Unit 1.6 Periodic table

Question 1

How would u write the electron configuration based on s, d and p blocks?

Answer 1

Wherever an element is in the periodic table based on them s, d, p blocks, it means their last electron is within that orbital

e.g. silicon is in p block and has 14 electrons
therefore arrangement:
1s^2, 2s^2, 2p^6, 3s^2, 3p^2

Question 2

How to write down first ionisation energy?
e.g. for Na(g)

Answer 2

Na(g) -> Na+(g) + e^-

Question 3

Why does nuclear charge increase across the period?

Answer 3

Each element has an additional proton
technically
N° of protons increases

Question 4

Why does electron shielding remain the same across the period?

Answer 4

Electrons added to the same shell

Question 5

What’s the trend in ionisation energy increase across a period

Answer 5

It goes up?

Question 6

Tell me about the trend in electronegativity?

Answer 6

Increase across a period
- Nuclear charge increases
- Shielding remains constant

Decreases down a group
- Effective nuclear charge remains constant
- Shielding increases

Question 7

Define electronegativity

Answer 7

The ability of an atom within a covalent bond to pull the pair of electrons towards itself

Question 8

Explaining melting point trends for groups 1 & 2

Answer 8

Group 1 & 2:
- Decreases down the group due to decreased electrostatic attraction
- Atomic radii is greater
- Electron more weakly held
- Metallic bond is weaker
-> Less heat energy to overcome bonds
-> lower mp

Question 9

Explain melting point trends for group 7

Answer 9

Group 7:
- Increases down the group as VDW forces stronger between molecules
- More electrons causing movement of electrons
- Increased temporary dipoles
- Increased VDW forces
-> More heat energy to overcome bonds
-> Higher mp

• States change O_O

Question 10

State the state changes down group 7 elements

Answer 10

• Fluorine = gas
• Chlorine = gas
• Bromine = liquid
• Iodine = solid

Question 11

Explain the period 3 trend in melting & boiling points
(graph type shi)

Answer 11

Na, Mg, Al (rising):
- They all have metallic bonding
- Melting points increase due to increasingly positively charged ions
- More electrons released as delocalised electrons
- Therefore attractive electrostatic forces increase

Si (the highest):
- Apparently it’s a macromolecular
- Has a very strong covalent structure
- The covalent bonds need a lot of energy to break
- Therefore Si has a very high melting point

P, S, Cl, Ar (the lowest’s):
- P, S & Cl are simple covalent
- They have weak Van Der Waals forces
- Intermolecular forces don’t need much energy to overcome
- Therefore low melting points
- Tho, melting point decreases from S to Ar
- Due to decreasing size of the molecules

For argon:
- It’s a noble gas
- Has individual atoms with a full outer shell of electrons
- Atom is very stable
- Van Der Waals forces between them very weak
- Soooo, melting point very low
- It’s just a gas at room temperature

Question 12

What do u call elements along the border line

Answer 12

Metalloids
Properties of both metal & non-metals

Question 13

What are redox reactions?

Answer 13

Remember GCSE?
The use of oil rig:
- Reduction is gain of electrons
- Oxidation is loss of electrons

Question 14

What are agents in terms of redox?

Answer 14

• Oxidising agent causes oxidations of another species and are themselves reduced
• Reducing agent causes reduction of another species and are themselves oxidised

Question 15

How to tell if it’s a redox reaction?

Answer 15

How to tell if it’s a redox reaction?:
- If the oxidation number changes, it’s a redox
e.g.
CH4 + O2 -> CO2 + H2O

Notice the oxygen before and after
If an element isn’t combined with anything, it’s oxidation number is 0
But now that it has combined, the oxidation number changed
Get it?

Question 16

How to identify the agent?

Answer 16

Icl, use ur knowledge D:
(that’s what he said, I’ll see what I can do)

Question 17

What’s the reaction called for if an element is both oxidised and reduced

Answer 17

Disproportionation reaction

Question 18

What are group 2 elements?

Answer 18

• They are called alkaline (earth) metals
• Reactivity increases down the group because:
  Nuclear charge increases
  Level of shielding increases
  Distance of 1st electron from nucleus increases
  Increases reactivity down the group
• In addition, they always lose 2+ electrons to form 2+ ions every reaction
• To achieve a full outer shell

Question 19

Explain to me about group 2 elements reacting with water

Answer 19

• Produces a redox reaction
• Producing metal hydroxide
  e.g. Mg + 2H2O -> Mg(OH)₂ + H₂
  OIL occurs on Magnesium, and it’s the same with other group 2 elements
• Magnesium reacts very slowly with liquid water
• Therefore they use steam, which also provides extra energy
• During reaction, magnesium burns with a white flame
• Forming hydrogen & magnesium oxide, a white powder

Question 20

Explain to me about group 2 elements reacting with oxygen

Answer 20

• Reacts with pure oxygen
• Producing simple oxides
• But they would always have an oxidation number of 2+
• to balance with oxygen (-2)
  e.g. 2Mg + O2 -> 2MgO
• Barium is an exception to this
• Forms barium peroxide
• Oxygen has -1 oxidation number
  e.g. Ba + O2 -> BaO2

Question 21

Explain to me about group 2 elements reacting with hydroxides

Answer 21

• Reacts with OH^- ions
• Forms hydroxides
• When writing formula, 2 hydroxide ions must be bracketed
  e.g. Mg²⁺ + 2OH^- -> Mg(OH)₂
• Solubility of group 2 hydroxides increases down the group
• Mg(OH)₂ = least soluble, Ba(OH)₂ = most soluble
• ∴, Magnesium hydroxide used in medicine as an antacid
• as it’s alkaline and can neutralise acids
• Or in agriculture to neutralise acidic soil

Question 22

Explain to me about group 2 elements reacting with sulfates

Answer 22

• Reacts with SO₄^2- ions
• Also forms hydroxides
• The difference is solubility decreases down the group
• and no brackets
• MgSO4 = most soluble, BaSO4 = least soluble
  e.g. Ba²⁺ + SO₄^2- = BaSO₄
• Barium sulfate used in medicine as barium meals
• as a medical tracer, imaging internal tissues or organs
• Although the substance is toxic, it’s insoluble
• Therefore can’t be absorbed in bloodstream and is safe

Question 23

What are the flame tests with group 2 elements?

Answer 23

Wait…. no need

Question 24

How are carbonates formed with group 2 salts?
_{(thermal decomposition)}

Answer 24

• Carbonates are formed by:
  Soluble group 2 salts
  +
  A solution containing carbonate ions (CO₃^2-)
  e.g. Ca(NO₃)₂ aq + Na₂CO₃ aq -> CaCO₃ s + 2NaNO₃ aq
  tho ionic equation:
  Ca²⁺ (aq) + CO₃^2- (aq) -> CaCO₃ (s)
• However, they undergo thermal decomposition
  (Splitting up a compound by heating it)
• To give the metal oxide & carbon dioxide gas
• All carbonates are white solids
• Same with the oxides

Question 25

Explain thermal stability of carbonates down the group

Answer 25

• As you go down the group, carbonates needs to be heated more strongly to decompose (thermal stability increases)
• Due to a smaller cation having a greater charge density
• Polarising the CO₃^2- ion more effectively
• Resulting the metal oxide to form
  e.g. CaCO₃ (s) △-> CaO(s) + CO₂ (g)
  △ = heat
  Better way:
• Down the group, metal ion increases in size
• Charge has a less polarising effect on carbonate ion
• Harder for carbonate to decompose
• Therefore need stronger heat

Question 26

Explain thermal stability for group 2 elements with hydroxides?

Answer 26

• Thermal stability also increases down the group
• And they also form oxides during thermal decomposition too
• Same reason to why it increases too
  e.g. Mg(OH)2 (s) △-> MgO(s) + H2O (g)

Question 27

Tell me about the basic character of oxides & hydroxides

Answer 27

• We know group 2 metals can react with oxygen to form metal oxides
• And react with water to form group 2 hydroxides
• They can also form group 2 hydroxide from:
  Soluble group 2 salts
  +
  Solution containing hydroxide ions (OH^-)
  e.g. MgNO₃(aq) + NaOH(aq) -> Mg(OH)₂(s) + NaNO₃(aq)
  the ionic equation:
  Mg²⁺(aq) + 2OH^-(aq) -> Mg(OH)₂ (s)
• Once made, group 2 oxides & hydroxides have basic character
  (They can react with acids to form salt and water)
  Group 2 oxides = bases
  Group 2 hydroxides = alkalis (soluble in water)
• S-block oxides are all basic and react
  Group 1 = Na₂O(s) + H₂SO₄(aq) -> Na₂SO₄(aq) + H₂O(l)
  Group 2 = CaO(s) + 2HCl(aq) -> CaCl₂(aq) + H₂O(l)
  same with any other metal…
• S-block hydroxides are also basic and react
• Those which are soluble in water form alkaline solutions
  Group 1 = NaOH(aq) + HCl(aq) -> NaCl(aq) + H₂O
  Group 2 = Ca(OH)₂ (aq) + H₂SO₄ (aq) -> CaSO₄ (aq) + H₂O
  same with any other metal too…

Same with the balanced equations above, should know how to write them with any group 2 oxide/hydroxide with common acids:
Hydrochloric acid = HCl
Sulfuric acid = H₂SO₄
Phosphoric acid = H₃PO₄

Question 28

How are group 1 oxides/hydroxides stronger bases than group 2?

Answer 28

• Group 2 ions have 2+ charges
• Larger charge = greater force of attraction in compound
• Harder for hydroxide/oxide ions to be released

Question 29

What are group 7 elements?

Answer 29

• They are called Halogens
• Highly reactive non-metals (they needa gain 1 electron to form a 1- ion)
• Volatility decrease down the group (tendency of a substance to vaporize)
• In the same way, reactivity decreases & melting point increases down the group:
  More energy needed to break bonds as
  Van deer waal’s forces increases in strength as
  More electrons ∴
  Greater wobble ∴
  Larger temporary dipoles ∴
  larger van deer waal’s force with neighbour ∴
  More heat energy needed ∴
  Higher M.P/B.P

Question 30

How are halogens diatomic molecules? (X₂)

Answer 30

7 electrons in outer shell so they react with other atoms to complete their shell

Question 31

Although their covalent bond strong, why mp so low for fluorine?

Answer 31

Weak intermolecular forces (vdw’s) between molecules

Question 32

Explain to me about halogens with metals

Answer 32

• Salt makers
• Halogen + metal = ionic salts called metal halides
• Halogen causes oxidation of metal; halogen itself being reduced
  e.g. 2Na(s) + Cl₂(g) -> 2NaCl(s)
  Half equations:
  Oxidation of sodium = 2Na -> 2Na⁺ + 2e^-
  Reduction of chlorine = Cl₂ + 2e^- -> 2Cl^-
• All reactions = exothermic
• Fluorine w/ metals = typically very violent
• All of group 1 & 2 metals react with halogens
• Forming white ionic salts
  e.g. Ca(s) + Br₂(g) -> CaBr₂(s)
• In addition, transition metals can react with halogens to form coloured salts
  e.g. Ba²⁺(s) + Cl₂^- -> BaCl₂(s)
  Half equations:
  Oxidation = Ba -> Ba²⁺ + 2e^-
  Reduction = Cl₂ + 2e^- -> 2Cl^-

Question 33

Explain to me about halogens and their oxidising power

Answer 33

• Good oxidising agents
• They accept electrons from species being oxidised
• and are reduced as a result
• Decreases down the group
• Ability to attract electrons decrease cuz:
  Shielding increases
  Greater atomic radius
  (distance between nucleus and outer shell increases)

Question 34

What are halogens states, colour & reaction w/ iron wool?

Answer 34

Fluorine = gas = pale yellow = burns with cold iron wool
Chlorine = gas = greenish-yellow = burns vigorously with hot iron wool
Bromine = liquid = red-brown/vapour red brown = burns quickly with hot iron wool
Iodine = solid = lustrous grey-black/vapour purple = reacts slowly w/ hot iron wool producing iodine vapours)

Question 35

The origins for testing for aqueous halide ions

Answer 35

• There’s already summin made for this
• But not the origins soooo:
• Acidified silver nitrate contains aqueous Ag⁺ ions
• Used to test for halide ions
  e.g. NaCl_(aq) + AgNO_3(aq) -> AgCl_(s) + NaNO_3(aq)
  Ionic equation:
  Cl^-(aq) + Ag⁺(aq) -> AgCl(s) spectator ions not included
  It’s the same with any other halide ions so be aware

Question 36

Explain to me about the displacement reactions of the halogens?

Answer 36

• So we know that halogens are generally oxidising agents gaining electrons to form corresponding halide ions
• But as reactivity decreases down the group, a more reactive halogen (e.g. F₂) will oxidise the halide ion of a less reactive halogen
  For example:
  When chlorine gas/chlorine water added to aqueous potassium bromide, red brown colouration of bromine is observed.
  Half equations:
  Cl₂(g) + 2e^- -> 2Cl^-
  2Br^-(aq) -> Br₂ + 2e^-
  Overall ionic equation:
  Cl₂(g) + 2Br^-(aq) -> Br₂(l) + 2Cl^-(aq)
  Full reaction equation w/ spectator ions; those species in solution don’t undergo any redox changes:
  Cl₂(g) + 2KBr^-(aq) -> Br₂(l) + 2KCl^-(aq)
• It’s actually just the same with any other halogens yano?
• These reactions are often called displacement reactions
• Examples of redox reactions
• Each case, halogen = gained electrons to become halide ion
• & has been reduced
• And the aqueous halide ion has lost an electron
• & has been oxidised
• VOILA! redox reaction

Question 37

What are the relative oxidising strength of the halogens?

Answer 37

• A halogen will displace any halide beneath it in the periodic table:
  Cl₂ = displace Br^- & I^- ions
  Br₂ = displace I^- ions
  I₂ = won’t displace any halide ions ;P

Question 38

Tell me about chlorine & fluorine in water treatment

Answer 38

Chlorination in water treatment:
- Used as a disinfectant to kill pathogens such as:
bacteria & viruses
- Doesn’t kill all pathogens but reduces them to a level
- Which is considered to be safe for human consumption
- Combined with filtration, chlorination = excellent method of disinfecting drinking water supplies

• Ethical objections to chlorination
• As it results in the formation of very low levels of toxic by-products
• But no evidence for such by-products having negative health effects in humans

In addition with chlorination, it’s actually a disproportionation reaction

Question 39

Tell me about fluorine in water treatment

Answer 39

Fluorination in water treatment:
- Controlled addition of fluoride to drinking water
- To reduce tooth decay
- Low levels of fluoride in saliva reduce rate of tooth enamel demineralisation
- Leading to the formation of fewer tooth cavities

• Countries where fluoridation has been introduced….
• Rates of tooth decay fallen dramatically
• But, fluoridation considered unethical by some ppl
• Constitutes compulsory medication without informed consent
• HOWEVER, no scientific evidence of health risks associated
• With water fluoridation at the low fluoride concentration required
• To maintain dental health

Question 40

Tell me about soluble & insoluble salts

Answer 40

• Formed by reacting acids with soluble/insoluble bases
• However, soluble salt can be separated (“recovered”) from a solution by crystallisation[/evaporation, not gonna write about it….]:

Crystallisation:
NaOH(aq) + HCl(aq) -> NaCl(aq) + H₂O(l)
- Salt solution slowly heated in an evaporating dish
Equipment: bunsen burner, tripod, gauze, evaporating dish
- Evaporates some of the solvent
- Salt crystals start to form which then
- Heat is removed and solution left to cool
- Once crystals are formed, crystals are filtered out of the solution and left to dry in a warm place (oven/desiccator)

Question 41

Tell me about insoluble salts

Answer 41

• Formed by a precipitation reaction
  Occurs when 2 solutions containing different salts are reacted together

Precipitation:
Pb(NO₃)₂(aq) + 2KI(aq) -> PbI₂(s) + 2KNO₃(aq)
Equipment: filter funnel, filter paper, conical flask
Ig the lead iodide residue found inbetween filter paper & funnel
- Lead iodide residue would need to be washed with water
- To remove any potassium nitrate impurity
- The washed solid residue will need further drying (oven/desiccator)
- May be used in a gravimetric analysis afterwards

Gravimetric Analysis:
- Used to determine mass percentage of an ion in an impure compound which is normally soluble in water
Unfortunately we’ll use a different example:
- Gravimetric Determination of Chloride in Table salt
1. Completely dissolve a known mass of table salt in water
2. Add a precipitating agent
- This a species which precipitates the analyte complete
- E.g. silver nitrate precipitates chloride ions as
- Solid silver chloride AgCl
- Important that the precipitating agent only causes analyte ion to precipitate
- If table salt was suspected of containing iodide impurities
- Silver nitrate = poor choice of precipitating agent
- Will cause iodide ions to precipitate as solid silver iodide
(So basically, just choose one that will only affect the analyte specifically)
3. Filter and dry the precipitate
4. Weigh the precipitate

Number of moles of precipitate can be calculated and used to determine the mass percentage of chloride in the table salt sample
(I assume that’s where mass/mr = n comes in)