C4 - Inorganic Chemistry and the Periodic Table

Group 2 elements

• Form 2+ ions

• Can be seen by their electron configurations

• The electrons in the outer s sub-shell are relatively weakly bound

  • Are lost during reactions

• Loss of two electrons results in a stable noble gas electron config

G2 FIE

• Decreases down group 2

• Energy is a measure of the strength of electrostatic attraction between the outer electron and nucleus

• Decreases due to

  • Nuclear charge increases

    • More protons

  • Atomic radius increases

    • More electron shells are added

  • Electron shielding increases

    • More inner electron shells reduce nuclear attraction

  • Shielding and atomic radius are stronger than nuclear charge

G2 Reactivity (water, oxygen and chlorine)

• Reactivity increases down the group

• Reactions of group 2 elements down group with water, oxygen and chlorine become more vigorous

• Reaction with water

  • Produce metal hydroxide and hydrogen

  • Group two metal is oxidised from 0 to +2

  • Reactivity trend

• Reaction with oxygen

  • When group 2 metals burn in oxygen, a solid white oxide is formed

    • 2M + O₂ ---> 2MO

    • Metal is oxidised from 0 to +2

    • Oxygen is reduced from 0 to -2

• Reaction with chlorine

  • M + Cl₂ ---> MCl₂

  • Metal is oxidised from 0 to +2

  • Chlorine is reduced from 0 to -1

Solubility of G1 and G2 compounds

• Solubility of group 2 metal hydroxides and sulfates follow opposite trades

  • Group 2 metal hydroxides become more soluble

  • Sulfates become less soluble down the group

Reaction of G2 oxides and hydroxides

• Are basic and alkaline

• Group 2 metal oxides readily react with water to form metal hydroxides (dissolve)

  • MO + H₂O ----> M(OH)₂

  • Releases OH- ions which makes the solution alkaline

  • pH of the resulting solution increases down the group

    • Solubility of hydroxide increases

    • Exception of MgO which reacts slowly and has low solubility

• Group 2 metal hydroxides react with dilute acids such as HCl to form water and salt

  • Reactions become more vigorous down the group

G1-2 compounds thermal decomposition

• Group 1 compounds

  • Are very stable and don’t decompose with Bunsen burner heating

  • Only lithium carbonate does

    • Li₂CO₃ ----> Li₂O + CO₂

  • Group one nitrates decompose to nitrites and oxygen for example

    • 2NaNO₃ ---> 2NaNO₂ + O₂

    • Exception of lithium nitrate

      • 2LiNO₃ ----> Li₂O + 2NO₂ + 1/2 O₂

• Group 2 compounds

  • Group two carbonates decompose to oxides and carbon dioxide

  • Group 2 nitrates decompose into oxides, nitrogen dioxide and oxygen

Thermal stability of carbonates and nitrates of G1-2

• Increases as you go down G1 and G2

  • Due to ion polarisation and polarising power

• Cation size

  • As we move down group, cations become larger

  • Larger cations have a lower charge density

  • Results in lower polarising power, causing less anion distortion

  • Less distortion leads to stronger C-O and N-O bonds, so more stable

• Cation charge

  • G2 cations have higher charge than G1

  • Higher charge results in greater polarising power

  • Makes G2 less stable than G1 counterparts

• Testing thermal stability of carbonates and nitrates

  • G2 Carbonates

    • Measure time taken to produce CO₂ gas that can turn limewater cloudy

  • G1 nitrates

    • Tiem taken to produce enough O₂ gas to relight a glowing splint

  • G2 nitrates

    • Measure time taken to produce brown NO₂ gas

Flame test procedure

• Take a platinum wire loop

  • Unreactive so it will not affect the results

  • Really high melting point so wont melt

• Clean it with dilute hydrochloric acid, distilled water and heat it

  • Gets rid of any leftover ions

  • So the results do not get messed up

• Hold it in the blue part of the flame

  • Use BLUE otherwise you won't see a colour change

Flame test colours

• G1

  • Li - red

  • Na - orange/yellow

  • K - lilac

  • Rb - red

  • Cs - blue

• G2

  • Ca - orange-red

  • Sr - crimson

  • Ba - green

Tests for carbonate/hydrogen carbonate ions

• Add dilute nitric acid dropwise to  sample

• Observe for effervescence

  • CO₃2- + 2H+ ----> CO₂ + H₂O

  • HCO₃- + H+ -----> CO₂ + H₂O

• Pass bubbles of gas through limewater, does it turn cloudy?

  • Calcium carbonate forming

Identifying sulfate ions

• Add nitric dropwise to a sample

• Add a few drops of solution containing Ba2+ (barium chloride)

• Observe for white precipitate (barium sulfate)

Identifying halides

• Add nitric acid and silver nitrate dropwise

• Chloride - white

• Bromine - cream

• Iodine - yellow

• Add concentrated ammonia or dilute ammonia

Identifying ammonia ions

• Add NaOH dropwise and mix

• OH react with ammonia ions to produce ammonia gas

• Moisten a piece of red litmus paper and hold over the test tube

• Paper will turn blue if ammonia gas is present

Halogen physical properties

Highly reactive non-metals and tend to form diatomic molecules with one covalent bond

Melting and boiling points

• Increase down the group

  • Molecules become larger in size and relative mass

  • Allows for stronger DP-DP forces to develop between molecules

  • Stronger IMF require amounts of energy to overcome

  • Halogens melt and boil at higher temperatures down the group

Electronegativity

• Related to atomic size, electronegativity of halogens shows an overall decrease going down the group

• Refers to the tendency of an atom to attract a bonding pair of electrons

• Because

  • Nuclear charge increases

  • Atomic radius increases

  • Electron shielding increases

Halogen properties

• Gain an electron

• Results in the halogen being reduced as ON decreases from 0 to -1

• Halogens act as oxidising agents

Reactivity of G7

• Atomic radius increases down the group as more electron shells are added

• Increasing size leads to the other electrons being father from the positive nucleus

• Outer electrons also experience more shielding from inner electron shells

• Electrostatic attraction between outer electrons and nucleus gets progressively weaker

• Increase in atomic radius and shielding outweigh the increase in nuclear charge

  • Harder for larger halogens to attract the electron needed to form a negative ion

Oxidising power of halogens

• Decreases down group 7

• Because increasing size and shielding makes it harder for larger halogens to remove electrons

• Fluorine is the strongest oxidising agent

• Halogens displace less reactive halide ions

  • Displacing halogen is reduced as it gains an electron to form halide

  • Displaced halide is oxidised so it loses an electron to form a halogen molecule

  • Halogen will displace any halide ion below it

• Oxidising ability

  • Chlorine can oxidise both bromide and iodide

  • Bromine can oxidise iodide but not chloride

  • Iodine cannot oxidise any of them

Reaction of halogens with G1 and G2 metals

• G1

  • React with halogens in a 2:1 ratio

  • 2Na + F₂ ---> 2NaF

  • Na is oxidised and fluorine is reduced

• G2

  • Group 2 metals react with halogen in a 1:1

  • Mg + Cl₂ ---> MgCl₂

  • Mg is oxidised and chlorine is reduce d

Colour changes in halogen displacement

• Colour changes in aqueous solution

  • If Br- is displaced and Br2 is formed, solution turns yellow

  • Iodide is displaced, solution turns orange and brown

  • If no reaction occurs, aqueous solution remains colourless

• Colour changes in organic solution

  • Halogen products can be more clearly seen with an organic solvent

  • If it separates out on the top layer of the aqueous solution

    • Bromide is displaced, organic layer turns orange

    • Iodide is displaced, organic layer turns purple

    • No reaction, reaming colourless

Halide ions reaction

• Are negatively charged ionic forms of the halogen fluorine, chlorine bromine and iodine

• React by losing an electron to form neutral halogen

• Electron loss results in halide being oxidised

  • As it loses its electron, it causes another substance to be reduced

Reducing power of halides

• Increases down the group 7

• Because

  • Ionic radius increases down halogens as more electron shells are added

  • Increasing ionic radius leads to the outer electron being farther from the positive nucleus

  • Outer electrons experience more shielding

  • Electrostatic attraction between the outer electrons and nucleus gets weaker

  • Increase in ionic radius and shielding outweighs the increase in nuclear charge

• Fluorine is weakest reducing agent and iodide is the strongest

Reaction of halides with sulfuric acid: fluoride and chloride

• Misty white fumes of HF or HCl gas are seen

  • NaF + H₂SO₄ ---> NaHSO₄

  • NaCl + H₂SO₄ ---> NaHSO₄ + HCl

• These halides have low reducing power, so no further redox reactions occur

  • Oxidation number of sulfur remains at +6

Reaction of halides with sulfuric acid: bromide

• Misty fumes of HBr gas are produced initially

  • NaI + H₂SO₄ ----> NaHSO₄ + HI

• Bromide ions then reduce sulfuric acid to sulfur dioxide

  • Orange bromine vapour and choking gas are observed

  • 2HBr + H₂SO₄ ----> Br₂ + SO₂ + 2H₂O

  • Sulfur is reduced as its ON decreases from +6 to +4

  • Bromine is oxidised as ON increases from -1 to 0

Reaction of halides with sulfuric acid: iodide

• Iodide ions then reduce H₂SO₄ to SO₂

  • Violent iodine vapour and choking SO₂ gas are observed

  • 2HI + H₂SO₄ ----> I₂ + SO₂ + 2H₂O

• Iodide ions then reduce SO₂ further to produce H₂S gas

  • Violet iodine vapour and the rotten egg smell of H₂S gas are observed

  • 6HI + SO₂ ---> H₂S + 3I₂ + 2H₂O

• Sulfur is reduced as its ON decreases from +4 to -2

• Iodide is oxidised as ON increases from -1 to 0

Testing for halide ions

• Method

  • Add dilute nitric acid to sample to remove any ions

    • May interfere by also forming precipitates

  • Add silver nitrate to form silver halide precipitate

• Results

  • Chlorine - white

  • Bromide - cream

  • Iodide - yellow

• Adding excess dilute NH₃

  • Chloride - dissolves

  • Bromide - remains insoluble

  • Iodide - remains insoluble

• Adding concentrated NH₃

  • Chloride - dissolves

  • Bromide - dissolves

  • Iodide - remains insoluble

Reactions of hydrogen halides

• Reactions with water

  • Hydrogen halides are highly soluble in water, form acidic solutions

  • Reaction produces strong acids as the hydrogen halides fully dissociate in water

    • Resulting solutions will turn blue litmus paper red

  • When exposed to moist or water vapour, hydrogen halides form misty fumes due to the formation of tiny droplets of aqueous acid solution

• Reactions with ammonia

  • React with ammonia gas to form ammonium halide salts

    • Acid-base reaction producing white fumes

  • HX + NH₃ ---> NH₄Cl

    • Occur readily due to strong affinity between acidic hydrogen halides and basic ammonia molecule

Disproportionation reactions

Occur when a substance is simultaneously oxidised and reduced in same chemical reactions

Chlorine to disinfect drinking water

• When chlorine is bubbled through water, disproportionation occurs

  • Cl₂ + H₂O ----> HCl + HClO

• Hypochlorous acid dissociates

  • HClO ---> H+ + ClO-

• Hypochlorite ions act as a disinfectant which kills microorganisms

• In sunlight chlorine gas react with water to form HCl and oxygen gas

  • Repletes the hypochlorite ion disinfectant

Making bleach using disproportionation

• Mixing chlorine gas with cold dilute sodium hydroxide solution

  • Cl₂ + 2NaOH ---> NaClO + NaCl + H₂O

  • ON of chlorine increases from 0 to +1 in NaClO and decreases to -1 in NaCl

• Bleach contains chlorate ions that act as oxidising agents to kill bacteria

• If chlorine is reacted with hot sodium hydroxide, NaClO₃ is formed

  • 3Cl₂ + 6NaOH ----> NaClO₃ + 5NaCl + 3H₂O

  • ON of chlorine from 0 to +5 and decreases to -1