AS inorganic

describe how we can use dilute NaOH to identify group 2 cations:

• add 6 drops of dilute NaOH solution to test tube containing metal ion solution (e.g. magnesium sulfate, Mg²⁺ etc.) and observe precipitate (if any) formed

• keep adding dilute NaOH until it is in excess - record any changes you see

give 3 ways we can identify group 2 metal cations:

• flame test (not in RP4)

• by adding dilute NaOH

• by adding H₂SO₄

what is seen for each of the group 2 metal cations when dilute NaOH is added?

what would be the initial observation for a group 2 metal compound?

colourless solution

how can we test for ammonium (NH₄⁺) ions?

• add substance you are testing to test tube

• add NaOH solution and shake

• warm gently using a water bath

• ammonia should be released - test fumes released by holding damp litmus paper at mouth of test tube

• if ammonium ions present, damp litmus paper turns from red → blue

give the ionic eqn for the ammonium ion test:

NH₄⁺ _(aq) + OH^- _(aq) → NH_{3 (g)} + H₂O _(l)

describe how we can use H₂SO₄ to identify group 2 metal cations:

• add 10 drops of H₂SO₄ solution to test tube containing metal ion solution (e.g. magnesium sulfate, Mg²⁺ etc.) and observe precipitate (if any) formed

• keep adding H₂SO₄ until it is in excess - record any changes you see

what is seen for each of the group 2 metal ions when H₂SO₄ is added?

give 2 ways we can test for OH^- ions:

• using litmus paper

• using ammonia solution and litmus paper

how can we test for OH^- ions with just red litmus paper?

• dip a piece of red litmus paper into the solution

• if OH^- ions are present, the paper will turn blue

how can we confirm a solution is ammonia based on the test for OH^- ions?

• take 5 drops of solution you are testing and place on filter paper

• place this inside a petri dish w/ a lid

• dampen a piece of red litmus paper w/ deionised water and place on other side of petri dish

• replace lid and observe over a few minutes - if OH^- ions present, litmus paper should turn blue as ammonia vapours produced

how does the test for OH^- ions and litmus paper work?

• we add deionised water as OH^- ions form when NH₃ comes into contact w/ water

• this turns red litmus paper blue as ammonia solution vapours are alkaline

how can we test for CO₃^2- ions in aqueous solution?

• add 2 cm³ Ca(OH)₂ (limewater) to a test tube

• to a different test tube, add 3 cm³ substance you are testing and an equal volume of dilute HCl

• immediately place a delivery tube from this to the limewater test tube

• if CO₃^2- ions present, limewater goes cloudy

explain how HCl can help detect CO₃^2- ions and give the eqn:

• limewater turns cloudy if carbonate ions present as CO₂ is produced in the reaction w/ dilute HCl (H⁺ ions)

• CO₃^2-_(aq) + 2H⁺ _(aq) → CO_{2 (g)} + H₂O _(l)

how can we test for sulfate ions in aqueous solution?

• add dilute HCl and acidified BaCl₂

• white precipitate of BaSO₄ forms if SO₄^2- ions present

• (BaCl₂ acidified by HCl but is in MS)

give the ionic eqn for the test for sulfate ions in aqueous solution:

Ba²⁺ _(aq) + SO₄^2- _(aq) → BaSO_{4 (s)}

how can we test for halide ions in aqueous solution?

• add 10 drops of substance you are testing to a clean, dry test tube

• add approx 5 drops of dilute nitric acid and shake

• add 10 drops silver nitrate solution and record observations

• add dilute ammonia (or ammonia solution if testing for I^- ions and if so, work in a fume cupboard) and record further observations

why do we add nitric acid during the test for halide ions?

to remove CO₃^2- and OH^- ions, which would also form precipitates and interfere w/ the test

give the general eqn for the reaction of halide ions w/ silver nitrate:

Ag⁺ _(aq) + X^- _(aq) → AgX _(s)

give the results of the silver nitrate test:

AgCl - white ppt

AgBr - cream ppt

AgI - yellow ppt

give the results of the silver nitrate halide ions test when aqueous ammonia/ammonia solution is added:

AgCl - redissolves

AgBr - redissolves slowly and needs a lot of ammonia

AgI - does not dissolve

how can we test for halide ions in solid salts?

carry out in a fume cupboard and wear gloves:

• place a small spatula of the solid you are testing into a clean, dry test tube

• slowly add a few drops of concentrated H₂SO₄ and record observations

• test gas w/ moist blue litmus paper and record observations

• (repeat w/ acidified K₂Cr₂O₇ / lead (II) nitrate solution)

give the observations for the tests for halide ions in solid salts:

(not part of the practical but useful!) how can we test for all acids? what denotes a +ve result?

• add Na₂CO₃

• +ve result = effervescence - as CO₂ produced

(not in RP4 but useful to know) identify a reagent/test that could be used to distinguish between aqueous solns of SO2 and SO3 w/ the same concs:

use universal indicator:

• SO₂: orange-red (higher pH)

• SO₃: red (lower pH)

how does the reactivity of the group 2 elements w/ water change going down the group?

increases

does Be react w/ water? if so, how? + give the eqn

no :(

does Mg react w/ water? if so, how? + give the eqn

• reacts very slowly w/ water but readily w/ steam

• water: Mg + 2H₂O _(l) → Mg(OH)₂ + H₂

• steam: Mg + H₂O (g) → MgO + H₂

state two observations you would make when magnesium reacts w/ steam:

• bright white light

• white solid (NOT ppt)

does Ca react w/ water? if so, how? + give the eqn

• reacts steadily w/ water

• Ca + 2H₂O → Ca(OH)₂ + H₂

does Sr react w/ water? if so, how? + give the eqn

• reacts fairly quickly w/ water

• Sr + 2H₂O → Sr(OH)₂ + H₂

does Ba react w/ water? if so, how? + give the eqn

• reacts rapidly w/ water

• Ba + 2H₂O → Ba(OH)₂ + H₂

how does solubility of group 2 compounds change down the group?

• compounds of group 2 w/ singly charged anions - OH^- - increase in solubility down the group

• compounds of group 2 w/ doubly charged anions - SO₄^2- - decrease in solubility down the group

is BaSO₄ soluble in water? why is this significant?

• not soluble in water

• precipitates out of solution so forms a white precipitate when BaCl₂ and dilute HCl are added if SO₄^2- ions are present

summarise the test for sulfate ions:

• add dilute HCl and BaCl₂

• white precipitate of BaSO₄ forms if SO₄^2- ions present

why is HCl added in the test for sulfate ions?

to remove carbonates and sulfites, as they would also produce a white precipitate

how is Ti extracted from its ore? + give the eqn

• Mg is used to extract Ti from rutile

• TiO₂ is first converted to TiCl₄, which is then reduced by Mg in a furnace at 1000^oC

• TiCl₄ + 2Mg → Ti + 2MgCl₂

name and describe a use of BaSO₄:

barium meals:

• BaSO₄ is opaque to X-rays

• patients swallow a barium meal (suspension of BaSO₄) which coats the tissues so they show up on an X ray

• BaSO₄ is insoluble so passes through your body after

how is sulfur removed from flue gases? + give the eqns

flue gas desulfurisation:

• powdered CaO or CaCO₃ mixed w/ water to make an alkaline slurry

• when flue gases mix w/ alkaline slurry, acidic SO₂ gas reacts w/ calcium compounds to form CaSO₃, which is harmless

• CaO + 2H₂O + SO₂ → CaSO₃ + 2H₂O

• CaCO₃ + 2H₂O + SO₂ → CaSO₃ + 2H₂O + CO₂

give 3 general uses of group 2 compounds:

• to neutralise acids

• antacids (indigestion remedies)

• to neutralise acidic soils

give the general ionic eqn for the reaction of a group 2 metal w/ water:

X²⁺ _(s) + 2H₂O _(l) → Ca²⁺ _(aq) + 2OH^- _(aq) + H_{2 (g)}

state the role of water when reacting w/ group 2 metals:

oxidising agent

explain why different observations are made when aq BaCl₂ is added separately to aq MgSO₄ and aq Mg(NO₃)₂:

BaSO₄ is insoluble but Ba(NO₃)₂ is soluble

solutions of barium hydroxide are used in the titration of weak acids.

state why magnesium hydroxide solution could not be used for this purpose (1)

Mg(OH)₂ is insoluble

how does atomic radius change down group 2?

• increases

• as more e^- shells are added

• so size increases

how does 1st IE change down group 2? why?

• decreases

• outer e^- in higher E level / increase in shielding / atoms larger / more shells

• ∴ weaker attraction between ion and lost e^-

• (there is more +ve charge in the nucleus but this is overridden by the extra shells)

how does reactivity change down group 2? why?

• increases

• as 1st and 2nd IE decrease down the group - it is easier to lose e^- and form 2+ ions as you go down

what is the general trend of mpt down group 2?

• generally decreases

• as the atoms (and ions) get bigger as you go down the group, delocalised e^- are further from the +ve nuclei

• so it takes less energy to break metallic bonds

• Mg does not follow this trend (because of how its ions are arranged)

how reactive are the halogens in general?

fairly reactive

how does reactivity change down the group? why?

• decreases down the group

• as halogens want to gain e^- to become -ve ions so reactivity decreases down the group as shielding increases down the group

how does the bpt of the halogens change down the group? why?

• increases down the group

• as molecules get larger

• so VDWs between molecules get stronger

how does electronegativity change down group 7? why?

• decreases down group

• as nucleus is more shielded

• so weaker attraction between nucleus and bonding pair

what is a disproportionation reaction?

reaction where a single element is both oxidised and reduced

give the full and ionic eqn for making bleach. what type of reaction is this? what observations are made?

2NaOH + Cl₂ → NaClO + NaCl + H₂O

ionic: 2OH⁻ + Cl₂ → Cl⁻ + ClO⁻ + H₂O

disproportionation, yellow green gas Cl → colourless solution

give the full and ionic eqn for the reaction of chlorine w/ water. what type of reaction is this?

ionic: Cl₂ + H₂O ⇌ 2H⁺ + Cl^- + ClO^-

Cl₂ + H₂O ⇌ HCl + HClO

disproportionation

give the eqns for the reaction of chlorine w/ water IN SUNLIGHT:

2Cl₂ + 2H₂O ⇌ 4HCl + O₂

2HClO → 2HCl + O₂

why is Cl added to water, despite its limitations?

• ClO^- ions kill bacteria → safe to swim in/drink

• only a small amount is added/benefits outweigh risks

give the strengths of adding Cl to water:

• kills disease causing microorganisms

• prevents algae growth

• removes discolouration

give the limitations of adding Cl to water:

• Cl gas is harmful

• liquid Cl is corrosive

• Cl can react w/ organic compounds in water to make chlorinated compounds (which can be cancerous)

describe a halogen displacement reaction and give the eqn for the displacement of Br by Cl₂ in KBr:

• a solution of a more reactive halogen can displace a less reactive halide from its salt

• 2KBr + Cl₂ → 2KCl + Br₂

how does oxidising ability of halide ions change down the group?

• attraction between outer e^- and nucleus decreases down group as ions get bigger

• so oxidising ability decreases

give the eqns for halides reacting w/ concentrated sulfuric acid:

NaX + H₂SO₄ → NaHSO₄ + HX

give the further reaction for some halides reducing sulfuric acid:

2HX + H₂SO₄ → X₂ + SO₂ + 2H₂O

give the overall eqn for halides reacting w/ concentrated sulfuric acid and reducing it further:

2NaX + 3H₂SO₄ → 2NaHSO₄ + X₂ + SO₂ + 2H₂O

give the half eqns for some halides reducing sulfuric acid:

2X^- → X₂ + 2e^-

H₂SO₄ + 2H⁺ +2e^- → SO₂ + 2H₂O

name and explain the observations for the reactions of NaF and NaCl w/ H₂SO₄:

• HF and HCl are not strong enough reducing agents to reduce the H₂SO₄ further

• not a redox reaction - oxidation states of halogens and sulfur remain the same

• observation: misty white fumes of HF / HCl

name and explain the observations for the reaction of NaBr w/ H₂SO₄:

• HBr is a strong enough reducing agent to reduce H₂SO₄ further

• second eqn = redox: Br oxidised (from -1 to 0) and S reduced from (+6 to +4)

• observations: misty fumes of HBr, red-brown vapour Br₂, choking fumes of SO₂

what are the overall and ionic eqns for the overall reaction of sodium bromide w/ sulfuric acid, w/ the sulfuric acid being further reduced?

ionic: 2Br⁻ + 3SO₄^2- + 6H⁺ → 2HSO₄^- + Br₂ + SO₂ + 2H₂O

overall: 2NaBr + 3H₂SO₄ → 2NaHSO₄ + Br₂ + SO₂ + 2H₂O

give and explain the eqns for NaI reacting w/ sulfuric acid:

NaI_(s) + H₂SO_{4 (l)} → NaHSO_{4 (s)} + HI_(g)

HI can reduce H₂SO₄ further: 2HI_(aq) + H₂SO_{4 (l)} → I_{2 (s)} + SO_{2 (g)} + 2H₂O_(l)

HI is a very strong reducing agent and can reduce the SO₂ again to S and H₂S: 6HI_(g) + H₂SO_4(l) → 3I_2(s) + S_(s) + 4H₂O_(l)

6HI_(g) + SO_2(g)→ H₂S_(g) + 3I_2(s) + 2H₂O _(l)

give and explain the observations for NaI reacting w/ sulfuric acid:

• misty white fumes (HI)

• purple vapour (I₂)

• yellow solid formed (S)

• rotten egg smell (H₂S)

• black solid formed (I₂)

• choking fumes (SO₂)

give the overall and overall ionic eqns for NaI reacting w/ H₂SO₄ and the observations^_:

overall: 8NaI + 9H₂SO₄ → 4I₂ + 8NaHSO₄ + H₂S + 4H₂O
ionic: 8I⁻ + SO₄^2- + 10H⁺ → 4I₂ + H₂S + 4H₂O

observations: purple vapour released, black solid formed and a rotten egg smell

how can we test for halide ions in aqueous solution?

• add 10 drops of substance you are testing to a clean, dry test tube

• add approx 5 drops of dilute nitric acid and shake

• add 10 drops silver nitrate solution and record observations

• add dilute then concentrated ammonia and record further observations

give the general eqn for the reaction of halide ions w/ silver nitrate:

Ag⁺ (aq) + X^- (aq) → AgX (s)

give the results of the silver nitrate test:

AgCl - white ppt

AgBr - cream ppt

AgI - yellow ppt

give the results for the silver halide ppts dissolving in ammonia:

AgCl - dissolves in dilute ammonia to form a colourless solution

AgBr - dissolves in concentrated ammonia to form a colourless solution

AgI - does not dissolve

how do we classify elements? how is this determined?

• according to their block - s, p, d or f

• block is determined by proton no. (or last orbital in e^- configuration)

how does atomic radius change across a period? why?

• decreases

• as no. of protons increases, +ve charge of nucleus increases

• this means outer e^- are pulled closer, making radius smaller

• w/ same shielding (as even though we are adding extra e^-, they are added to the same shell, so they don’t increase the size or shield the +ve charge of the nucleus)

  

what is first ionisation energy?

minimum energy needed to remove 1 mole of e^- from 1 mole of atoms in the gaseous state

what is the general trend of ionisation energy across a period? why?

• generally increases - takes more energy to remove an e^-

• as no. of protons increases, so +ve charge of nucleus increases

• so pulling force on outer e^- increases

  

why may first IE sometimes decrease across a period?

an e^- may be being removed from a different subshell, so requires less energy to be removed

how does the mpt change across period 3 from Na → Al? why?

• there is a general increase from Na → Al as metallic bonding gets stronger, so more energy is required to break the metallic bonds

• this is because we go from Na⁺ to Mg²⁺ to Al³⁺ so there is a smaller radius and more delocalised e^-

• so there is a stronger attraction between the +ve ions and the delocalised e^-

how does the mpt change across period 3 when we get to Si? why?

• increase - has highest mpt

• as Si is a macromolecular compound w/ very strong covalent bonds

• which require a large amount of energy to break

how does the mpt change across period 3 from P → Ar? why?

• P₄, S₈ and Cl₂ are all molecular substances, so their mpt depends on the strength of the IMF - in this case VDW - between the molecules

• the more atoms, the stronger the IMF, so S₈ > P₄ > Cl₂

• Ar is monatomic so has a very low melting point

which ion has the largest radius and why? F^-, Mg²⁺, Na⁺, O^2-

O^2- - most -ve, fewest protons so weakest attraction between outer e^- and nucleus