AS Chemistry

Edexcel

Unit 2

Application of Core

Principles

Unit 2.1 Shapes of molecules & ions (p144 – 153 in text book)

1. Predicting Shapes of molecules The bonding of atoms will impact on the shape of the molecule formed. The shape of a molecule can determine both physical and chemical properties. In particular, enzyme activity is controlled by the shape since this decides the active site.

2. Shapes of ionic compounds Ionic bonds are non-directional. The ions will always be electrostatically attracted to each due to their charge. Conversely, they will also repel. As a result a lattice structure is formed to maximise attraction and minimise repulsion.

3. Shapes of covalent compounds Covalent bonds are highly directional. As a result, these compounds have a very specific shape which governs the chemical and physical properties. The shape of a molecule depends on:

the number of electron pairs around the central atom electron pair repulsion

The electron pairs align themselves in such a way to minimise the amount of repulsion. The bond angles depend on the number of electron pairs around a central atom.

(Refer to table 2.1.1 p145)

Number of electron pairs around the central atom

2 3 4 5 6

7

Example BeCl2 BF3 CH4 PCl5 SF6 IF7

Dot & cross diagram

Shape

Bond angle

Name of shape

4. The effect of lone pairs on shapes of molecules Some molecules have lone pairs of electrons around the central atom as well as the bonding electrons. These lone pairs impact on the shape since they too become involved in electron repulsion. However, the lone pair electrons are closer to the nucleus than the bonding electron pairs. As a result, the repulsion between a lone pair and a bonding pair is greater than that of just 2 bonding pairs. This impacts on bond angles. Examples Ammonia NH3 Water H2O Xenon tetrafluoride XeF4

5. Multiple bonds These are treated in the same way as single bonds. The electron pairs are still arranged to minimise repulsion. Examples Carbon dioxide CO2 Sulphur dioxide SO2

6. Shapes of ions The ions are formed when electrons are lost or gained. However, electron pair repulsion can still be used to predict the shape of the ion by viewing the electron pairs around the central atom. Examples Ammonium NH4

+ Sulphate SO42- Nitrate NO3

-

7. Shapes of simple organic molecules These molecules generally have a tetrahedral structure, but features such as double bonds will have an impact on the shape. (Refer to p148 – 149)

a. Alkanes All of these molecules show a tetrahedral shape around each carbon atom with a bond angle of 109.5° E.g. Ethane

b. Alkenes These molecules have a planar structure around the double bond E.g. Ethene

c. Alcohols These molecules have a tetrahedral structure, but lone pair repulsion decreases the angle around the oxygen atom E.g. Methane

d. Carboxylic acids Methanoic acid is planar. In all other carboxylic acids the functional group is planar, with the rest of the molecule displaying tetrahedral structure.

e. Halogenoalkanes These molecules are tetrahedral. E.g. Bromomethane

f. Aldehydes & ketones These molecules have a planar structure around the carbonyl group, with tetrahedral structure elsewhere in the molecule. E.g. Propanal & propanone

8. Allotropes of carbon (Refer to p150 – 153) Allotropes are different forms of the same element that exist in the same physical state. This is due to different bonding between the atoms resulting in different shapes.

a. Diamond Pure diamond is made from interlocking crystals of carbon covalently bonded in a tetrahedral structure. Each carbon atom has 4 bonds. (Refer to fig 2.1.16 p150) This structure is giant molecular and means that diamond is extremely hard allowing it to be used for cutting tools. Pure diamond is colourless. However, mineral contaminates can colour them from opaque to black. Diamonds are chemically inert and poor conductors of electricity. Since 1955 diamonds have been able to be manufactured synthetically. This is done by subjecting graphite to high temperature and pressure. These were only good enough for industrial processes until 1970. However, the synthetic diamonds of gem quality are so expensive to make they aren’t used for jewellery.

b. Graphite This is a black/grey lustrous substance that easily crumbles due to the tendency of the layers to cleave. Three carbon atoms are bonded together in hexagonal layers, leaving delocalised electrons to flow between the layers. The layers are held together by weak London forces. (Refer to fig 2.1.17 p151) Graphite is chemically inert, but less so than diamond. However, whilst a poor heat conductor, it is a good conductor of electricity.

c. Amorphous carbon This is a deep black powder, often seen as soot. It can be formed through heating organic substances with a lack of oxygen. It is the most reactive form of carbon and easily burns in air. It appears to be similar in structure to graphite. It is used for black pigment in paints and inks, along with stabilising filler for rubber and plastics.

d. Fullerenes Buckminster fullerene is a spherical structure with 60 atoms bonded together. Other fullerenes exist with different numbers of carbon atoms. They have delocalised electrons so are able to conduct electricity.

9. Nanochemistry (Refer to p152 & 153 especially the HSW boxes) Nanotubes were developed in 1991. These are elongated cage-like structures which can form highly complex shapes. Various uses are being trialled for nanotubes and nanoparticles to include;

i. superconductors ii. composite materials iii. nanoprobes to identify cancerous cells iv. cosmetic preparations

Unit 2.2 Intermediate Bonding & Bond Polarity (p154 – p159 in text book)

1. The position of the electrons in a covalent bond When the 2 atoms in a covalent bond are identical to each other, e.g. in a hydrogen molecule, the bonding electrons are shared equally between the 2 atoms. However, when the 2 atoms in a covalent bond are different to each other, e.g. in a hydrogen chloride molecule, the bonding electrons are no longer shared equally between the 2 atoms due to different attractions from the nuclei. Electronegativity is a measure of the attraction of an atom in a molecule for a pair of electrons in a covalent bond.

2. The Pauling scale The most common scale used to measure electronegativity is the one produced by Linus Pauling. It is a relative scale, so without units, between 0 and 4. The higher the number, the more electronegative the element is. The most electronegative element is fluorine with a value of 4, while the least electronegative element is caesium with a value of 0.7. (Refer to fig 2.2.4 p155) There are several trends in electronegativity;

non-metals have higher electronegativity values than metals electronegativity increases across a period electronegativity decrease down a group

The noble gases do not have electronegativity values since they don’t have an affinity for electrons. Electron density maps can be used to view the uneven sharing of electrons in a covalent bond. (Refer to HSW box p155)

3. Bond character (Refer to table 2.2.1 p156) Ionic and covalent bonding models represent two extremes. In fact, they sit at each end of a continuum. By calculating the electronegativity difference in a bond, the character of the bond can be calculated i.e. how much covalent character it exhibits and how much ionic character it exhibits.

4. Polar bonds

a. Covalent bonds (Refer to fig 2.2.7 p157) The uneven sharing of bonding electrons results in a distortion of the electron cloud. As a result, small charges occur at each end of the bond. These are designated as δ- and δ+. This produces a polar bond e.g. hydrogen chloride.

b. Ionic bonds (Refer to fig 2.2.8 p157) Ionic bonds may be distorted if the cation has a high charge density since it can polarise the electron cloud of the anion. If the distortion is great enough, the electron clouds can overlap to resemble a covalent bond.

5. Polar molecules It is possible to have several polar bonds in a molecule without the molecule itself being polar. Polar liquids can be deflected by an electric charge. This property is dependant on the symmetry of the molecule.

Liquids deflected Liquids undeflected

For a molecule to be polar, it needs to be asymmetrical so that there is an uneven distribution of charge across the molecule. The charge separation makes it a dipole. This polarity is measured as its dipole moment – the amount of charge separation multiplied by the distance between the two centres of charge. The unit of the dipole moment is the debye, D. Examples

i. Boron trifluoride, BF3 Although each B-F bond is polar, this is a non-polar molecule since the molecule has a net polarity of zero.

ii. Water, H2O The molecule is v-shaped so it isn’t symmetrical. Both sets of bonding electrons are attracted towards the oxygen atom. As a result there is an overall dipole moment.

6. Predicting whether a substance is polar or not

Formula

Description Example

Polar

Non-polar

Unit 2.3 Intermolecular Forces (p160 – p169 in text book)

1. Polarity & boiling temperature (Refer to fig 2.3.1 p160) Non-polar molecules have much lower boiling points than polar molecules. The forces between polar molecules must therefore be much bigger since it takes more energy to break these.

2. Comparing intermolecular forces Hydrogen Bonds London Forces Permanent

Dipole-dipole Strength Strongest Weak Strong Melting and Boiling Points

High Low High

Bonding Molecules

Hydrogen and Oxygen, Fluorine or Nitrogen

All atoms and molecules

Polar Molecules

Reason When hydrogen bonds with any of the above elements the electrons are pulled towards the said element because of the large electronegativity difference. The lone pairs on the oxygen, fluorine or nitrogen are then weakly bonded to the hydrogen atoms.

Electrons move around fast in an atom and at any given time they are more likely to be at one side of an atom. This causes a dipole which then induces a dipole in neighbouring atoms or molecules. These dipoles are weakly attracted.

Some molecules have an uneven share of electrons in there bonds due to polarisation. This creates a dipolar molecule. The charged ends of these molecules attract each other.

3. Trends in physical properties

a. Trends in alkanes

(Refer to table 2.3.3 p164) The melting and boiling points of the straight chain alkanes increase as the chain length increases. The forces between alkanes are weak but they grow larger as the chain length increases due to a greater surface area. (Refer to table 2.3.4 p165) Branching of alkane chains generally lowers the melting and boiling points. The side chains interfere with the packing of the molecules and also reduce the number of points where London forces can occur.

b. Comparing alkanes & alcohols (Refer to table 2.3.5 p165) Alcohols have higher melting and boiling points than alkanes. This is due to the hydrogen bonding that can take place between alcohol molecules.

c. Trends in boiling temperature of the hydrogen halides (Refer to fig 2.3.12 p166) Boiling temperatures decrease from HI to HCl as you move up group 7. However, HF has a much higher boiling point than any of the other hydrogen halides. The electronegativity difference between hydrogen and fluorine is greater. The result is the formation of hydrogen bonds between the molecules so more energy is required for the state change.

4. Solubility Water is a very good solvent since it dissolves a wide variety of substances. However, not all substances will dissolve in water e.g. chlorophyll. Other solvents, often called non-aqueous solvents, will dissolve many of the substances that water won’t. Solubility strictly means the mass of a solute that dissolves in 100g of solvent at a particular temperature.

a. Patterns in solubility

There are some general rules about solubility;

highly polar solids such as ionic salts dissolve in water (a polar solvent) but not in hexane (a non-polar solvent)

polar organic substances such as glucose dissolve in water but not in hexane

non-polar solids such as candle wax dissolve in hexane but not in water

non-polar liquids such as petrol and diesel mix completely. They are miscible

polar liquids such as water and ethanol are miscible a polar liquid and a non-polar liquid such as water and oil do not mix

together. They are immiscible and form separate layers

b. Dissolving an ionic solid in water (Refer to fig 2.3.14 p168) Sodium chloride dissolves in water with no apparent change in energy. The ionic solid has a lattice energy which needs to be applied to break down its structure. However, the polar water molecules align themselves around the ions through the process of hydration which provides a hydration enthalpy. In this case the lattice energy and hydration enthalpy are balanced, hence the apparent lack of energy change. In sodium hydroxide the lattice energy is less than the hydration enthalpy so it dissolves in water as an exothermic process. Lithium fluoride, however, is insoluble in water since its lattice energy is greater than the hydration enthalpy.

c. The solubility of alcohols in water (Refer to fig 2.3.15 p169) Alcohols are able to dissolve in water since they contain the polar OH group. However, their solubility decreases as the length of the carbon chain increases. Carboxylic acids follow the same pattern.

d. Why are non-polar substances insoluble in water? Non-polar substances do not dissolve in water. The London forces between the non-polar substances are weaker than the hydrogen bonds in water. As a result, the non-polar substances are unable to disrupt the arrangement of the water molecules. (Refer to fig 2.3.16 p169)

Even halogenoalkanes are unable to dissolve in water since most of the molecule is non-polar.

e. Mixing two organic liquids (Refer to fig 2.3.17 p169) Hexane and octane are 2 organic liquids that mix together. They both have weak London forces so when they are mixed together the London forces extend throughout the mixture.

Unit 2.8 Organic chemistry – alcohols & halogenoalkanes (p208 – 221 in text book)

1. Alcohols These are very important organic compounds in industry. They are good solvents and make suitable raw materials. They have the general formula CnH2n+2O (also written as CnH2n+1+OH). They have an alkyl chain with an OH group attached and are named in the same way as other organic molecules. (Refer to table 2.8.1 p208)

Alcohol Displayed formula Skeletal formula

Methanol

Ethanol

Propan-1-ol

Butan-2-ol

2-methylbutan-2-ol

2. Types of alcohols Alcohols can be classified as either primary, secondary or tertiary depending on the position of the functional OH group. (Refer to fig 2.8.2 p209)

Primary Secondary Tertiary

Example

Definition

3. Reactions of alcohols

Alcohols will undergo a number of different reactions.

a. Combustion Just like hydrocarbons, the products of complete combustion are carbon dioxide and water. For incomplete combustion, carbon monoxide will be formed. E.g. Propan-1-ol

b. Reaction with sodium When sodium is added to an alcohol, it reacts with the hydrogen atom of the OH group to form H2. The other product is a salt. The alcohols are behaving like weak acids. The strength of the ‘acid’ decreases down the homologous series. E.g. Propan-1-ol & sodium

c. Test for alcohols / Making halogenoalkanes Alcohols react with PCl5 to form a chloroalkane and hydrogen chloride gas. E.g. CH3CH2OH(l) + PCl5(s) CH3CH2Cl(l) + POCl3 + HCl(g) This particular reaction results in white fumes being released (the hydrogen chloride) which can be used as an observation to confirm the presence of an OH group. Similar reactions occur with other phosphorous halides e.g. Propan-1-ol & PBr5 Phosphorous (III) halides will also react with alcohols to produce halogenoalkanes E.g.

3 CH3CH2OH + PF3 3 CH3CH2F + H3PO3

4. Oxidation of alcohols These can be used to distinguish between primary, secondary & tertiary alcohols since they each react differently with common oxidising agents such as acidified potassium dichromate (K2Cr2O7) and acidified potassium permanganate (KMnO4).

a. Primary alcohols These are quickly oxidised to aldehydes which can be rapidly changed into carboxylic acids. The common product is the acid unless the aldehyde is separated during the reaction. E.g. propan-1-ol

b. Secondary alcohols These are readily oxidised into ketones, but no further oxidation occurs so only one product is made. E.g. Propan-2-ol

c. Tertiary alcohols These aren’t readily oxidised by any of the common reducing agents

d. Colour changes (Refer to fig 2.8.6 p211)

Oxidising agent Start colour End colour if oxidation occurs

Potassium dichromate

Potassium permanganate

Colourless

5. Extracting the products of oxidation of alcohols

a. Reflux (Refer to fig 2.8.8 p211) This method is used to prepare carboxylic acids from primary alcohols. The reflux condenser prevents vapours from leaving the flask so that a full reaction can take place.

b. Distillation (Refer to fig 2.8.7 p211) This method is used to prepare an aldehyde from a primary alcohol. The aldehyde can be collected before it has the chance to react further to form a carboxylic acid. This method is also used to prepare a ketone from a secondary alcohol.

6. Halogenoalkanes

These organic molecules are rarely found naturally occurring on Earth. Most are synthesised. They are important in medicine, plastic production and agriculture. However, most of them are responsible for high levels of environmental damage, both in the atmosphere and on the ground. At least one of the hydrogen atoms from an alkane is replaced with a halogen to form these molecules. Consequently there are many different isomers. (Refer to fig 2.8.10 p212)

Primary Secondary Tertiary

Example

Definition

7. Uses of halogenoalkanes

There are many uses of these chemicals.

a. Anaesthetics (Refer to HSW box p213) These have been developed to aid surgery over the years since it is much easier to operate on an unconscious patient than a conscious one. In 1841 ethoxyethane (ether) was discovered to be an anaesthetic but it is highly flammable. In 1844 nitrous oxide was used in dentistry. Whilst not flammable or toxic it is only a mild anaesthetic. Towards the end of the 1840s trichloromethane (chloroform) was discovered to be a better anaesthetic than these earlier two. The middle of the 20th century saw the discovery of 2-bromo-2-chloro-1,1,1-trifluoroethane (halothane) as an effective deep anaesthetic.

b. Dry cleaning The most common dry cleaning fluid is tetrachloroethene.

c. Refrigerants (Refer to p214 & 215) These are liquids which circulate in a refrigerator, constantly changing phase from liquid to gas and back again in order to transfer heat from inside the fridge to the outside. A good refrigerant should; CFCs (chlorofluorocarbons) have been widely used refrigerants. There are so many of these refrigerants that they have had a separate naming system developed.

Rightmost digit

Tens digit

Hundreds digit

Thousands digit

Any remaining bonds not accounted for are occupied by chlorine atoms

A suffix of a lower case letter a, b or c indicates unbalanced isomers e.g. R134a

When CFCs are released into the atmosphere they begin to react once they reach the ozone layer. Stage 1 – homolytic breakdown under UV light e.g. Stage 2 – reaction of a chlorine radical Since the chlorine atom isn’t used up in the reaction, one molecule of CFC can destroy thousands of ozone molecules. In 1987 the Montreal Protocol set out plans for the reduction of CFC use. Instead, HCFCs (hydrochlorofluorocarbons) have been synthesised to replace CFCs. Although they pose a much lower threat to the ozone layer they are very potent greenhouse gases.

d. Fire retardants (Refer to HSW box p216 & 217) These were used in fire extinguishers designed to tackle electrical fires and were identified in green containers. However, their contribution to the ozone layer depletion and the narcotic effect of the gases has seen them withdrawn from use. Instead they have been replaced by the blue CO2 extinguishers. TBBPA contains approximately 59% bromine and is used to cover circuit boards as a fire retardant. Synthetic fibres often have some halogen content in order to reduce the size of the flame should they catch fire. Such a benefit means that children’s clothing is less likely to catch fire near a naked flame.

e. Insecticides DDT was widely used as an insecticide but causes toxic build up higher up in food chains so its use has now been prohibited.

8. Reactions of the halogenoalkanes The chemistry of these compounds depends on the halogen atom present along with its position on the alkyl chain. Most reactions are nucleophilic substitutions but under some conditions elimination will occur.

a. Reaction with aqueous hydroxide or water This is a nucleophilic substitution reaction but can also be referred to as hydrolysis. The reaction with water is slow at room temperature but is much quicker with an OH- ion. The ideal conditions are under reflux with heat. The main product will be an alcohol. E.g. Equation for chloroethane with water Equation for bromomethane with NaOH The rate of reaction is inversely proportional to the electronegativity difference of the RX bond i.e. as the electronegativity of the halogen decreases, the faster the rate of reaction.

b. Reaction with alcoholic potassium hydroxide In this case elimination will take place under reflux with heat. The main product will be an alkene. E.g. reaction of ethanolic potassium hydroxide with 2iodobutane

c. Reaction with alcoholic ammonia Ammonia behaves as a nucleophile to undertake a substitution reaction to form an amine. E.g. reaction of 1bromopropane with ammonia However, the resulting amine also possesses a lone pair of electrons so further reaction can take place. This reaction must take place in a sealed tube since the ammonia would otherwise escape through the condenser if attempted under reflux.

9. Identifying halogenoalkanes When these compounds react it usually results in the release of a halide ion. This can then be tested for using silver nitrate and ammonia.

Halide ion Observation with silver nitrate

Observation with the addition of ammonia

Chloride

Bromide

Iodide

Plotting the production of these halide ions can be used to determine the rate of reaction. The further down group 7 the faster the halogenoalkane will react. Also, the tertiary halogenoalkane reacts faster than the secondary, which in turn is faster than the primary i.e. more branching gives a faster reaction.

10. Preparing halogenoalkanes This can be done by reacting an alcohol under reflux. (Refer to fig 2.8.18 p221) The flask is put into a beaker of cold water and concentrated sulphuric acid is added slowly from a dropping funnel. The flask is cooled since the reaction is exothermic at this stage. Refluxing on a water bath takes place for 30 minutes. (Refer to fig 2.8.19 p221) The mixture is then distilled with the distillate collected in two layers – an upper aqueous layer and a lower organic layer. The aqueous layer is discarded while the organic layer is purified and redistilled to give a pure product.

Unit 2.10 Mass spectra and IR absorption (p230 – 235 in text book)

1. The mass spectrometer This was studied in unit 1.3. For organic chemistry it can be applied in the use of pharmacology, space research, radioactive dating and catching drug cheats in sport.

2. Using a mass spectrometer to analyse organic compounds

Identifying an organic compound from the huge numbers that exist can be a daunting and long task. The use of mass spectrometry and infrared spectroscopy can make it much quicker. The organic compound can form a molecular ion when bombarded with the high speed electrons. An m/z value produces a molecular ion peak which is usually the peak with the highest m/z value. Other peaks can occur due to fragmentation of the molecule. Some peaks are more difficult to interpret than others. Some can be easily identified as being from the loss of hydrogen, a methyl group, an OH group etc. However, even in simple molecules, it is impossible to identify all the peaks due to complicated rearrangements of the cations inside the machine. (Refer to specific examples of butane, ethanol, propanal and propanone on p230 – 231)

3. Absorbing infrared (Refer to fig 2.10.8 p232) If a substance is irradiated with IR its molecules will absorb some of this radiation. This absorption is as a result of molecules vibrating due to either a stretch or a bend in the bonds. The amount of energy required to make them vibrate depends on;

- the bond strength - the bond length - the mass of each atom involved in the bond

A range of IR frequencies are passed through the sample in an infrared spectrometer. The spectrum identifies a series of troughs where energy is absorbed, although these are referred to as peaks. The frequency of the IR absorption is measured in wavenumbers (cm-1) shown on the x axis. The percentage transmission is shown on the y axis. One beam of IR is passed

through the sample while a second is passed through a reference cell to ensure only IR absorption in the compound is observed on the spectrum.

4. Bending & stretching A higher frequency of IR is required to make stronger bonds vibrate. The bonds can bend, stretch symmetrically or stretch asymmetrically. Each type of vibration has a corresponding peak. (Refer to fig 2.10.9 p233) Only molecules which change their polarity will absorb IR. Consequently, molecules such as H2 and Cl2 will not have any peaks on an IR spectrum but HCl would. It is polar molecules in the atmosphere such as CO2, and H2O that can absorb IR which imparts on the greenhouse effect. Functional groups within organic molecules can be recognised from their characteristic absorption peaks which are recorded in reference tables. (Refer to table 2.10.3 p233) (Refer to specific examples on p234 – 235)

Unit 2.5 The Periodic Table – Groups 2 & 7 (Refer to p178 – 195 in text book)

1. Group 2: the alkaline earth metals Group 2 elements behave in a similar way to group 1 elements, with some clear patterns as you go down the group.

a. Ionisation energies of group 2 elements There is a large increase for all the group 2 elements between the second and third ionisation energies. They readily lose 2 electrons to form 2+ ions, but the large third ionisation energy makes it impossible to form a 3+ ion. Losing 2 electrons gives them noble gas configuration. (Refer to table 2.5.1 p179) The ionisation energy decreases down the group. There is an increasing nuclear charge down the group, but there is also an increase in atomic radius. This means that the outermost electrons experience a weaker attraction to the nuclear charge as you move down the group alongside a greater shielding effect from inner electron shells.

b. Flame tests These are carried out to test for the presence of specific metal ions. Concentrated hydrochloric acid is used to clean platinum/nichrome wire which is then used to hold a sample in a hot Bunsen flame. The colour of the flame can be used to identify the cation.

Cation Flame colour

Lithium

Sodium

Potassium

Magnesium

Calcium

Strontium

Barium

The flame colour is produced when the electrons in the cation return to their ground state after being excited. However, some cations don’t produce a colour since there is too large a gap between energy levels for the electrons to become excited.

A flame photometer can be used to provide a more accurate result by measuring the wavelength of light emitted by the cation. It is often used to find sodium and potassium levels in blood samples and wine.

c. Reactions of the group 2 metals There are also obvious chemical trends down group 2 as well as the physical trends.

i. Reaction with oxygen Reactivity with oxygen increases down the group to form a metal oxide e.g.

ii. Reaction with chlorine Reactivity with chlorine also increases down the group to form solid metal chlorides e.g.

iii. Reaction with water The reactivity increases down the group. (Refer to table 2.5.4 p181) Magnesium reacts with steam to form magnesium oxide and hydrogen. Calcium, strontium and barium react with water to form a hydroxide and hydrogen e.g. Group 2 oxides and hydroxides aren’t naturally occurring compounds. However the cations are often found in carbonates which can undergo thermal decomposition to form the oxide, which then reacts with water to form the hydroxide.

d. Reactions of the group 2 oxides

i. Reaction with water Beryllium oxide doesn’t react with water, while magnesium oxide will only slightly react. However, calcium, strontium and barium oxides will react vigorously with water. The calcium hydroxide (slaked lime) produced is used in water treatment, neutralising acidic soils and making whitewash, mortar and plaster.

ii. Reaction with dilute hydrochloric and nitric acids Group 2 metal oxides will react with these acids to form a salt and water e.g.

e. Solubility of group 2 compounds Barium sulphate is used in barium ‘meals’ or enemas. Since it is insoluble in water it shows up well on X-rays without poisoning patients. All the group 2 nitrates and chlorides are soluble. Group 2 salts where the anion has a charge of -2 (such as sulphates) are generally insoluble with the exception of a few magnesium and calcium salts. The solubility of group 2 compounds tends to decrease down the group as atomic number and ionic size of the cation increase. However, the solubility of hydroxides increases. (Refer to table 2.5.5 p183) Barium sulphate is used to test for sulphate ions. First dilute hydrochloric acid is added to the sample to destroy any carbonates. Then barium chloride or barium nitrate is added. A white precipitate of insoluble barium sulphate confirms the presence of sulphate ions. Group 1 metals produce hydroxides, sulphates, chlorides, nitrates and carbonates that are soluble in water. Group 2 compounds are likely to have larger lattice energies due to the increased ionic charge, but don’t have high enough hydration enthalpies for solubility.

f. Thermal stability of the salts of the s-block elements The stability of ionic compounds increases as;

- cationic radius decreases - the charge on the ion increases

The lattice energies of the carbonates and oxides fall as you go down both groups 1 and 2 because the cations are increasing in size. The distances between the ions increases so the attractions become weaker. The lattice energies of the carbonates and oxides fall at different rates due to the different sizes of the anions. The oxide ion is relatively small whereas the carbonate ion is much larger. In a Bunsen flame group 1 carbonates are thermally stable with the exception of lithium carbonate. However, the group 2 carbonates decompose to form stable oxides. The temperature at which this happens increases down the group. Beryllium carbonate doesn’t exist at room temperature it is so thermally unstable. The s-block nitrates decompose upon heating in a Bunsen flame. Group 1 nitrates (except lithium nitrate) form the corresponding nitrite e.g. Group 2 nitrates (and lithium nitrate) form the corresponding oxide e.g. The hydroxides of groups 1 and 2 follow the same pattern as the carbonates and nitrates. All group 1 metal hydroxides are stable up to quite high temperatures. All of the group 2 metal hydroxides decompose to give the corresponding oxide e.g.

2. Group 7: the halogens

a. Physical properties of the halogens The halogens are diatomic molecules which are very reactive and are strong oxidising agents. The most common oxidation state is -1, although others exist (except for fluorine). Down the group they become less reactive, darker in colour and less volatile.

Halogen Colour State at room temperature

Fluorine

Chlorine

Bromine

Iodine

Their solubility in water decreases down the group. Chlorine reacts in water to form a mixture of hydrochloric acid and chloric acid. Bromine reacts in a similar way. However iodine is almost insoluble in water but is soluble in potassium iodide solution due to the formation of I3- ions. (Refer to table 2.5.9 p188 for other physical properties) Halogens are non-polar so they are more soluble in hydrocarbon solvents than in water. (Refer to fig 2.5.13 p189)

b. Oxidation reactions of the halogens

i. with metals When halogens react with metals they become reduced to halide ions e.g.

ii. with non-metals This is done through covalent bonding e.g.

iii. with iron (II) chloride solution The halogen oxidises the pale green iron (II) chloride solution to brown iron (III) ions e.g.

iv. Disproportionation The reaction of chlorine, bromine and iodine with sodium hydroxide depends on the temperature. In cold temperatures a mixture of halide and halate (I) ions are formed. The halate (I) ions may decompose to form more halide ions and halate (V) ions. The rate of decomposition depends on the halogen used and the temperature. Chlorine decomposes slowly at 15oC but rapidly at 70oC. For bromine both reactions happen rapidly at 15oC and at 0oC the decomposition is prevented. However, for iodine, both reactions are rapid at 0oC so it is difficult to separate the products. These are all examples of disproportionation since the halogen is both oxidised and reduced in the same reaction.

v. Redox reactions with the potassium halides Potassium halides will react with another halogen in a displacement reaction e.g.

c. Reactions of the hydrogen halides

i. with ammonia This reaction produces ammonium halides e.g.

ii. with concentrated sulphuric acid Halides will react to form the hydrogen halide since the sulphuric acid is a proton donor e.g. However, with bromide and iodide compounds the sulphuric acid acts as an oxidising agent to produce the halogen