Year 4 Chemistry Notes (compiled by Jing En) Mole Concept Mass – no. of moles x molar mass Volume of gas = no. of moles x molar volume Concentration = mass/volume or no. of moles/volume Percentage yield = actual yield/theoretical yield x 100% Percentage purity = mass of pure substance in sample/mass of sample x 100%

Concentration: g/dm3 or mol/dm3

1dm = 10cm

Intermolecular Forces 1. Van der Waals/instantaneous dipole-induced dipole

- All atoms and molecules experience VDW forces arising from fluctuations in the electron distribution within atoms or molecules - At any instant, the random motion of electrons within an atom or molecule may cause the electrons to be clustered more at one end of the particle, giving that end a very small partial negative charge. - VDW forces exist between ALL molecules, but are the ONLY forces that exist between non-polar molecules.

2. Permanent dipole-permanent dipole

- Attractions between opposite partial charges in the permanent dipoles of polar molecules

3. Hydrogen bonds

- Formed between a hydrogen atom covalently bonded to a highly electronegative atom and the lone pair of a nearby electronegative atom (F, O, N)

Id-id < pd-pd < H-bond < Covalent & Ionic

Melting/boiling points

STRENGTH OF INTERMOLECULAR FORCES OF ATTRACTION.

Both F2 and H20 are simple covalent molecules. There are hydrogen bonds formed between H20 molecules. As F2 is non-polar, only Van der Waals forces exist between the molecules. More energy is required to overcome the relatively stronger H-bonds than weak VDW forces, hence H20 has a higher boiling point than F2.

Qualitative analysis Acids Bases/alkalis Sour, contains H+ ions Bitter, contain oxide or hydroxide ions Important chemical reactions

Acid + Base Salt + Water Acid + Carbonate Salt + Water + Carbon Dioxide

Acid + Reactive metal Salt + Hydrogen Alkali + Ammonium salt Ammonia + Salt + Water

Characteristics of colours of some chemicals

In general, compounds of transition metals have characteristic colours Group I, II, III metals generally form white salts that dissolve to form colourless solutions.

Blue – Hydrated Copper (II) Dark green – Chromium salt Light green - Iron (II), Copper (II)

Colour Inferences Colourless Dilute acids, alkalis and sotluions of salts of Group I, II, III metals White Solid salts of NA+, K+, NH4+, Ca2+, Zn2+, Pb2+, Al3+ Green Copper (II) salts Blue Copper (II) salts + solution (Hydrated) Yellow PbI2, AgI, Iron (III) Chloride, CHLORIDE Dirty green ppt Light green solution

Iron (II)

Reddish Brown Iron (III) Pale pink Manganese (II) salts Purple KMnO4 Solubility of ionic compounds

Soluble Insoluble Salts of NH4+, K+, Na+ Nitrates* Sulfates* Except Pb, Ba, Ca (slightly) Halides – Chlorides*, Bromides, Iodides* Except Pb, Ag Except those of Na+, K+ (Group I), NH4+ Ca is also slightly soluble. So if asked for possible ions present when SOME precipitate dissolved when excess NaOH was added for example, include Ca2+.

Carbonates* Hydroxides Oxides

Test for Cations Metal hydroxides which are amphoteric can react with both acids and bases

Able to dissolve in excess NaOH: Zn, Al, Pb Able to dissolve in excess NH3: Cu (deep blue solution), Zn (colourless)

Test for Anions

Reagents Purpose Observations Dilute Acids CO32- Effervescence, gas observed Add BaCl3 and HCl OR Ba (NO3)2 and HNO3

SO42- CO32-

White ppt formed, insoluble in acid White ppt formed, soluble in acid, Co2 produced

Add AgNO3 and divide ppt into 2 separate portions:

I. Add dilute HNO3 II. Ii. Add aq NH3

Aq AgNO3 and dilute HNO3

Cl- I- Br- CO32-

White ppt formed, insoluble in acid, soluble in aq NH3 White ppt formed, insoluble in acid, insoluble in aqNH3 Cream ppt formed, insoluble in acid, partially soluble in aqNH3 White ppt formed, turned pale yellow, soluble in dilute acid,

CO2 produced Aq Pb(NO3)2 Cl-

I- SO42- CO32-

White ppt formed, soluble in boiling water but reappeared as white crystals on cooling Yellow ppt formed. Soluble in boiling water but reappeared as yellow crystals on cooling. White ppt formed White ppt formed

Aq NaOH, Devarda’s alloy or zinc or aluminium. Warm.

NO3- Effervescence. NH3 evolved.

Test for Gases

Add acid before testing with silver nitrate for chloride ion to eliminate CO3 from the mixture.

Gas Observations Ammonia Colourless and pungent gas evolved. Turns moist red litmus paper blue. Chlorine Pale-yellowish green and pungent gas evolved. Turns moist red litmus red, then

bleached. Moist red litmus was bleached. Carbon Dioxide Colourless and odourless gas evolved. White ppt formed when gas is bubbled

into limewater,soluble in excess gas, forming a colourless solution. Oxygen Colourless and odourless gas evolved. Glowing split rekindled. Hydrogen Colourless and odourless gas evolved. Lighted splint was extinguished with a

‘pop’ sound.

Thermal Dissociation Loss of water of crystallization

Thermal Decomposition (Irreversible) Always anion that is broken down.

Metal Heat on Oxide Heat on Hydroxide Heat on carbonate Heat on nitrate Potassium Stable to heat Stable to heat Stable to heat Decomposed to the

metal nitrite NO2−

and oxygen Sodium

Calcium Decomposed to the metal oxide and water by heat

Decomposed to the metal oxide and carbon dioxide by heat

Decomposed to the metal oxide, nitrogen dioxide and oxygen by heat

Magnesium Aluminium Zinc Iron Tin Lead Copper Mercury Decomposed to the

metal Hydroxides do not exist

Decomposed to the metal, carbon dioxide and oxygen by heat

Decomposed to the metal, nitrogen dioxide and oxygen by heat

Silver Gold

Don’t evaporate to dryness because there may still be impurities left on it.

Heat to obtain saturated solution and cool to obtain crystals. Filter, wash the crystals with distilled water and dry them between filter papers.

Redox reactions

Reduction Oxidation Loss of oxygen Gaining of oxygen Gaining of hydrogen Loss of hydrogen Gain of electrons Loss of electrons Decrease in oxidation state Increase in oxidation state

A spectator ion is an ion that is present in the reactants and the products in a chemical equation.

EXCEPTIONS:

Element Oxidation state Exceptions F -1 (always) O -2 +1 (when bonded to F)

F is more electronegative than O hence the electrons will move towards it, hence having a negative OS while O has a positive OS +1 (when in peroxide/H2O2)

Cl -1 +1 (when bonded to F) Positive OS (when bonded to O)

H +1 -1 (when bonded to metals as hydrides, H-, e.g. NaH) How to answer: Carbon monoxide is oxidised as the oxidation state of carbon increases from +2 in CO to +4 in CO2. Iron (III) is reduced as the oxidation state of iron decreases from +3 in Fe2O3 to – in Fe.

Substance Formula Observation during reaction Use Oxidising agents (gets reduced)

Acidified potassium manganate (VII) (aq)

KMnO4 Purple solution turns colourless. MnO4- (aq) + 8H+ (aq) + 5e- Mn2+ (aq) + H2O (l) Purple Colourless

Test for reducing agents

Acidified potassium dichromate

K2Cr2O7 Not for SPA

Orange turns green Cr2O72- + 14H+ + 6e 2Cr3+ + 7H2O

Test for reducing agents

Hydrogen Peroxide

H2O2 H2O2 + 2H + 2e 2H2O (If reacts with KI, observation comes from Iodine)

Chlorine (Halogens tend to gain electrons, oxidizing agent)

Cl2 Solution turns darker in colour when reacted with bromide or iodide.

Oxidises Br- to Br2 and I- to I2

Substance Formula Observation during reaction Use Reducing agents (gets oxidised)

Potassium Iodide (aq)

KI Solution turns brown. 2I- (aq) + Cl2 (aq) 2Cl- (aq) + I2 (aq) Pale yellow brown

Test for oxidizing agents

Reactive Metals

e.g. Al, Zn

Less reactive metal is produced Displacement of less reactive metals

Hydrogen H2 Reddish-brown solid (Cu) is formed Reduces copper (II) oxide to copper

Disproportionation Reactions: Oxidised and reduced to different OS simultaneously

2H2O2 (l) 2H2O (l) + O2 (g)

Electrochemistry Anode is where oxidation takes place. Cathode is where reduction takes place.

In a spontaneous redox reaction, electrons are transferred and energy is released. Potential energy is generated due to the potential difference between the 2 electrodes which is then converted into electrical energy. The further apart the 2 metals used as electrodes are in the reactivity series, the higher the voltage produced.

Electrolytic cells

Poison sausages can make a zebra ill

therefore let highly clever man slaughter

good pigs.

Electrical energy is used to cause a non-spontaneous reaction.

During the electrolysis of water, a dilute acid (H2SO4 or HNO3) is added because water is a very poor conductor of electricity.

Molten Salts

Ions are always LIQUID in molten state.

Aqueous solutions

For example, the electrolysis of aqueous sodium chloride

Cathode (reduction) Anode (oxidation) Na+ + e- Na 2H+ + 2e- H2 Possible products are sodium metal or hydrogen gas.

2Cl- Cl2 + 2e- 4OH- O2 + 2H2O + 4e- Possible products are chlorine gas or oxygen gas.

How to tell which is preferentially discharged (selective discharge):

1. The position of the metal in the reactivity series Less reactive a metal, the easier its cation is reduced.

2. The nature of the electrode Inert – will not participate in electrolysis but merely connects circuit Metal electrodes at the anode – oxidized and dissolved as ions, electrode participation. E.g. Cu can be oxidized to u2+.

3. Concentration of the anion in the electrolyte The ion present in greater concentration will be preferentially discharged, even if its higher in the series. NO3- and SO42- ions are never oxidized during electrolysis.

Possible types of electrolytic cell:

With inert electrodes (graphite/platinum) - Molten electrolyte (discharge of ions)

- Aqueous electrolyte (discharge of ions) - Aqueous electrolyte with high concentration of halide ions

With reactive metal electrodes - Molten electrolyte - Aqueous electrolyte - Aqueous electrolyte with high concentration of halide ions

Universal indicator Acid: Red. Alkali: Purple. Neutral: Green.

Inert electrodes: Carbon, platinum, gold

Answering Questions

Describe an observation/change. - Describe the appearance of the deposit e.g. colour: shiny, grey - Do not name the metal itself

Effervescence at anode in Cell 1 where hydroxide ions are preferentially discharged to form oxygen gas as OH- is lower than SO42- in the electrochemical series but not Cell 2 where copper is oxidized to CU2+ ions at the anode. The blue electrolyte in Cell 1 (inert carbon electrodes) will be of lighter blue than Cell 2 (copper electrodes) as the CU2+ which is discharged forms back copper.

Salt bridge maintains electrical neutrality so that there is no build up of positive charges at the oxidation half cell and negative charges at the reduction half cell.

No visible reaction observed. Copper: A brown/pink solid is formed. Blue colour solution turns pale blue/colorless. Effervescence observed.

Metal becomes smaller/decrease in mass.

Energetics

Enthalpy is a measure of the total energy of a thermodynamic system. The total enthalpy, H, of a system cannot be measured directly. Thus, change in enthalpy, ΔH is more useful.

- Usually measure temperature of surroundings

Exothermic – ΔH is negative (Chemicals have lost energy to surroundings) Combustion of fuels Respiration

ΔH bond breaking < ΔH bond making

Endothermic - ΔH is positive (Chemicals have gained energy from surroundings) Photosynthesis Photodecomposition of silver bromide Thermal decomposition of limestone

ΔH bond breaking > ΔH bond making

ΔH = total energy for bond breaking – total energy for bond-making = Hp (energy content of products) – Hr (energy content of reactants)

Energy profile diagram

Exothermic reaction Endothermic reaction

Energy level diagram

Exothermic reaction Endothermic reaction

CHEMICAL EQUILIBRIUM (refer to handwritten notes)

Chemical Kinetics Rate of reaction (always positive unless rate at which amount of reactants change) is defined either as increase in the concentration of one of the products per unit time or as the decrease in the concentration of one of the products per unit time. Formula = Amount of product formed or used up/time taken

Methods to measure

If reactants produce gas (in cm3)

1. Collect the gas using a graduated gas syringe and measure the volume of gas produced at regular time intervals.

Why is a magnesium strip placed in the small test tube? It is to control when the reaction starts, ensure that the reaction does not take place until the apparatus is properly set up and that all the gas produced is measured accurately in the syringe.

How does rate of reaction change over time? At t=0, the rate of reaction is the greatest because the gradient of graph is the steepest. At t=0 to t=45, the gradient is less steep and rate also decreases. At t=45, the gradient and rate is both zero. Reaction is completed.

2. Allow the gas to escape and monitor the mass/loss in mass (of the reaction flask and its contents) at regular time intervals

What is the ball of cotton wool at the mouth of the flask for? As effervescence occurs, some solution may splash out of the glass hence cotton wool prevents acid from splattering out during reaction and any loss of mass should only be caused by loss of the gas and not the acid.

(reach 0 eventually)

Average rate is the rate measured over a long time interval. Instantaneous rate is the rate of reaction at any instant. (calculate gradient)

Average speed = total volume of gas/total time

If the reactants are in aqueous solutions and one of the products is insoluble in water The formation of a precipitate could be measured against time by monitoring the opacity of the mixture. - Since the amount of ions in the solution is decreasing during the reaction (due to the formation of insoluble precipitate) the concentration of ions in the solution can be monitored using a conductivity meter.

If reactants are a different colour from the products, then the change in colour intensity could be measured against time A colourimeter can be used to measure the intensity of the blue colour over time.

If the reaction involves an acid or alkali, then a change in pH could be measured against time. - A pH meter could be used to monitor the change in pH over time. - Samples of the mixture could be removed at regular intervals until the reaction stopped and the amount of acid left in the mixture is analysed by titration. (Add cold water to slow down the reaction/another reagent to take away one reagent)

If the reaction is exothermic or endothermic, then a change in temperature could be measured against time The temperature of the reaction could be measured at regular time intervals.

Collision Theory

For a reaction to occur, reacting particles must: - Collide with each other - Possess a certain minimum amount of energy (KE) to overcome the activation energy - Steric factor: Must collide in the correct orientation/geometrical alignment, must bring the reactive parts of the molecule into contact in the correct way (esp. large organic molecules e.g. carboxyl groups) Effective collision

Factors affecting reaction rates:

Mainly affect the collision rate

Concentration of solutions (aqueous)/pressure of gases

- Higher concentration/High pressure means more particles per unit volume - The reacting particles collide with one another more frequently - Frequency of effective collisions increases and hence rate of reaction increases

- As reactants get used up, concentration decreases (the products interfere and get int hew ay of reactants too) Rate of reaction gets slower

Some exothermic reactions do initially speed up if heat is given out more than compensates for decrease in concentration.

Surface area or particle size (when one or more reactants are a solid)

- For solid, only particles on surface can come into contact with surrounding reactant - If powdered, the surface area increases with more surface area for reacting particles to collide - The reacting particles collide with one another more frequently - Therefore, increases the frequency of successful collisions and hence the rate.

Instantaneous reactions/explosions may occur

Mainly affect the proportion with required activation energy

Temperature

As the temperature increases, the particles gain more kinetic energy and move faster, resulting in more frequent collisions. At the same time (more importantly), more particles will possess the necessary activation energy to react upon collision. This causes an increase in the frequency of successful collisions, thus increasing the rate.

- Particles move faster, more Kinetic Energy - More frequent collisions between reactant particles - More colliding particles possess necessary activation energy to bring about a greater frequency of

successful collisions

Effect of a catalyst

Catalysts provide an alternative pathway with lower activation energy, increasing the chances of successful collisions for reactions to react as more particles possess equal to or higher than the activation energy, rate of reaction hence increases with the presence of catalysts.

- Increase rate of chemical reaction without being chemically changed at the end of the reaction - Bring reactive parts of reactant particles into close contact - Alternative pathway for reaction by lowering the activation energy - More reactants will possess this low activate energy so rate increases

The greater the number of reactant molecules with energy greater than the activation energy of the reaction, the faster the reaction. (necessary discussion when temperature + catalyst)

Organic Chemistry Catenation: Ability of an element to form covalent bonds with itself, forming rings, chains or cage molecules Carbon is a tetravalent atom that has 4 valence electrons available for chemical bonding. Homologous series: Group of compounds with the same general formula, similar chemical properties and show a graduation in physical properties as a result of increase in size/mass e.g. melting, boiling points, viscosity, flammability and solubility. Functional group: Group of atoms that give a molecule its characteristic chemical properties.

Hydrocarbons are organic compounds that contain only hydrogen and carbon atoms. C-C bonds are stronger, leading to greater stability in compounds containing C-C bonds.

Displayed/full structural formula: Shows all bonds joining atoms Condensed structural formula: Shows how atoms are joined together but without showing bonds. Repeated units

in straight chain ca be shortened to e.g. (CH2)n in brackets to represent branching. Skeletal formula – Shows bonds, omitting all the C and H joined to C

Heterocyclic – Cyclic compounds in which one or more of the ring atoms is not carbon E.g. heterocyclic ring system

Naming organic compounds

PREFIX-PARENT-SUFFIX

1. Parent: longest continuous carbon chain – number of carbon atoms

Number of carbon atoms

Parent Number of carbon atoms

Parent

1 Meth- 6 Hex- 2 Eth- 7 Hept- 3 Prop- 8 Oct- 4 But- 9 Non- 5 Pent- 10 Dec-

2. Prefix: Location of functional groups e.g. methyl (alkyl groups are alkanes with one hydrogen removed to bond with other atoms)

3. Suffix: Functional group of the highest priority is indicated by the suffix while others are indicated by their prefixes

Group Family of compound Structure Prefix Suffix A Carboxylic acid

carboxy- -oic acid

Alcohol

hydroxy- -ol

B Alkane

- -ane

Alkene

- -ene

C Alkyl Alkane with 1 H removed alkyl- - Halogen F, Cl, Br, I fluoro-, chloro-,

bromo-, iodo- -

Steps in naming a compound

1. Name the parent hydrocarbon. Find longest continuous carbon chain. a. Add prefix cyclo- if most carbon atoms are joined in ring structure

2. Number the carbon atoms such that the carbon attached to the functional group is assigned the lower number of beginning with end nearest to a substituent

3. If carbon chain includes multiple bonds, replace “ane” with “ene” Designate position with multiple bonds with the number of the first carbon of the multiple bond.

4. If molecule includes Group A functional groups, replace last “e” with suffix of highest priority functional number + position number. (Position of –COOH group is always on first C atom hence do not need to state)

5. Indicate all group C substituents and group A functional groups of lower priority with a prefix with appropriate position numbers + alphabetical order before root name.

6. Numbers are used to indicate position of any substituent. Each substituent gets a unique number (e.g. 3, 3)

Commas separate numbers. Dash separate numbers from letters e.g. 3, 3-dimethylpropane.

2, 2, 2-trichloromethanoic acid 4-bromo-1,1,1-trifluorobutane 3,4-dimethylpentanoic acid

Things to take note of

- For cycloalkene, the double-bonded must be counted one after another. - Family takes priority and should have a smaller number e.g. 4-methylcyclopent-1-ene not 1-

methylcyclopent-5-ene (anyway this goes against the first rule) - Alphabetical order for multiple functional groups - Add up the sum of all the numbers – smallest sum is to be used - Carboxyl group is always numbered the first C atom

Isomers

Structural isomers: Same molecular formula, different arrangement of atoms

Skeletal isomers: Different carbon skeletons (skeletal isomers) straight/branched-chain The longer the alkane chain, the higher the number of isomers

Positional isomers: The functional group attached to different carbon atoms e.g. different positions

Functional isomers: Contain different functional groups e.g. belong to different homologous series

Stereoisomers: Same molecular and structural formulae, but different arrangement of atoms in space Geometric isomerism arises from (triangular arraignment of 2 C=C atoms):

- Unsaturated organic compound (alkenes): restricted rotation about C=C bonds - 2 different groups on left-hand end of the bond and 2 different groups on right-hand end

Tips: Draw straight chain, minus 1 carbon and move the methyl group around

Saturated organic compound (alkanes): free rotation about C-C bonds

Isomer with both atoms on: same side – cis isomer, opposite sides: trans isomer

Types of Organic Reactions

Addition reaction: 2 reactants add together to form 1 product (Unsaturated Saturated) Elimination reaction: Single reactant splits into 2 products (Saturated Unsaturated) Substitution reaction: 2 reactants exchange parts to form 2 new products

Alkanes (CnH2n+2)

Saturated hydrocarbons (contain only single C-C and C-H covalent bonds) Each carbon atom bonded to the maximum possible number of hydrogen atoms. Simple covalent molecules held together by weak intermolecular forces (id-id interactions)

Physical Properties

1 to 4 carbon atoms: gases 5 to 16: liquids 17 or more carbon atoms: solids

Boiling points:

Alkanes are simple covalent molecules held together by weak instantaneous dipole-induced dipole interactions, intermolecular forces. As the molecular size increases, the electron cloud size also increases. The strength of the intermolecular forces increases and requires more energy to overcome resulting in a higher boiling point.

Boiling and melting points increase down the group.

Branched-chain alkanes are more spherical in shape than straight-chain alkanes. Less surface area of contact leads to weaker intermolecular interactions. Less energy is required to overcome these forces and branched chain alkane has a lower boiling point.

Solubility in water: Alkanes are not soluble in water. The difference in electronegativity value between carbon and hydrogen is very small and hence, alkanes are not polar molecules. Water, a polar solvent, can only dissolve polar molecules.

Shorter chain alkanes are preferred as fuels due to their high flammability. Hence, the extent of complete combustion is higher.

Density increases down the group: Increase in mass per unit volume due to increase in molecular size and decreasing volume due to changes in state.

Viscosity (resistance to flow) increases down the group: Intermolecular forces increase as molecular size increase, making it harder for bigger molecules to break apart from one another.

Flammability decreases down the group: Greater number of covalent bonds in the molecule so more energy (activation energy) is required to break those bonds.

Reactions of alkanes: - Quite unreactive due to strong C-C and C-H bonds and lack of polarity. - Chemically inert

Combustion

- Complete combustion of any hydrocarbon in excess air or oxygen produces carbon dioxide and water vapour

- If insufficient oxygen, combustion will be incomplete producing carbon monoxide and carbon (soot) - Alkanes are very exothermic hence making them good fuels

Cracking (breaking down large hydrocarbon molecules into smaller and more useful molecules)

- Produces short-chain alkanes and alkenes o Short-chain alkanes are more useful, in higher demand + alkenes can be used in organic synthesis

- Thermal cracking: Heating large alkanes (long chains e.g. bitumen) at high temperatures and pressures without catalyst such that their molecules vibrate strongly enough to break and form mixtures of smaller molecules

- Catalytic cracking: Using a catalyst (e.g. Al2O3/SiO2) at lower temperatures and pressures - Fractional distillation can be used to separate petroleum into useful fractions.

o Petroleum is heated to a high temperature to vapourise the compounds present completely. o The vapour is allowed to pass through the fractionating column, compounds condense when the

temperature is lower than their boiling points. o The boiling points of the compounds increase down the coumn, corresponding to an increase in

the size of the molecules of the compounds.

Substitution Reactions

- Conditions: In the presence of ultraviolet (UV) light

o Provides energy to break down Cl2 to Cl atoms - Alkanes react with halogens to form mixture of halogenoalkanes - Usually hydrogen atom is replaced by e.g. chlorine atom

Why do alkanes not react with F2 and I2 under UV light? UV light does not provide sufficient energy to break F-F bonds. Reaction for alkanes and iodine has a positive change in temperature and is energetically nto feasible. Need to take in great amount of heat.

Alkenes (CnH2n) MUST WRITE CIS OR TRANS IF C=C BOND IS IN THE MIDDLE.

Unsaturated hydrocarbon (Carbon-carbon double bond, C=C) Has the same general formula as CYCLOALKANES. Hence to distinguise between cycloalkane and alkene Bromine test

Physical Properties

- Insoluble in water and soluble in organic solvents - The first few members of straight chain alkenes are gases at room temperature - Boiling point increases with increasing carbon number

o A few degrees lower than corresponding alkane. o Each alkene has 2 fewer electrons than an alkane with same number of carbon atoms.

- Branching lowers boiling point

Different in boiling and meltping points of cis- and trans- isomers (trans > cis)

Boiling point: Difference in shape of molecule or overall dipole moment. There is a net dipole moment in cis isomers hence it is polar and forms strong permanent dipole-permanent dipole interactions.

Melting point: Trans isomers pack better than cis isomers. The poorer packing in the cis isomers means that the intermolecular forces are not as strong as they should be. Less energy is needed to melt the molecule resulting in lower melting point for cis isomers.

Chemical Reactions

Presence of C=C bonds make alkenes more reactive than alkanes.

Addition reaction

- Reactant is added to the 2 C atoms that form the C=C bond, converting the C=C bond to a C-C single bond.

Hydrogenation

- Conditions: 200 °C + Presence of Nickel catalyst + Hydrogen - Addition of hydrogen to alkenes to form alkanes

Changing unsaturated fats into saturated fats (butter to magarine) More convenient because margarine will not melt so easily. Saturated fats have higher melting point harder to melt so it is easier to pack and sell

Bromination: Addition of aqueous bromine to alkenes to form dibromoalkanes

- Conditions: Aqueous bromine - Alkenes decolourise reddish brown (concentrated) or yellow (dilute) bromine solution rapidly. - Chemical test for presence of unsaturated hydrocarbons (e.g. alkenes)

Hydration: Addition of steam to alkenes to form alcohols

- Conditions: 300°C + Pressure of 60 atm + Presence of catalyst, phosphoric (V) acid, H3PO4 - Ethene reacts with steam to form ethanol

Why does the combustion of ethane release a higher content of soot and carbon monoxide than the combustion of ethane? ethane has a lower flammability compared to ethane due to the greater amount of energy required to break C=C bond. Hence, alkanes are preferred as fuels. Alkenes give more pollution problems.

AlcoholS (CnH2n+1OH)

Contain hydroxyl (-OH) functional group -OH group is covalently bonded to carbon atom and does not dissociate in water. Hence, alcohols are NOT alkalis.

Physical Properties

Alcohols have much higher melting and boiling points than corresponding alkanes

- Alcohol molecules are held by hydrogen bonds in addition to id-id interactions. More energy is required to overcome the intermolecular forces between alcohols.

Solubility: Decreases with increasing number of carbon atoms because the longer the carbon chain, the more extensive the instantaneous dipole – induced dipole intermolecular forces. However, alcohol molecules are polar and can dissolve in polar solvent.

Boiling point increases as number of carbon increases

Higher densities compared to hydrocarbons with the same number of carbon atoms - Higher Mr - Greater number of molecules per unit volume due to difference in state

Alcohols have lower flammability Greater energy is required to break a greater number of bonds.

Manufacturing ethanol

1) Hydration of ethane (Addition of steam) Can be used to make other alcohols 2) Fermentation of carbohydrates (However, the yield of ethanol is low because at ethanol levels more or

equals to 20%, yeast is killed) a. Fermentation is a chemical process in which microorganisms such as yeast act on carbohydrates

e.g. starch and sugars to produce ethanol and carbon dioxide. b. Yeast contains enzymes (biological catalysts) that cause starch and sugars to break down c. Conditions: Yeast, 37°C, absence of oxygen gas d. Glucose Ethanol + carbon dioxide

3) Ethanol can be oxidized by atmospheric oxygen to form ethanoic acid. a. However, reaction takes a long time.

Ethanol is used as a solvent, fuel and constituent of alcoholic beverages.

Chemical Reactions

Combustion

- All alcohols readily burn in excess oxygen to form carbon dioxide and water - In some countries, ethanol is combined with gasoline (petrol) to produce a fuel for cars called gasohol - Alcohols produce less energy per unit mass as compared to alkanes and are less energy efficient

Oxidation (primary alcohols only)

- Alcohols (e.g. ethanol) can be oxidized to carboxylic acids by heating it with oxidizing agents o Acidified potassium dichromate (VI) solution, K2Cr2O7, that will be reduced to chromium (III)

ions, Cr2+. The colour change will be from orange to green. o Acidified potassium manganite (VII) solution, KMnO4, that will be reduced to Mn2+ ions. The

colour change will be from purple to colourless. - Oxidation of ethanol by atmospheric oxygen can also give carboxylic acids.

CH3CH2OH + O2 CH3CO2H + H2O

CH3CH2OH + 2[O] CH3CO2H + H2O Oxygen atoms donated by KMnO4

Carboxylic acid (CnH2n+1OH)

Contains the carboxyl (-COOH) functional group

Physical Properties

At room temperature and pressure, short chain carboxylic acids are colourless liquids that are soluble in water.

Boiling point increases

Why are the boiling points of carboxylic acids much higher than alkanes of similar length?

Molecules of carboxylic acids are held by hydrogen bonds in addition to id-id interactions. More energy is required to overcome these intermolecular forces resulting in carboxylic acids having higher boiling points than alkanes.

Why are the carboxylic acids soluble in water?

Presence of very electronegative atoms, polar bonds molecular symmetry that determines if there is a net dipole. Therefore carboxylic acid molecules are polar and like dissolves like, hence it can dissolve in a polar solvent.

Chemical Reactions

Carboxylic acids are weak acids that dissociate partially to form hydrogen ions. (always aqueous)

Acid + reactive metal Salt + Hydrogen gas Salt formed can be named by replacing ‘oic acid’ with ‘-oate’ (e.g. sodium ethanoate)

Acid + carbonate Salt + Carbon dioxide gas + Water

Acid + base Salt + Water

How to test for carboxylic group? Add sodium carbonate to the carboxylic acid. Effervescence will be observed and when gas produced is bubbled through limewater, white precipitate will be formed.

Esterfication

- Reacts with alcohols in the presence of strong acids to produce esters in a reaction called esterification - Condensation reaction as water is formed as a side product - First part of name (ethyl): Alcohol used is ethanol, second part (ethanoate) tells us that the acid used is

ethanoic acid Alcohol first then carboxylic acid

In terms of bond strength: Ester (permanent dipole – permanent dipole) < Alcohol (hydrogen bonds) < Acid (hydrogen bonds) O is very electronegative.

Boiling point depends on the type of intermolecular forces Non-polar: Instantaneous dipole - induced dipole Polar: Permanent dipole – permanent dipole F/O/N and H: Hydrogen bonds

Solubility depends on polar molecules/not Like dissolves like. Polar molecules dissolve in water.

Polymers: Large molecule of very high relative molecular mass (macromolecule) that joins many small repeating units known as monomers together by covalent bonds - POLYMERISATION

Natural polymers Synthetic polymers Proteins Poly(ethene) e.g. plastic bags, clingfirm, toys Carbohydrates Nylon Fats Terylene Polyvinyl chloride (PVC)

Addition polymer: Successive linking together of unsaturated monomers

E.g. Poly(ethene) - formed from ethane molecules at high temperature and pressure in the presence of a catalyst

1. One bond in each double bond breaks 2. Each monomer forms single bonds with 2 other monomers

Name of monomer

Structural formula

Name of polymer Structural formula Uses

Chloroethene (vinyl chloride)

Poly(chloroethene) Polyvinyl chloride or PVC

Waterproof and insulating material, records

Propene

Poly(propene) Plastic bottles, containers, etc.

Phenylethene (styrene)

Poly(phenylethene) Packaging, ceiling tiles

Tetrafluoroethene

Poly(tetrafluoroethene) Non-stick saucepans, bridge, bearings

Condensation Polymers: Links together 2 types of monomer molecules, each containing 2 functional groups, with the elimination of a small molecule such as water

Amide linkage (Synthetic polyamide) Ester linkage (Synthetic polyester) Diagram

Type of molecules

Dicarboxylic acid and diamine Dicarboxylic acid and diol

Property Strong, resistant to attack by chemicals, fungus Strong, resistant to attack by chemicals and fungus, does not crumple easily

Uses - Make fabrics and garments e.g. socks and stockings

- Strong ropes

- Make fabrics for manufacture of clothing - Polyester is not as absorbent and cool to

wear compared to cotton so it is often combined with cotton to produce polyester cotton fabrics

- Make audio and video tapes and plastic soft bottles

Example Nylon Terylene By-products

Water molecules!

To find out the monomers, break the weaker C-O bonds and “add water” back.

Man-made fibres like these are used in clothing, curtain, fishing line, parachutes, and sleeping bags

- Carbohydrates, proteins and fats are natural macromolecules - Proteins have the same amide linkages as nylon but with different monomer units - Fats are esters possessing the same linkages as Terylene but with different monomer units

Plastic and Pollution

Generally describe many synthetic polymers

Plastics are preferred to natural materials because:

- Relatively cheap - Easily moulded into various shapes - Light - Durable (resistant to corrosion and chemical attack)

Threats to the environment:

- Plastics are non-biodegradable and cannot be broken down by bacteria in the soil. o When disposed can accumulate and cause land pollution

- Burning of plastics may lead to release of toxic gases, causing air pollution - Disposal of plastics results in pollution problems.

Application of energy from chemicals (fuels)

Suggest the advantages of using ethanol over propane and over octane. - Ethanol is a liquid and propane is a gas. - Using ethanol over propane gives greater ease of storage and transport (Need to store the gas under high pressure Liquified propane gas LPG) - Octane & Propane are components of fossil fuels which are non-renewable while ethanol is obtained from plants, which are renewable (e.g. sugar cane) - Burns cleanly in oxygen gas. Less sooty compared to petrol.

Suggest the disadvantages of using ethanol over propane and over octane - More ethanol is required to produce the same amount of heat, compared to octane and propane - Lower amount of energy per mole for ethanol

How to answer questions:

Briefly describe a chemical test that can be used to distinguish. Pass the 2 gases separately to aqueous bromine in the absence of UV light, at room temperature. For C4H10, aqueous bromine remains brown while aqueous bromine decolourises when added with C4H8. (bromination is the name of the reaction not the test)

Qualitative Analysis: A green precipitate will be formed. On standing, the green precipitate turns reddish brown. Fe9OH)2 is oxidised by air, oxygen to form Fe(OH)3.

Suggest why it is difficult to investigate the rate of reaction in this experiment using powdered calcium carbonate instead of a lump. The reaction occurs too rapidly and It becomes difficult to study the gradient of the graph.

Suggest why at the start of each experiment, the rate was very slow. The carbonates have been covered with dust/impurities on the surface.

The rate of decomposition of a metal carbonate is inversely proportional to the reactivity of the metal. If reactive (prefer to exist as ions) Harder to decompose

Low pH is acidic, high pH is basic.