C7-8. Electronic Structure and Periodicity

Question 1

What is the oxidation number of an element on its own?

Answer 1

This is always zero.

Question 2

How do we assign oxidation numbers in compounds?

Answer 2

• Each atom within the compound is assigned an oxidation number, which is usually identical to its ionic charge.
• The sum of these is always equal to the charge of the complex ion.

Question 3

What are the oxidation numbers of the following exceptional cases of atoms in compounds?
1. H in hydrides
2. O in peroxides
3. O bonded to F

Answer 3

1. This is -1, such as in NaH.
2. This is -1, such as in H2O2.
3. This is +2, such as in F2O.

Question 4

What can we say about the oxidation number of an element bonded to a more electronegative element?

Answer 4

The oxidation number of the less electronegative atom will have a positive oxidation number.

Question 5

What convention must we follow when writing oxidation numbers?

Answer 5

The sign always comes before the number itself, unlike with ionic charges where it comes afterwards.

Question 6

How do we name atoms and compounds which can have different ionic charges? Give an example.

Answer 6

• We present the oxidation number of the atom or compound as a Roman numeral given after the name of an element.
• Nitrate (III) i.e. nitrite, NO2- has a +3 oxidation number, and nitrate (V) i.e. NO3- has a +5 oxidation number.

Question 7

How do we define an oxidation reaction?

Answer 7

This is the loss of electrons, and the gain in oxidation number. Remember the OILRIG mnemonic.

Question 8

How do we define a reduction reaction?

Answer 8

This is the gain in electrons, but the loss in oxidation number. Think of it as, the oxidation number reduces.

Question 9

What is a redox reaction defined as?

Answer 9

A reaction in which species are both oxidised and reduced - i.e. if one of the processes happens, so must the other. This could be for example a metal-acid reaction.

Question 10

What is first ionisation energy defined as?

Answer 10

• The energy required to remove one electron from each atom In one mole of gaseous atoms of an element, in order to form one mole of gaseous 1+ ions.
• Think of it is a measure of how easily an atom loses electrons to form ions.

Question 11

Why must ionisation occur in the gaseous state?

Answer 11

This is because the removal of electrons is only achievable in high-energy states.

Question 12

What is the general formula for 1st ionisation energy?

Answer 12

X(g) = X+(g) + e-

Question 13

Why does first ionisation energy decrease down the group?

Answer 13

• The atomic radius increases due to the addition of new shells
• This also gives rise to more shielding
• Therefore the attraction between valence electrons and the nucleus is reduced

Question 14

Why does first ionisation energy generally increase across the period?

Answer 14

• Atomic radius decreases due to the addition of valence electrons in the same shell, and shielding is the same
• Nuclear charge increases due to the addition of protons
• Increasing attraction between valence electrons and nucleus

Question 15

How can we explain the group 2-3 dip in ionisation energies across the period?

Answer 15

• An electron is removed from a lower energy s-orbital in Group 2, whereas in Group 3 an electron is removed from a higher energy p-orbital.
• This requires less energy.

Question 16

How do we explain the group 5-6 dip in ionisation energies across the period?

Answer 16

• The Group 5 atom’s orbitals are singularly occupied, so the electron added in Group 6 must pair up, giving rise to repulsion as electrons have like charges.
• This means that less energy is required to overcome forces of attraction.

Question 17

What are successive ionisation energies defined as?

Answer 17

The energy required to successively remove electrons from each atom In one mole of gaseous atoms of an element.

Question 18

What do the ‘jumps’ on a graph of successive ionisation energies represent?

Answer 18

These represent the loss of an energy level (shell). It requires more energy to remove an electron from a shell closer to the nucleus, as there is less shielding, and nuclear attraction is greater.

Question 19

Why do successive ionisation energies increase logarithmically?

Answer 19

• As each electron is removed, the valence electrons are drawn closer to the nucleus since there are more protons than electrons.
• This gives rise to an increase in nuclear attraction, due to reduced atomic radius.

Question 20

How do we define periodicity?

Answer 20

A retreating trend of physical, chemical and atomic properties across the period, caused by regular and predictable variations in atomic structure.

Question 21

How are elements classified in to ‘blocks’ of the periodic table?

Answer 21

The ‘block’ an element occupies is defined by which sub-shell its highest-energy electron occupies: s, p or d.

Question 22

Explain the trend in atomic radius across period 3.

Answer 22

• This exhibits a decrease
• Nuclear charge increases, and electrons are added to the same shell; shielding is unchanged
• Greater attraction, electrons drawn closer to nucleus

Question 23

Explain the trend in electronegativity across period 3.

Answer 23

• This exhibits an increase
• Nuclear charge increases and atomic radius decreases
• Giving rise to stronger attraction between the nucleus and the two bonded/shared electrons

Question 24

Describe the trend in boiling point across period 3.

Answer 24

• Increases from sodium to magnesium
• Peaks at silicon
• Steep decrease to sulfur, then slow decrease to chlorine
• Lowest at argon

Question 25

Explain the increase in boiling point between sodium and magnesium.

Answer 25

• These elements undergo metallic bonding
• As we move across the period, the cation gains a higher charge; more delocalised electrons
• The atomic radius is also smaller
• This leads to greater electrostatic attraction

Question 26

Explain why the boiling point of Period 3 metals peaks at silicon.

Answer 26

• Silicon has a giant covalent structure
• This contains a very large (indefinite) number of covalent bonds which are very strong
• Therefore lots of energy is required to overcome the forces of attraction

Question 27

Why are the boiling points of elements from phosphorus to chlorine lower than those before them in the period?

Answer 27

• These elements have a simple molecular structure
• We only have to overcome much weaker London forces between molecules
• So less energy is required to overcome the intermolecular bonds

Question 28

Explain why sulfur has the highest boiling point out of itself, phosphorus and chlorine.

Answer 28

• Consider the molecular formulae of these elements: S8, Cl2, P4
• Sulfur has the most electrons in a relative large molecule
• So London forces are much stronger, requiring more energy to overcome

Question 29

Why does argon have the lowest boiling point of the Period 3 elements?

Answer 29

• Argon is monatomic
• We only have to overcome very weak London forces between atoms
• These require little energy to overcome

Question 30

Why does 1st ionisation energy of G2 metals decrease down the group?

Answer 30

• Atomic radius increases
• Shielding increases
• Electrostatic nuclear attraction decreases

Question 31

Why does melting point decrease down group 2?

Answer 31

The size of the ion increases. Therefore, the overall charge density must decrease for a given charge. Therefore electrostatic attraction decreases.

Question 32

What properties characterise group 2 metals?

Answer 32

• They have high melting points
• They are light with low density
• They can form colourless compounds

Question 33

Why are G2 metals good reducing agents?

Answer 33

They lose 2 electrons to form 2+ ions, so they are oxidised during reactions - I.e. they gain 2 electrons from other atoms.

Question 34

Why does the reactivity of G2 metals decrease down the group?

Answer 34

• Atomic radius (size of atom) increases
• Increase in shielding
• Highest energy electrons contained in energy levels further from nucleus
• Electrostatic attraction decreases

Question 35

What happens when oxides/hydroxides of G2 metals dissolve in water? How is this affected as w ego down the group?

Answer 35

They dissociate to release OH- ions, making the resulting solution more alkaline, with a pH of 10-12. As we go down the group, solubility increases, so more OH- ions are released and solutions formed are more alkaline.

Question 36

Name a use of a G2 hydroxide (compound).

Answer 36

• Farmers use calcium hydroxide to neutralise acidic soil.
• Magnesium oxide is used in antacids to prevent indigestion.

Question 37

What happens to the decomposition of G2 carbonates as we go down the group?

Answer 37

The carbonates require more energy to decompose thermally as we go down the group.

Question 38

What are the general appearances of the main halogens?

Answer 38

• Fluorine: yellow gas
• Chlorine: pale green gas
• Bromine: brown liquid, forms orange vapour
• Iodine: grey crystalline solid, forms purple vapour

Question 39

Why does the electronegativity of halogens decrease down the group?

Answer 39

The bonded pair of electrons is held further from the nucleus as more energy levels are added. This means that the nuclear attraction between these shared electrons and the nucleus must decrease.

Question 40

Why does the boiling point of G2 metals increase down the group?

Answer 40

The increase in number of energy levels means that more electrons are contained within the atom. These therefore give rise to greater fluctuation, and therefore stronger instantaneous dipole-dipole interactions (London forces). These require more thermal energy to overcome.

Question 41

Why does the first ionisation energy of halogens decrease down the group?

Answer 41

As energy levels are added, nuclear attraction falls and shielding increases, giving rise to greater electron-electron repulsion and therefore the highest nervy electron is more easily lost.

Question 42

What do we mean when we say that halogens are good oxidising agents?

Answer 42

This means that halogens can easily be reduced (gaining electrons), I.e. they are able to cause other elements to be oxidised and lose electrons.

Question 43

How is the oxidation ability of the halogens affected as we go down the group?

Answer 43

Down the group, oxidation ability is decreased, since a larger atoms with more shells has a weakened nuclear electrostatic attraction, so the ability to gain oppositely charged electrons falls.

Question 44

What is a disproportionation reaction?

Answer 44

This is a type of redox reaction in which an element both gains and loses electrons (is both oxidised and reduced).

Question 45

What are the equations for the reactions between…
- Halogens and water
- Halogens and sodium hydroxide

Answer 45

• Cl2 + H2O = HClO + HCl
• Cl2 + 2NaOH = NaClO + H2O

Question 46

Describe the test for carbonate ions.

Answer 46

• Add a dilute acid (such as nitric acid, hydrochloric acid) to your sample on a compound in a test tube.
• Observe for effervescence (carbon dioxide gas).
• Bubble this gas through limewater which should turn from colourless - cloudy as a white ppt. is formed.

Question 47

Describe the test for sulphate ions.

Answer 47

• Add a few drops of barium chloride solution to the compound in a test tube.
• Positive test result is indicated by the formation of a white ppt.

Question 48

Describe how we would test for halide ions.

Answer 48

• Add acidified silver nitrate (silver nitrate + nitric acid) to a test tube containing the compound to be tested.
• Observe for the formation of the following precipitates: white ppt. - chloride, cream ppt. - bromide, yellow ppt. - iodide.
• Further tests are conducted using dilute and concentrated ammonia. Chloride compounds redissolve in dilute ammonia, and bromide compounds do so only in concentrated ammonia. Iodides never redissolve.

Question 49

Describe how you would test for ammonium ions.

Answer 49

• Add sodium hydroxide solution to a test tube contains the compound to be tested.
• Swirl the test tube carefully.
• Using tongs heat the solution gently over a Bunsen burner.
• A pungent gas (ammonia) should begin to evolve, and damp red litmus paper should turn blue.

Question 50

What colour solution do chloride ions form? Answer in terms of in water and cyclohexane environments.

Answer 50

Pale-green in both water and cyclohexane.

Question 51

What colour solution do bromide ions form? Answer in terms of in water and cyclohexane environments.

Answer 51

Orange in both water and cyclohexane.

Question 52

What colour solution do iodide ions form? Answer in terms of in water and cyclohexane environments.

Answer 52

Brown in water, and violet in cyclohexane,

Question 53

What is the main observation for the reaction between chlorine and water?
- Cl2 + H2O = HClO + HCl

Answer 53

• This is a disproportionation reaction which results in the formation of sodium chlorate, a bleaching agent.
• The formation of hydrochloric acid will turn red litmus paper blue.
• However, the simultaneous formation of sodium chlorate will then bleach the litmus paper, causing a colour change from blue - white.
• The water here represents the water added to red litmus paper to make it damp. So this equation essentially represents the litmus test for chlorine.