e- 2H2

At anode : 4OH- 2H2O + O2 + 4e-

Electrolysis of brine

Brine= concentrated solution of sodium chloride.

2Nacl + 2H2O 2NaOH + Cl2 + H2

Use diaphragm cell, lets ions through but gases apart. Anode of titanium. Cathode of steel.

At cathode: 2H+ + 2e- H2

At anode: 2Cl- Cl2 + 2e-

Chlorine : poisonous yellow gas. - Platic PVC, dyes, paints, bleaches, medical drugs, pesticides

Sodium Hydroxide: alkaline corrosive- Soaps, detergents, dyes, paper

Hydrogen, colourless flammable gas- Nylon, hydrogen peroxide, harden vegetable oil to make margarine,

fuel hydrogen cells.

Uses of not inert electrodes:

Refining copper

Anode impure copper, cathode pure copper. Electolyte dilute copper sulphate solution.

Copper in anode dissolves, impurities drop to floor. Pure copper builds at cathode.

Pure copper for: wires in cars, cooking utensils..

At anode:

Cu Cu2+ + 2e-

At cathode:

Cu2+ + 2e- Cu

Electroplating

Anode silver, cathode object to electroplate. Electrolyte silver nitrate solution.

At anode silver dissolves into ions, at cathode silver coat the object.

At anode:

Ag Ag+ + e-

At cathode:

Ag+ + e- Ag

Topic 10. How fast are reactions?

Rate: measure of the change that happens in a single unit of time. To find rate measure: amount of reactant used up per unit of time : amount of product produced per unit of time.

Measure using a gas syringe.

Collision theory: particles must collide with each other Collisions must have enough energy to be successful

1. Concentration. Higher= faster, steeper curveMore particles = more chances successful collision

2. Temperature.Higher = faster, steeper curveEnergy = move faster= collide more often with more strength = more successful collisions.

Experiment: dilute HCL + Sodium thiosulphate solution. Yellow precipitate sulphur forms, cross can´t be seen.

3. Surface area.Greater= fasterIn small particles, more atoms exposed and more chances of collisions.

Experiment: HCL + calcium carbonateCaCO3 + 2HCL CaCl2 + H2O + CO2 Large/ small chips.

In flour industry surface area a problem, also flour is combustible. Spark, kit match easily put it on fire.

4. LightSome reactions photochemical.Silver bromide yellow, darkens on wxposure to light as breaks to silver2AgBr 2Ag + Br2

Photosynthesis. 6CO2 + 6H2O C6H12O6 + 6O2

Catalyst: substance speeds up a chemical reaction byt remains chemically unchanged itself. Eg: 2H2 + O2 2H2O Use: platinum wire. Molecules absorbed into platinum wire, less energy needed for bonds to break, called activation energy.

Catalysts usually transition metals or their oxides. Used in industry:Iron- ammonia productionVanadium pentoxide- sulphuric acid

Catalysts in carsPetrols is a hydrocarbon produces:Oxides of nitrogen = acid rainCarbon monoxide = poisonousUnburnt hydrocarbons = cancer

Use catalytic converter:In chamber A harmfull compounds reduced2NO N2 + O2

In chamber B harmful compounds oxidized2CO + O2 2CO2

Enzymes: biological catalysts . Proteins. Work by key& lock model.Work best in conditions of living cell where they belongUsually temperature 22-45ºCIf temp. too high enzyme denatured, looses shapeUsually work best at ph 7.

Use enzymes:- Making ethanos by anaerobic respiration of yeast. Fermentation

C6H12O6 2C2H5OH + 2CO2 + energy

- Making bread also fermentation. CO2 makes dough expand. Heat kills yeast.

- Biological detergents: grease, sweat…

Topic 11. Energy changes and reversible reactions.

Bond energy: energy needed to break bonds or released when these are formed.

Exothermic Reactions

More energy is released in bond making than the energy needed for bond breaking

Endothermic reactions

More energy released in bond breaking than the needed for bond making.

To calculate energy changes:

1. Eg: between Hydrogen and chlorine

H-H + Cl-Cl 2H-Cl

Energy in

1x H-H 436kJ

1x Cl-Cl 242 kJ

Total: 678kJ

Energy out

2x H-Cl 862

Energy in- energy out = 678kJ – 862kJ = -184kJ

Reaction gave out 184kJ. Exothermic reaction.

2. Eg: decomposition of ammonia : 2NH3 N2 + 3H2

Energy in

6x N-H 2345kJ

Energy out

1x N =N 946kJ

3X H-H 1308kJ

Total energy out= 2254 kJ

Energy in- energy out = 2346kJ – 2254 kJ = 92 kJ

So reaction takes in 92kJ of energy. Endothermic reaction.

Activation energy: the energy needed to break bonds

Good fuel: lots of heat given out

Little pollution

Available

Easy store & transport ( not very flammable)

Not very expensive

Fuel+oxygen

Energy plenty energy given out

Oxides

New fuels:

Ethanol: C2H5OH Used in car engines or mixed with petrol

Hydrogen: burns explosively in oxygen- lots of energy. Space rockets. Also in fuel cells to power homes/ cars.

Nuclear fuels: have unstable atoms or radioisotopes which break down. Force them tto break by shooting neutrons at them Give out radiatin and energy used to heat water, steam, drive turbines…

High amount energy + no CO2 or greenhouse gases

Waste can be radioactive, explosion affect large area

Reversible reactions:

Copper sulphate

Forward = CuSO4. 5H2O CUSO4 + 5H2O

Backward = CuSO4 + 5H2O CuSO4.5H2O

CuSO4.5H2O CuSO4 + 5H2O

Blue crystals White powder

Hydrated Anhydrous

Endothermic Exothermic

Thermal decomposition of ammonium chloride

NH4Cl NH3 + HCL

When heated on test tube ammonium chloride decomposes. Ammonia and hydrogen chloride gases rise combine again at end of tube where its cool.

In a closed system a reversible reaction reaches dynamic equilibrium= forward and backward reactions take place at same rate, no overall change.

Manufacture of ammonia

N2 + 3H2 2NH3

It reaches equilibrium, every ammonia molecule formed, another two break into nitrogen and hydrogen, final level of ammonia remains unchanged.

Improve yield by:

High pressure- molecules will oppose by forming more ammonia to reduce number of molecules and space occupied and pressure.

Low temperature- hig temperature will make backward reactin (endothermic) quicker so more ammonia will break down. Also reaction will be slower as less energy provided

Remove ammonia- by cooling and running it off as a liquid. Warm again and nitrogen and hydrogen left will again join into ammonia

Get decent rate by :

Raise temperature: more energy so quicker (compromise)

Use catalyst: iron.

Unit 12. Behaviour of metals

Properties:

Strong, malleable, ductile, sonorous, shiny, good conductors heat/electricity, high density, high melting/boiling point, react with oxygen to form oxides (which are bases) form positive ions.

Non – metals:

Not strong malleable ductile or sonouros, Brittle. Low melting, boiling points. Poor conductors heat/electricity, low densities, react with oxygen to form oxides which aren’t bases ( dissolve in water to give acidic solutions) Form negative ions.

Very reactive metals react with water to form hydrogen and a hydroxide When reacting with water: sodium wizes over the surface

: calcium gives some bubbles of H2 and a cloudy solution of Ca(OH)2

: magnesium react slowly, quickly when heating and glows.

When metal reacts water or HCL, displaces hydrogen if more reactive. Not reactive metals ( silver, copper) don’t displace / react.

Competing for oxygen:

- When a metal is heated with oxide of less reactive metal, acts as reducing agent in an exothermic reaction.

Competing to be compound in solution

- A metal will always displace a less reactive metal from solutions of its compounds

Competing with carbon for oxygen

- If metal less reactive than carbon, carbon will displace it. 2PbO + C 2Pb + CO2

Competing with hydrogen for oxygen

- Hydrogen reduces the oxides which are less reactive than itself:

CuO + H2 Cu + H2O

The more reactive metal the more stable its compounds are and the more difficult to extract from its ores.

Some compounds breakd down on heating = thermal decompositions. On heating: - Carbonates- decompose to oxide & carbon dioxide

- Carbonates of sodium/potassium don´t decompose - Strong heating for calcium carbonate (reversible reaction)

- Less reactive, easier carbonates break down CuCO3 CuO + CO2

Green black

- Hydroxides- decompose to oxide & water

Zn(OH)2 ZnO + H2O - Hydroxides potassium/ sodium don’t decompose - Less reactive , easier breakdown

- Nitrates- all decompose on heating, different products: - Potassium/ sodium nitrates break to nitrites and oxygen 2NaNO3 2NaNO2 + O2

- Other metals break to oxides + nitrogen dioxide (brown) . + oxygen 2 Pb(NO3)2 2 PbO +4 NO2 + O2 - Less reactive, easier break down

Using reactive series: 1. Thermite process

- Repairing railway linesPowdered aluminum and iron oxide put in container over damaged rail. Mixture lit aluminum reduces iron oxide to molten iron. Iron runs into cracks & gaps hardens. Fe2O3 + 2Al 2Fe + Al2O3

- Also used to extract chromium. Ore converted to chromium oxide, heated with aluminum.

Cr2O3 + 2Al Al2O3 + 2Cr

2. Sacrificial protection of ironIron reacts with oxygen and water to form iron oxide (rust)Iron joined to more reactive metal gives up more electrons2 Zn 2 Zn 2+ + 4e-O2 + 2H2O + 4e- 4OH- 2Zn + O2 + 2H2O 2 Zn(OH)2

Zinc oxidized instead of iron, sacrifized.3. Galvanising

Same as sacrificial protection in iron roofs, steel in car bodies. - Iron / steel plated with zinc by electrolysis ( galvanized)- Zinc coatings keeps air and moisture away- If coating broken, zinc will still protect iron by sacrificial protection.

4. Making cells

Less reactive metal is negative pole, atoms give up electrons go into solution as ions. The metals (electrolytes) must be differentThe further the metals are in reactivity series the more voltage

Reactivity series:

PotassiumSodiumCalciumMagnesiumAluminum

CarbonZincIronLead

HydrogenCopperSilverGold

Topic 13. Making use of metals.

Ores: rocks from which metals are obtainedBauxite- ore alluminium1. Mine ore

2. Extract metal; metal compound- reduced to metal by electrolysis or reducing agent.

Carbon can be used as reducing agent- reducing metal oxides (most ores are oxides) Carbon monoxide also used.

For metals above carbon must use electrolysis

Eg:

Carbon

- Iron (III) oxide + Carbon monoxide iron + carbon dioxide- Fe2O3 + 3CO 2Fe + 3CO2

Electrolysis

- Aluminium oxide Aluminium + oxygen- 2Al2O3 4Al + 3O2

Zinc sulphide First roasted in air to give zinc oxide

- Zinc sulphide + oxygen zinc oxide + sulphur dioxide- 2ZnS + 3O2 2ZnO +2 SO2

Then…

- Zinc oxide + Carbon monoxide zinc + carbon dioxide- ZnO + CO Zn + CO2

Zinc must be fractionall distilled to remove impurities

Zinc also made by electrolysis. Zinc oxide very high melting point, so dissolved in dilute sulphuric acid, giving zinc sulphate. Then pure zinc by electrolysis.

Blast furnace: extracting iron from ore.

Mixture (charge) added trough top, hot air blasted in bottom. After reactions liquid iron collects at bottom.

Charge - iron ore, haematite (iron II oxide, Fe2O3 mixed with sand)

- Limestone ( CaCO3)- Coke ( coal, nearly pure carbon)

1. Coke burns, gives off heatBlast hot air starts burning, gives CO2

Carbon + oxygen Carbon dioxideC + O2 CO2

2. Carbon monoxide madeCarbon dioxide reacts with more coke, give COCarbon dioxide + Carbon Carbon monoxideCO 2 + C 2CO

3. Iron III oxide reducedCarbon monoxide + iron III oxide Iron + Carbon dioxide3CO + Fe2O3 2 Fe + 3 CO2

4. Limestone:

5. Reacts with sand (sillica) in ore

Limestone + silica calcium silicate + carbon dioxideCaCO3 + SiO2 CaSi

Topic 16. Non- metals: oxygen, sulphur, chlorine, carbon

Oxygen- Clear, colourless gas, no smell- Slightly soluble in water

- Very reactive, makes oxides

20% of the airMaking in industryFrom air by fractional distillationMaking in labFrom hydrogen peroxide, use manganese oxide as catalystH2O2 2H2O + O2UsesRespiration - C6H12O6 + 6º2 6CO2 + 6H2O + energyBurning of fuels – we burn them to obtain energy. Eg: methaneCH4 + 2O2 CO2 + 2H2O + energyTest: wooden splint

1. Basic oxides. Metals react with oxygen to form basic oxides2. Acidic oxides. Non-metals react with oxygen to form acidic oxides3. Amphotheric oxides. Both acidic and basic. Alluminium & zinc. 4. Neutral oxide. Neither acidic or basic. Carbon monoxide, dinitrogen oxide, N2O

Sulphur- Brittle, yellow solid- Two different forms , allotropes ( rhombic and monoclinic) - Molecular, low melting point- Doesn´t conduct electricity- Insoluble in water- Reacts with metals to form sulphides- Non metal- In large underground beds. In rims volcanoes.- Occurs as compound in many metal ores. Eg Galena ( lead sulphide)- Occurs naturally in fossil fuels.

Extracting sulphur- From oil and gas. Eg in natural gas methane & hydrogen sulphide

2H2S + O2 2S + 2H2O- From sulphur beds. Superheated water pumped down to melt sulphur

and carry to surface.

Uses:

- For sulphuric acid- Added ro rubber for car tyres. Vulcanizing- Make drugs, pesticides, matches, paper.- Added to cement to make sulphur concrete.

Sulphur dioxide:- When sulphur burns in air- Colourless heavier than air , strong choking smell- Acidic oxide, dissolves in water to form sulphurous acid- H2O + SO2 H2SO3

- Acts as bleach when damp/solution, removes colour form compounds= reduces acid rain

Uses :- sulphuric acid- Bleach wool, paper, silk- Sterilizing agent in soft drinks and jam.

Sulphuric acidMaking :

Contact process: raw materials- sulphur, air, water- Sulphur dioxide, ,air, water. -

Sulphur1. Burned in air

Sulphur dioxide, SO2 : S + O2 SO2

2. Mixed with more air3. Passed over four separate beds of catalyst ( pellets vananium oxide) at 450ºC

Sulphur trioxide SO3: 2SO2 + O2 2SO3

4. Dissolved in concentrated sulphuric acid

Thick fuming = oleum5. Mixed with water

Concentrated sulphuric acid H2SO4 H2O + SO3 H2SO4

Step 3 reversible: 4 beds for mre chances to react + sulphur trioxide removed. Step 3 reaction exothermic. 45ºC compromise = for catalyst/ yield To keep temp. down pies cold water to carry away heat Step 2 use acid not water, this would make thick+ dangerous mist of acid.

Uses:

- Fertilizers. Ammonium sulphate- Paints, pigments, dyestuffs- Fibres and plastics- Soaps detergents

Concentrated: thick oily liquid. Dehydrating agent: removes water. Removes water of crystallization from blue copper (II) sulphate crystals giving anhydrous salt.CuSO4.5H2O CuSO4 Blue white

Chlorine- Green- yellow gas- Very reactive- Not found as a free element, occurs as compound rock salt (sodium

chloride)- Heavier than air- Soluble in water chlorine water. Acidic chlrine reacts with water to form

2 acids. C2 + H2O HCL + HOCLHydrochloric hypochlorous Hypochlorous acid decomposes giving acid. 2HOCL 2HCL + O2

- Chlorine water is bleach. Hypochlorous acid loose oxygen to other substance- oxidizes. Many loose colour when oxidizes.

Displacement reaction of chlorineCl2 + 2KBr 2KCl + Br2 Colourless orange

Chlorine acts as an oxidizing agent ( Bromine looses 2 electrons) A more reactive halogen always displaces less reactive halogen from solutions

compounds

Made in industry- Molten sodium chloride- Brine, concentrated solution of sodium chloride

Uses- Plastic PVC- HCL- Solvents- Medical drugs- Sterilize water

Carbon

- Free element diamond & graphite. Diamond hard clear solid. Graphite dark+ greasy solid

Carbon dioxide: ocurs naturally in air and in:- Carbon compounds burning in plenty of air eg methane:

CH4 + 2O2 CO2 + 2H2O

- Respiration- Reaction dilute acids & carbonates

CaCO3 + 2HCL CaCl2 + CO2 + H2O- Colourless- Much heavier tan air- Doesn’t support combustión- Slightly soluble in water forms carbonic acid - H2CO3

- Used in fire extinguishers- Put in soft drinds- Solid for frozen food.- Lime water turns milky, precipitate calcium carbonate

Ca(OH)2 + CO2 CaCO3 + H2OIf more CO2 precipitate dissolves, calcium hydrogen carbonate solubleCaCO3 + CO2 + H2O Ca(HCO3)2

Carbon monoxide: when carbon compounds burn in too little oxygen. Eg: 2 CH4 + 3O2 2CO + 4H2O

- Poisonous, no smell. Moxes haemoglobin. Carbonates: compounds contain carbonate ion: CO3 -2

Eg: calcium carbonate CaCO3

- Insoluble in water ( except sodium, potassium, ammonium carbonates)- React acids to form salt, water and carbon dioxide- Most break on heating to oxide and carbon dioxide.

CaCO3 CaO + CO2

Organic compounds: carbon compounds found in or derived from living things. Thousands: proteins, carbohydrates, fats, incrude oil and coal, plastics and medical drugs,

Limestone Limestone CaCO3 Crushed limestone grind Powdered limestone

heat Used in: Used in:

Extracting iron from ore neutralize acidity in soil, water Make sodium carbonate neutralize SO2 gases power stations Material for road building making glass

Stone chips for concrete

Heat & clay Quicklime, CaO add water Slaked lime Ca(OH)2

Used: Used in;

Making steel from iron Neutralise acidity soil/water due to acid rain Neutralize acidity in soil Drying agent in industry

Cement Used to make concrete

QuiklymeCaCO3 CaO + CO2

Calcium carbonate calcium oxide carbon dioxide

Slaked limeCaO + H2O Ca(OH)2

ConcreteCement mixed with aggregate ( mixture small stone chips and sand) and water. Compounds in cement bond to form hard mass of hydrated crystals. Exothermic.

Topic 17. Organic chemistry

Oil:- Non- renewable- fossil fuel- made from remains dead sea plants/ animals, buried form

layer of sediment turn into oil and gas. Movements in Earth’s crust cause sea floor to raise, water drains, sea floor = land.

- mixture different organic compounds ( started off in living things). Most also hydrocarbons

- used for: transport fuel, burned for heat, chemicals for plastics, paint…- Refining oil = separate oil into compounds with molecules of similar size- Refining in lab:

1. Heat oil, compounds evaporate, ones smaller molecules first2. Hot vapours rise and thermometers reading too. Vapours condense

in test tube.3. When temp. 100 replace first tube with empty one. Liquid in first

tube first fraction from the distillation.4. Collect three more fractions same way replacing tybe at 150ºC,

200ºC, 300ºC

Fraction Boiling point range flows? volatile? burns? Size molecules

1 up to 100º very runny volatile very easily small

2 100ºC to 150ºC runny less volatile easily

3 150ºC to 200ºC not very runny even less volatile not easily

4 200ºC to 300ºC viscous least volatile only with wick large

Refining in industry: fractional distillation:Hot at base, coll at top.Crude oil pumped in at base, compounds start to boil off, smaller molecules boil off.

Gas – C1 to C4 separated into fuels methane/propane/ethane/butanePetrol (gasoline) – C5 to C6 fuel for carsNaphta C6 to C10 starting point / feedstock many chemicals and plasticsParaffin (kerosene) C10 to C15 fuel aircraft, oil stoves, lampsDiesel C15 to C120 fuel for diesel enginesFuel oil C20 to C30 fuel power stations and ships and heating systemsLubricating fraction: C30 to C50 oil for car engines and machinery, waxes and polishesBitumen C50 + : road surfaces, roofs

Fractions must be treated to:Remove inpurities, sulphur compounds, if left in fuel burn to sulphur dioxide gasSome fractions separated further compuds eg: methane, ethanePart fraction cracked, breaking molecules into smaller ones

Cracking in lab: Heat mineral wool soaked with hydrocarbon oil. Use alluminium oxide catalyst & heat.

Thermal decomposition,gas product no colour, strong smell, burns readily, many chemical reactions. Reactant high boiling point, not flammable ( large molecules, long carbon chain)Product low boiling point, volatile (small molecules, short carbon chains)Product hydrocarbon as nothing new added. Catalytic cracking as use of catalyst

- Craking makes best use of oil, eg crack naphta for more petrolDecane(Naphta) pentane, propene, ethaneH H H H H H H H HH-C-C-C-C-C-C-C-C-C-C-HH H H H H H H H H H

H H H H H H H H H HH-C-C-C-C-C-H + C=C-C-H + C=C H H H H H H H H H H

- Always produces short-chain compounds with carbon-carbon double bond= reactive

1. AlkanesMehtane ( CH4) ethane (C2H6) propane (C3H8) butane (C4H10)Homologous series, general formula: CnH2n+n

- Found oil natural gas.- First four gases at rtp, next twelve liquid, rest solids, boiling point

increases with length as force attraction between molecules increases- Each carbon forms 4 single covalent bonds, unreactive.- Burn well supply oxygen form CO2 and water vapour and heat, fuels:

CH4 + 2O2 CO2 + 2H2O - Less oxygen incomplete combustion give CO or carbon only

2CH4 + 3O2 2CO + 4 H2O + less heat energyCH4 + O2 C + 2H2O + even less heat energy

- Alkanes react with chlorine in sunlight, H HH - C - H + Cl2 light H - C - Cl + HCL H HChlorine atom takes place hydrogen = substitution reaction. Can be explosive

Isomers

Same formula, different structures

Branched isomers lower boiling points, branches stop molecules getting close

Branched isomers also less flammable

2. AlkenesDouble bond between carbon atomsCnH2n

Double bond carbon dictates how alkenes react = functional group of alkenes

- Made form alkanes by cracking- Unsaturated because double bonds- More reactive alkanes, double bond can form single bonds:

H H H H C=C +H2O heat + pressure H-C-C-HH H catalyst H H

Hydrogen adds on = addition reaction- Addition reaction with steam to form alcohols OH

H H H H C= C + H2O heat pressure H- C -C - OHH H catalyst H H

- Alkenes highly flammable burn readily in oxygenC2H4 + 3O2 2CO2 + 2H2O

- Alkenes more carbon so more likely incomplete combustion- Alkenes add to each other, long carbon chains = polyners. Reaction=

polymerization

Test unsaturation: + Br2

Bromine dissolved water,bromine water = orange. If unsaturated addition reaction transparent

H H Br Br

C=C Br2 H-C- C- H

H H H H

ethene orange dibromoethane

Also isomers in alkalenes : C4H8

H

H C H

H H H H H H H

C=C-C-C-H H-C-C=C-C-H C = C- C –H

H H H H H H H H H H

But-1-ene but-2-ene 2-methylpropene

3. Alcohols OH General formula CnH2n+1 OHOH group is functional group for alcoholsNames have -1 in it has carbon atom attatched 1 end of chain

Ethanol in alcoholic drinksGood solventEvaporates easily (parfums)Starting point many chemicalsCar fuel

Clear, colourless, boils 72ºCMisceable with waterBurns well in oxygen plenty of heatC2H5OH +3O2 2CO2 + 3H2O + heat

Can be dehydrated to ethane by passing its vapour over heated alluminium oxide (catalyst) H H H HH-C-C-OH -H2O C=C H H Al2O3 H H

Ethanol in air oxidised to ethanoic, sour taste

H H H OH-C-C-OH O H –C-C

H O-HEffect alcohol:Agressive, Little coordination, depression, ulcers, high blood pressure, brain liver damage, cáncer…

Ethanol made by:1. Fermentation yeast ( batch)

C6H12O6 2C2H5OH + 2CO2 + energyexothermic

2. Ethene hydration (continuous) H H H H C=C + H2O catalyst H- C- C – OHH H 570ºC 60-70 atm H H

Exothermic, reversible

3. Carboxylic acids

Formula: CnH2n+1 COOH

Functional group COOH, carboxyl group

Can be made from oxidation of alcohols

Solutions turn litmus red

Like inorganic, solution have H+ ions, gives them acidity. Ethanoic dissolves:

H O H O

H-C-C H-C-C + H+

H OH H O-

Ethanoic acid ethanoate ion hydrogen ion

Carboxyl acids weak acids, only some dissacociate React metals and bases to form salts:

Ethanoic acid + Sodium hydroxide sodium ethanoate + water

CH3COOH + NaOH CH3COONa + H2O

React with alcohols to give compounds called esters H O H H H H O

H-C-C + OH-C-C-C-H H-C-C H H H

H OH H H H H O-C-C-C-H

H H H

Ethanoic acid propanol propyl ethanoate

Condensation reaction= 2 molecules join to larger molecule eith loss one molecule,water

Esters: smell, taste fruit and vegetables

Oils and fats obtained plandt animals mixture esters

Artificial esters made for shampoos, perfums…

Topic 18.Polymers

Polymer: substance contains very large molecules formed when lots small ones join together. Made of macromolecules

Can be made from alkenes:

H H H H H H H H H H C=C C=C C=C C=C C=C

H H H H H H H H H HEthane molecules (monomers)

H H H H H H H H H H H -C-C-C-C-C-C-C-C-C-C-C- H H H H H H H H H H H Polythene

Synthetic polymer: man made : nylon, chewing gum, plastics…

Natural polymers: starch , cellulose (made glucose) Proteins (made amino acids)

ADDITION POLYMERISATION

H H H H

n C=C heat pressure - C – C –

H H catalyst H H n

- Monomers added together- Double bonds break and molecules ad on to each other- Double bond needed.- Long chains created

Monomer Part of polymer molecule eqt for reaction

H H H H H H H H H H H H H H H H H

C=C - C- C-C-C-C-C-C-C-C-C-C- C=C C C

H Cl H H H H H H H H H H H H Cl H Cl n

Chloroethene poly(chloroethene) PVC, polyvinyl chloride

F F F F F F F F F F F F F F F F F

C=C -C-C-C-C-C-C-C-C-C- n C=C -C-C-

F F F F F F F F F F F F F F F F F n

Tetrafluoroethene Poly(fluoroethene)

Teflon

C6H5 H C6H5 H C6H5 H C6H5 H C6H5 H C6H5 H

C=C - C - C – C – C – C – C - n C=C - C - C –

H H H H H H H H H H H H n

Phenylethene Poly(phenylethene)

Styrene poly(styrene)

CONDENSATION POLYMERISATION

Two different monomers

Double bonds break, monomers eliminate small molecules

1. Making nylon

H H H H H H H H H H H O H H H H O

N –C-C-C-C-C-C-C-C-C-N C- C-C-C-C- C

H H H H H H H H H H H Cl H H H H Cl

A- 1,6 diaminohexane B- Hexan- 1,6 dioyl chloride

Functional groups take part in reaction:

This is the macromolecule:

Nylon called polyamide because amide link( CHON)

2. Making Terylene

benzene, 1,4 dicarboxylic acid ethane, 1,2-diol

Now there is an ester link (COO) so Terylene called polyester

Synthetic polymers called plastics:

- Don’t conduct electricity, heat- Unreactive, safe- Usually light- Don’t break easily- Strong because long molecules- Don’t catch fire.

Change properties: more pressure : plastic branches, less dense : more heat: plastic makes long chains, very dense

Uses:Polythene- plastic bags, gloves, bowls, chairs..Polychloroethene- PVC water pipes, electricity cablesPolypropene – ropesPolystrene : fast-food cartons, packaging, insulation..Teflon : coated frying pans , flooring..Nylon: ropes, tents, curtainsTerylene: clothing

Plastics as unreactive can’t get rid of them. Specially polythene. Can recycle by: melting into new objects, melting so chains crack to small molecules, burned energy for electricity. Degradeable polythene: biodegrable: additives (like starch) that bacteria feed on. : photodegradeable: additives break down in sunlight into flakesBio-polymers: grow inside plants, made by bacteria.

Plants make polmers ( starch ) :6CO2 + 6 H20 C6H1206 + 6O2

You eat animals & olants and in digestion break macromolecules.

Carbohydrates

CHO, monosaccharide, simple sugar.Glucose molecule: maltose, dissacharide: starch polysaccharide:

Cellulose: polysaccharide , made from glucose. Makes fibre

ProteinsMade from amino acids, CHON and some Sulphur. Amino acids combining:

They make amide linkage ( CHON) Proteins make : enzymes, collagen, keratin, haemoglobin, hormones. Body needs 20 amino acids to make proteins. 9 are essential taken in food.

FatsContain fats and oils. Are esters ( formed from alcohol and acid)Alcohol is glycerol

Acids are carboxylic, fatty acids. Fats form:

Condensation reaction, eliminate water molecules. Ester linkage COOSome fats for energy, make energy, stored under skin. Unsaturated fats ( carbon-carbon double bond) are better.

Macromolecules broken by hydrolysis: by reaction with water.- Starch broken into glucose for respiration- Proteins broken into amino acids for proteins- Fats and oil broken into glycerol and fatty acids

Eg: hydrolysis of ester in vegetable oil

Enzymes ( amylases, lipases, proteinases) as catalysts,

Hydrolysis in lab:

- Boil in acid

Hydrolysis of esters in fats and oils in industry:

- Using sodium hydroxide = glycerol and sodium salts of fatty acids. Salts used as soaps.