Chemistry AS ISA- Re-sit

'The aim of the task is to investigate an enthalpy change'

3.2.1- Energetics

3.2.5- Group 7

3.2.6- Group 2

3.2.10- Alcohols


1. Energetics

Enthalpy Changes (delta H)- The heat energy transferred in a reaction at constant temperature

  • Units= KJ mol -1
  • Standard states and pressure (100KPa/1 atm, 298K/25)= o
  • Exothermic- The reaction gives out energy, delta H is negative
  • Endothermic- The reaction absorbs energy, delta H is positive
  • Bond Breaking- Energy is needed for bond breaking so the process is endothermic
  • Bond Making- Energy is released when bonds are made so the process is exothermic
  • Enthalpy change is the overall effect of these two changes e.g. if more energy is needed to break the bonds than make them, then the reaction is endothermic
  • Standard enthalpy change of reaction- The enthalpy change when a reaction occurs in the molar quantities shown in the chemical equation, under standard conditions in their standard states
  • Standard enthalpy change of formation- The enthalpy change when one mole of a compound is formed fomis constituent elements in their standard states and under standard conditions
  • Standard enthalpy change of combustion- The enthalpy change when one mole of a substance is completely burned in oxygen under standard conditions
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2. Energetics

Mean Bond Enthalpies

  • The enthalpy change for a particular bond varies depending on the molecule it is found and its position in that molecule
  • The mean bond enthalpy is an average for a big range of molecules so it isn't exact

Calculating Enthalpy Changes from Mean Bond Enthalpies

  • Enthalpy change of reaction= Total energy absorbed - Total energy released  
  • Add up the enthalpies of the reactant bonds (endothermic)
  • Add up the enthalpies of the product bonds (exothermic)
  • Take the total of the product bonds away from the reactant bonds
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3. Energetics

Enthapy Changes In the Lab

  • To measure the enthalpy change of a reaction, you need to know the number of moles of the substances that is reacting and the change in temperature
  • Can be done using a polystyrene beaker (reduces heat loss and gain) and a thermometer
  • Combustion reaction (trickier because the reactant is burnt in water)- A copper calorimeter containing a known mass of water is often used, the known mass of reactant is burnt and the temperature change of the water is measured


  • q= heat lost or gained (in joules), same as the enthalpy change if the pressure is constant
  • m= mass of water in the calorimeter, or solution in the polystyrene beaker (in grams)
  • c= specific heat capacity of water--> 4.18 Jg-1K-1
  • deltaT= change in termperature of water (combustion reaction) or solution
  • The enthalpy change is often different to the one calculated due to heat being lost to the environment or incomplete combustion
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4. Energetics

Hess's Law- The total enthalpy change is independant of the route taken

  • Indirect method of finding the enthalpy change or reaction using combustion or formation data
  • When using formation data- elements have a value of 0 because the element is being formed from its element meaning there is no change. The arrows point down
  • When using combustion data- oxygen has a value of 0 because the reactants are being burnt in oxygen. The arrows point up

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1. Group 7


  • Flourine (F2)- Colour: Pale yellow, Physical state: Gas
  • Chlorine (Cl2)- Colour: Green, Physical state: Gas
  • Bromine (Br2)- Colour: Red-brown, Physical state: Liquid
  • Iodine (I2)- Colour: Grey, Physical state: Solid
  • Boiling points- Increases down the group due to the increasing strength of Van der Waals foces as the relative size and mass of the atoms increases. Shown in the physical state changes of the elements: Flourine (gas)-->Iodine (solid)
  • Electronegativity- Decreases down the group due to the size of the atom increasing. This means it is harder for the atom to attract a pair of electrons

Displacement Reactions

  • Oxidising agents
  • Become less oxidising going down the group due to them becoming less reactive
  • A halogen will displace a halide from solution if the halide is below it in the periodic table
  • If bromine is diplaced, the solution will turn orange and if iodine is displaces the solution will turn brown
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2. Group 7


  • Mixing chlorine gas with dilute sodium hydroxide at room temperature froms sodium chlorate (l) solution--> bleach
  • 2NaOH + Cl2 --> NaClO + NaCl + H2O
  • The oxidation state of Cl goes up and down--> disproportionation

Chlorine and Water

  • Chlorine goes under disproportionation
  • Forms hydrochloric acid and chloric (l) acid: Cl2 + H2O <--> HCl + HClO
  • Aqueous chloric (l) acid ionises to make chlorate (l) ions: HClO + <--> ClO- + H3O+
  • Chloate (l) ions kill bacteria
  • Strengths- Kills disease-causing bacteria, prevents reinfection, prevents growth of algae
  • Limitations- Chlorine gas is very harmful when inhaled (irritates the respiratory system), react with organic compounds to form chlorinated hydrocarbons which are carcinogenic
  • Ethics- Forced mass medication
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3. Group 7

Halide Ions

  • Compounds with the -1 halide ion e.g. HCl, NaBr, KI
  • Reducing power- increases down the group due to the attraction between the nucleus and the outer electrons becomes weaker (ions get bigger, more shielding) meaning they can be lost more easily

Reactions with sulfuric acid

  • All halides react with conc. sulfuric acid to a give a hydrogen halide
  • NaF or NaCl with H2SO4- hydrogen flouride/hydrogen chloride gas is produced creating misty fumes, HF and HCl are not strong enough reducing agents to reduce H2SO4 so the reaction stops
  • NaBr with H2SO4- hydrogen bromide gas is produced, giving off misty fumes, HBr is a strong enough reducing reagent to react with H2SO4 in a REDOX reaction, the reaction produces choking fumes of SO2 and orange fumes of Br2 (2HBr + H2SO4 --> Br2 + SO2 + 2H2O)
  • NaI with H2SO4- hydrogen iodide gas is produced, giving off misty fumes, HI then reduces H2SO4 but keeps going and reduces SO2 to H2S (6HI + SO2 --> H2S + 3I2 + 2H2O)
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4. Group 7

Tests for Halides

  • Siliver nitrate test
  • Step 1: Add dilute nitric acid to remove ions which might interfere with the test
  • Step 2: Add silver nitrate solution (AgNO3), a pricipitate of silver halide is prodiced
  • Ag+ + X- --> AgX
  • The colour of the precipitate identifies the halide
  • Flouride (F-)- No precipitate
  • Chloride (Cl-)- White precipitate
  • Bromide (Br-)- Cream precipitate
  • Iodide (I-)- Yellow precipitate
  • Ammonia solution can also be used, with each silver halide having a different solubility in ammonia
  • Chloride (Cl-)- white precipitate, dissolves in dilute NH3
  • Bromide (Br-)- cream precipitate, dissolves in conc. NH3
  • Iodide (I-)- yellow precipitate, insoluble in conc, NH3
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1. Group 2

Atomic Radius

  • Atomic radius increases down the group, this is due to the extra electron shells as you go down

Ionisation Energy

  • Ionisation energies decreases down the group
  • Each element down group 2 has an extra electron shell compared to the one above
  • The extra inner shells shield the outer electrons from the attraction of the nucleus
  • The outer electrons move further away from the nucleus reducing its attraction
  • These factors make it easier to remove the outer electrons resulting in a lower ionisation energy


  • Increases down the group
  • Ionisation energies decrease going down the group meaning reactivity increases as electrons are more easily lost
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2. Group 2

Melting Points

  • Decreases down the group
  • Group 2 elements have typical mettalic structures with the electron in the outer shells being delocalised
  • Going down the group the metal ions get bigger (smaller charge/volume ratio) but the number of delocalised electrons remains constant, this means the delocalised electrons get more spread out
  • These two factors mean there's reduced attraction of the positive ions to the sea of delocalised electrons, it therefore takes less energy to break the bonds
  • The melting point of magnesium is much lower than expected, this is because the crystal structure (arrangement of metallic ions) changes

Reactions with water

  • When group 2 elements react, they are oxidised from 0 to +2
  • React with water to from metal hydroxide and hydrogen- M + 2H2O --> M(OH)2 + H2
  • React more readily going down the group 
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3. Group 2

Solubiltity of Compounds

  • Group 2 elements that contain singly charged negative ions (OH-), increase in solubility down the group
  • Grup 2 elements that contain doubly charged negative ions (SO4 2-) decrease in solubility down the group
  • Barium sulphate is insoluble

Testing for sulfate ions

  • If acidified barium chloride (BaCl2) is added to a solution containing sulfate ions, then a white precipitate of barium sulfate will form
  • *Barium chloride must fist be acidifies to get rid of any sulfites or carbonates

Barium meals

Used to make soft tissues such as the digestive system opaque so that they can be seen on X-rays. The patient swallows the barium meal in a suspension of barium sulfate

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4. Group 2


  • Known as alkaline earth metals
  • Many of the common compounds are used for neutralising acids
  • Calcium hydroxide (Ca(OH)2) is used in agriculture to neutralise acidic soils
  • Magnsium hydroxide (Mg(OH)2) is used in some indigestin tablets as an antacid
  • OH- + H+ --> H2O
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1. Alcohols

Primary, Secondary and Tertiary

  • Homologous series has the general formula CnH2n+1OH
  • Primary- OH group attached to a carbon with one alkyl group
  • Secondary- OH group attached to a carbon with two alkyl groups
  • Tertiary- OH group attached to a carbon with three alkyl groups

Industrial production of ethanol by fermentation

  • C6H12O6 --> 2C2H5OH + 2CO2
  • Exothermic process, carried out by yeast in anaerobic conditions
  • Yeast produces an enzyme (optimum temperature 30-40 degrees) that converts sugars such as glucose into ethanol and carbon dioxide
  • When the solution reaches 15% ethanol, the yeast dies, fractional distillation used to increase concentration of ethanol
  • Uses cheap equipment and is a renewable process
  • Ethanol must still be purified
  • Thought to be a carbon neutral process, however there are still carbon emissions when looking at the whole process
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2. Alcohols

Dehydration of Alcohols to form Alkenes

  • Ethene can be made my eliminating water from ethanol in a dehydration reaction
  • C2H5OH --> CH2=CH2 + H2O
  • The ethanol must be refluxed with conc. sulfuric acid and the process takes place in two stages:
  • Step 1- C2H5OH + H2SO4 --> C2H5OSO2OH + H2O
  • Step 2- C2H5OSO2OH --> CH2=CH2 + H2SO4
  • Sulfuric acts as a catalyst as it is unchanged at the end of reactin
  • Phosphoric acid (H3PO4) can be used as an alternative
  • This process allows you to form alkenes from renewable resources
  • Allows you to produce polymers without oil
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3. Alcohols


  • The oxidising agent acidified potassium dichromate (Vl) is used which turns from orange to the green chromium (lll) ion (Cr 3+) when reduced
  • Primary alcohols- Oxidise to aldehydes and then to carboxylic acids, you can control how far the reaction goes: aldehyde--> distilling the aldehyde off immediately using distilling apparatus, carboxylic acid--> vigorously oxidise the alcohol by heating it under reflux
  • Secondary alcohols- Oxidise to ketones under reflux conditions
  • Tertiary alcohols- Cannot be oxidised using potassium dichromate, the only ways is by burning them

Distinguishing between Aldehydes and Ketones

  • Fehling's solution of Benedict's solution- both deep blue complexes that turn brick red (CuO) when warmed with an aldehyde, remains blue when warmed with a ketone
  • Tollen's reagent (silver mirror)- reduced to silver when warmed with an aldehyde but not with a ketone 
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