Unit 1.6 Periodic table Flashcards

Question 1

Q

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

Answer

A

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

Q

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

Answer

A

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

Question 3

Q

Why does nuclear charge increase across the period?

Answer

A

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

Question 4

Q

Why does electron shielding remain the same across the period?

Answer

A

Electrons added to the same shell

Question 5

Q

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

Answer

A

It goes up?

Question 6

Q

Tell me about the trend in electronegativity?

Answer

A

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

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

Question 7

Q

Define electronegativity

Answer

A

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

Question 8

Q

Explaining melting point trends for groups 1 & 2

Answer

A

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

Q

Explain melting point trends for group 7

Answer

A

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

Q

State the state changes down group 7 elements

Answer

A

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

Question 11

Q

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

Answer

A

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

Q

What do u call elements along the border line

Answer

A

Metalloids
Properties of both metal & non-metals

Question 13

Q

What are redox reactions?

Answer

A

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

Question 14

Q

What are agents in terms of redox?

Answer

A

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

Question 15

Q

How to tell if it’s a redox reaction?

Answer

A

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

Q

How to identify the agent?

Answer

A

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

Question 17

Q

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

Answer

A

Disproportionation reaction

Question 18

Q

What are group 2 elements?

Answer

A

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

Q

Explain to me about group 2 elements reacting with water

Answer

A

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

Q

Explain to me about group 2 elements reacting with oxygen

Answer

A

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

Q

Explain to me about group 2 elements reacting with hydroxides

Answer

A

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

Q

Explain to me about group 2 elements reacting with sulfates

Answer

A

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

Q

What are the flame tests with group 2 elements?

Answer

A

Wait…. no need

Question 24

Q

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

Answer

A

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

Q

Explain thermal stability of carbonates down the group

Answer

A

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

Q

Explain thermal stability for group 2 elements with hydroxides?

Answer

A

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

Q

Tell me about the basic character of oxides & hydroxides

Answer

A

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

Q

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

Answer

A

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

Question 29

Q

What are group 7 elements?

Answer

A

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

Q

How are halogens diatomic molecules? (X₂)

Answer

A

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

Question 31

Q

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

Answer

A

Weak intermolecular forces (vdw’s) between molecules

Question 32

Q

Explain to me about halogens with metals

Answer

A

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

Q

Explain to me about halogens and their oxidising power

Answer

A

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

Q

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

Answer

A

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

Q

The origins for testing for aqueous halide ions

Answer

A

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

Q

Explain to me about the displacement reactions of the halogens?

Answer

A

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

Q

What are the relative oxidising strength of the halogens?

Answer

A

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

Q

Tell me about chlorine & fluorine in water treatment

Answer

A

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

Q

Tell me about fluorine in water treatment

Answer

A

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

Q

Tell me about soluble & insoluble salts

Answer

A

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

Q

Tell me about insoluble salts

Answer

A

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)

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Unit 1.6 Periodic table Flashcards

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