3. elements from the sea

electrolysis

electrolysis

process of passing an electrical current through a molten or solution of a ionic substance and it splitting into its ions

electrolyte

liquid required to conduct electricity in electrolysis, made up of free ions

anion

negative ion

cation

positive ion

anode

positive electrode

cathode

negative electrode

4, 2, 5, 1, 3

1. turn power supply on

2. use inert electrodes (eg platinum or carbon) so they don’t react

3. depending electrolyte products will form as metals (plating) or gases (bubbles)

4. use wires and clips to connect the electrode to power supply

5. place electrodes into beaker containing electrolyte (make sure they don’t touch)

half equation

equation showing movement of electrons

anode half equation

negative ions losing electrons forming atoms

cathode half equation

positive ions gaining electrons to form atoms

ions produced by a molten salt

only ions available are those in the substance

addition ions in aqueous solution

H+ and OH- ions from water

metal formed

if at the cathode the metal is less reactive than hydrogen eg silver or copper

hydrogen gas formed

if at cathode the metal is more reactive than hydrogen eg group 1 and 2 metals plus aluminium

oxygen formed

if the solution at the cathode doesn’t contain halide

halogen formed

if the solution at anode is concentrated and contains a halide

oxygen formed

if the solution at the cathode contains a dilute halide

causes the anode to shrink

due to the anode been made from an impure metal and the cathode made from pure metal

electrolysis of halide solutions

process to extract halogens

brine

solution of water with high concentration of salts, mainly sodium chloride but some bromine and iodine salts as well

chloride

halogen made by industrial electrolysis of brine

electrolysis of brine

cathode 2 H+ form H2 by accepting 2e-

anode 2 Cl- form Cl2 y gaining 2e-

sodium ions stay in solution as less reactive than hydrogen, they eventually form sodium hydroxide with OH

chlorine production

electrolysis of brine

inert electrode eg carbon or platinum

constant stream of brine

chlorine collected as gas

only extract from concentrated sodium chloride

bromine extraction

displacement reaction where more reactive chlorine gas is bubbled though brine and displaces the halogen product which is then collected, condensed into a liquid and purified.

iodine extraction

displacement reaction where more reactive chlorine gas is bubbled though brine and displaces the halogen product which is then collected and condensed into a grey solid

oxidation states

how many electrons an atom has donated or accepted to from an ion or bond

uncombined elements

oxidation state of 0

identical atoms

where atoms combine to always have an oxidation state of 0

monatomic ion

oxidation state same as charge

compound ions (ions with multiple atoms)

oxidation state is sum of atoms which is also the same as overall charge

neutral compound

oxidation state of 0

oxygen

oxidation state of nearly always -2

hydrogen

oxidation state is nearly always +1

roman numerals

used if a element has multiple oxidation states

-ate

used at end of ion name meaning it contains oxygen and another element

front of ion name

positive ion name

end of ions name

negative ion name

nitrate (V)

NO3-

sulfate (VI)

(SO4)2-

carbonate

(CO3)2-

manganate VII

(MnO4)-

hydroxide

OH-

ammonium

(NH4)+

hydrogencarbonate

(HCO3)-

sulfide

S2-

redox reaction

a reaction where electrons are transferred by reduction and oxidation happen simultaneously

oxidation

loss of electrons

reduction

gain of electrons

electrons lost

oxidation state increase- oxidation

electron gained

oxidation state decreases- reduction

half equations

equations which show whats been reduced and whats oxidised

oxidised half equation

element to ion and electron

reduction half equation

electron and electrons to ion

reducing agent

themselves are oxidised by accepting electrons

oxidising agents

themselves are reduced by donating electrons

balance charges

needed to be done in order to balance redox reactions, done using oxidation states

iodine sodium thiosulfate titrations

titration useful for working out concentration of oxidising agent

oxidise iodine with oxidising agent

25cm3 of potassium iodate

excess potassium iodide solution

forms iodine

find how many moles iodine has produced

titrate the iodine solution with sodium thiosulfate

when iodine solution fades to pale yellow add starch

end point is when it goes blue black to colourless

calculate the concentration of oxidising agent

1 mole of iodate V ions : 3 moles of iodine

iodate V moles is the same as potassium iodate V

causes inaccurate titrations

contaminated apparatus

no reading burette correctly

not using concordant results

fresh solutions

decreases down group 7

become less volatile (less easy to vaporise) due to increasing strength of instantaneous dipole-iduced dipole bonds

fluorines appearance at RTP

pale yellow gas

chlorine appearance at RTP

yellow green gas

bromine appearance at RTP

red brown liquid

iodine appearance at RTP

shiny grey solid

halogens natural state

covalent diatomic molecules

halogen solubility

due to been polar and covalent they have low solubility in water but dissolve in organic solvents like hexane

chlorine colour in water and hexane

virtually colourless

bromine colour in water

yellow/orange

bromine colour in hexane

orange/red

iodine colour in water

brown

iodine colour in hexane

pink/violet

halogen get less reactive down group

due to been larger so electrons are more shielded from proton so gain an electron less easily

halogen are oxidising agents

due to gaining an electron in p sub shell and are reduced

displacement reaction

halogens are able to swap with halide ions

more reactive halogen displaces less reactive halogen

chlorine is able to displace bromide and iodide

iodine is not able to displace any thing

hydrogen halide

halogen + hydrogen

hydrogen halide production

made by adding concentrated acid to a solid ionic halide eg concentrated phosphoric acid and sodium chloride

hydrogen chloride

can be made with concentrated sulphuric acid which get involved in redox reactions as it is a oxidising agent, cannot be used to make other hydrogen halides as the halide reduces sulfuric acid (H2SO4) to sulphur dioxide (SO2) or hydrogen sulfide (H2S)

won’t split with heat

hydrogen flouride ad hydrgen chloride

slipt slighty with heat

hydrogen bromide and iodide

thermal stability of hydrogen halides

decreases down group as they get larger meaning bonding electrons are further away from nucleus and are more shielded

acidic

hydrogen halides dissociate in water into their ions

ammonia halide

formed when ammonia accepts an proton from hydrogen halide causing it to gain both

hydrogen bromide and hydrogen iodide react

react with sulfuric acid to produce either SO2 or H2S

silver ions

used as a test for halides

add dilute nitric acid

add silver nitrate solution

forms precipitate

no precipitate

colour of precipitate formed by fluoride

white precipitate

colour of precipitate formed by chloride

cream precipitate

colour of precipitate formed by bromide

yellow precipitate

colour of precipitate formed by iodide

ammonia solution

added after silver ion test to help differentiate halides

dissolves to colourless

silver chloride when ammonia solution added as well

unchanged if dilute by colourless if concentrated

silver bromide when ammonia solution added

doens’t dissolve

silver iodide when ammonia solution added

chlorine gas

toxic, corrosive, increases fire risk, makes it difficult to transport so typically transported as a liquid under pressure

chlorine uses

water treatment (sterilises it), kills microorganisms, bleach

atom economy

useful product over sum off all products z 100

dynamic equilibrium

when rate of forward reaction equals reverse reaction, only in closed system

more reactants

equilibrium lies to left