inorganic chemistry 1

Question 1

define ‘periodicity’ (periodicity)

Answer 1

• the regular recurrence of the properties of elements when they are arranged in atomic number order

Question 2

name the metals in period 3 & what groups they are in (8) (periodicity)

Answer 2

• sodium (metal) = group 1
• magnesium (metal) = group 2
• aluminium (metal) = group 3
• silicon = group 4
• phosphorous (non-metal) = group 5
• sulfur (non-metal) = group 6
• chlorine (non-metal) = group 7
• argon = group 8

Question 3

what type of structure is found on the left side of the periodic table & what type of melting/boiling point does it tend to have? (2) (periodicity)

Answer 3

• giant structures
• tend to have high MP/BP

Question 4

what type of structure is found on the right side of the periodic table & what type of melting/boiling point does it tend to have? (2) (periodicity)

Answer 4

• molecular/atomic structures
• tend to have low MP/BP

Question 5

why do the melting & boiling points of metals increase across period 3? (Na-Al) (periodicity)

Answer 5

• the strength of the metallic bond increases
• from left to right the charge on the ion increases, so more electrons join the delocalised electrons that hold the giant metallic lattice together
• more energy is needed to break the stronger metallic bonds so the MP & BP are higher

Question 6

why does silicon have a high MP/BP? (3) (periodicity)

Answer 6

• silicon has a macromolecular structure (similar to diamond)
• each silicon atom is bonded to 4 other silicon atoms by strong covalent bonds
• these bonds must be broken for it to melt, which requires a lot of energy to overcome & therefore resulting in a high MP & BP

Question 7

why do the MP/BP of the non-metals decrease across period 3? (4) (periodicity)

Answer 7

• non-metals have molecular structures
• the MP & BP of phosphorous, sulfur & chlorine are lower than of silicon because they have a simple molecular structure with weak van der Waals forces
• this depends on the number of electrons in the molecule & how closely the molecules can pack together
• breaking these forces requires less energy than breaking covalent bonds, so MP & BP are lower

Question 8

why is atomic radius difficult to define? (periodicity)

Answer 8

• due to the uncertainty over the size of the electron cloud

Question 9

when can metals form covalent bonds? (periodicity)

Answer 9

• in the gas phase

Question 10

why don’t noble gases have covalent radii? (periodicity)

Answer 10

• the don’t bond covalently with one another

Question 11

why is atomic radius a periodic property? (2) (periodicity)

Answer 11

• it decreases across each period
• there is a jump when starting the next period

Question 12

why does atomic radii decrease across a period? (4) (periodicity)

Answer 12

• proton number increases across a period, but shielding remains constant
• this causes an increase in effective nuclear charge
• this leads to a greater attraction between the nucleus & outermost electrons
• this increased charge pulls electrons closer the the nucleus, resulting in a smaller radii

Question 13

why do atomic radii increase down a group? (2) (periodicity)

Answer 13

• going down a group atoms have 1 extra complete main level of electrons compared with the oe before
• therefore the outer main electron level is further from the nucleus, so atomic radii increases

Question 14

define ‘first ionisation energy’ (periodicity)

Answer 14

• the energy required to remove 1 electron from a gaseous atom

Question 15

what is the general trend for first ionisation energy across a period? (periodicity)

Answer 15

• generally increases across a period

Question 16

what molecules have the highest & lowest first ionisation energy values? (2) (periodicity)

Answer 16

• alkali metals (e.g. Na & Li) = lowest
• noble gases = highest

Question 17

what is the trend for first ionisation energy down a group? (periodicity)

Answer 17

• decreases down any group

Question 18

why does first ionisation energy increase across a period? (2) (periodicity)

Answer 18

• the number of proteins int he nucleus increases but the electrons enter the same main level
• the increased charge in the nucleus means that it gets increasingly difficult to remove an electron

Question 19

why does the first ionisation energy decrease down a group? (4) (periodicity)

Answer 19

• number of filled inner levels increases down the group
• results in an increase in shielding
• also, the electrons enter to be removed is at an increasing distance from the nucleus & is therefore held less strongly
• therefore the outer electrons get easier to remove as they are further away from the nucleus

Question 20

why is there a drop in ionisation energy from one period to the next? (2) (periodicity)

Answer 20

• sodium is at a new main level & so there is an increase in atomic radius
• the outer electron is further from the nucleus, so is less strongly attracted & therefore easier to remove)

Question 21

briefly outline the variations of first ionisation energy across period 3 (2) (periodicity)

Answer 21

• drops between 2 & 3 (so Al has a lower 1st IE than Mg)
• drops again between groups 5 & 6) (P & S)

Question 22

explain the drop in first ionisation energy between Mg & Al (3) (periodicity)

Answer 22

• the electron removed when Al is ionised comes from a 3p orbital
• this is of higher energy than the 3s orbital removed when Mg is ionised
• the removal of electrons from the higher energy 3p orbital requires less energy & therefore is easier to remove

Question 23

explain the croup is first ionisation energy between P & S (3) (periodicity)

Answer 23

• highest energy orbital for both is 3p
• S has paired electrons whereas P only has singularly occupied orbitals
• minimal repulsion between paired electrons means that less energy is required to remove one of them

Question 24

what do the number of electrons that are easy to remove tell us about an element? (2) (periodicity)

Answer 24

• the group number
• e.g. an element in group 2 would have a large jump in value after the second electron is removed

Question 25

define ‘second ionisation energy’ (group 2)

Answer 25

- energy required to remove 1 mole of electrons from 1 mole of +1 gaseous ions to form 1 mole of +2 gaseous ions

Question 26

outline the electron arrangement for group 2 metals (Mg-Ba) (3) (group 2)

Answer 26

- all have 2 electrons in their outer s orbital - outer s orbital becomes further from the nucleus going down the group - Be = 2s2, MG= 3s2, Ca = 4s2, Sr = 5s2, Ba = 6s2

Question 27

outline the size of Mg to Ba in group 2 (3) (group 2)

Answer 27

- increases us number of shells increases - this results in more shielding - outer electrons will be further from the positive nucleus & therefore there is weaker nuclear pull

Question 28

outline the melting points of metals Mg-Ba in group 2 (3) (group 2)

Answer 28

- decreases - atomic radius increases so sea of delocalised electrons is further from the positive nucleus - therefore there strength of the metallic bond decreases (less attractive force)

Question 29

outline the ionisation energies of metals Mg-Ba in group 2 (6) (group 2)

Answer 29

- decreases - form 2+ ions - X(g) —> X2+(g) + e- - nuclear charge/shielding & distance between nucleus & electrons increase - the further an election is from the nucleus the easier it is to remove - also additional shielding going down the group

Question 30

what happens to the reactivity of metals with water as you go down group 2? (group 2)

Answer 30

- metals get more reactive with water going down the group

Question 31

state the standard equation for the reaction of a group 2 metal with water, including oxidation states (2) (group 2)

Answer 31

- X (s) + 2H2O (l) —> X(OH)2 (aq) + H2 (g) - respective oxidation states = 0, +1, +2, 0

Question 32

what is magnesium hydroxide usually used for? (2) (group 2)

Answer 32

- treatment for indigestion/milk of magnesia - neutralises stomach acid

Question 33

state & explain the reaction of magnesium with steam (2) (group 2)

Answer 33

- Mg reacts very slowly with cold water but rapidly with steam to form an alkaline oxide & hydrogen - Mg (s) + H2O (g) —> MgO (s) + H2 (g)

Question 34

what is calcium hydroxide usually used for (2) (group 2)

Answer 34

- slaked lime - used to treat acidic soil

Question 35

are group 2 metals oxidised or reduced in all their reactions, & what happens to their oxidation states? (2) (group 2)

Answer 35

- oxidised - go from oxidation states 0 to oxidation +2

Question 36

what happens to the solubility of group 2 hydroxides going down the group? (4) (group 2)

Answer 36

- hydroxides become more soluble - Mg(OH)2 is almost insoluble - Ba(OH)2 dissolves to produce a strongly alkaline solution - Ba(OH)2 (s) + aq —> Ba 2+ (aq) + 2OH- (aq)

Question 37

what happens to the solubility of surfaces as you go down the group? (group 2)

Answer 37

- become less soluble, so barium sulfate is basically insoluble

Question 39

give 2 reasons why barium useful when it is insoluble (7) (group 2)

Answer 39

- can be taken by mouth as a barium meal to outline the gut in medical x-rays - test is safe, despite barium compounds being highly toxic (because barium sulfate is so insoluble) - also used in a simple test for sulfate ions in solution - solution is first acidified with nitric/hydrochloric acid - then barium chloride solution is added to the solution under test & is a sulfate is present, a white precipitate of barium sulfate is formed - Ba 2+ (aq) + SO4 2- (aq) —> BaSO4 (s) - addition of acid removed carbonate ions as CO2