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IB chemistry > topic 3- periodicity > Flashcards

topic 3- periodicity Flashcards

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1
Q

what does a periodic table consist of?

A

groups (vertical columns) and periods (horizontal rows)

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2
Q

how is the periodic table organised?

A

by order of increasing atomic number

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3
Q

elements in the same group…

A

have the same number of valence electrons

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4
Q

elements in the same period…

A

have the same number of electron shells

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5
Q

location of metals

A

right side

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6
Q

location of non-metals

A

left side

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7
Q

location of metalloids

A

border between metals and non metals

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8
Q

what are metalloids (give egs)

A

chemical elements whose physical and chemical properties fall in between the metal and non-metal categories (boron, germanium, silicon)

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9
Q

alkali metals

A

group 1

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10
Q

halogens

A

group 17

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11
Q

noble gases

A

group 18

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12
Q

groups 3-12

A

transition metals

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13
Q

top row of ‘island’ at bottom

A

lanthanides

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14
Q

bottom row of ‘island’ at bottom

A

actinides

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15
Q

first 2 blocks

A

outer electron is in s shell

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16
Q

middle block

A

outer electron is in d shell

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17
Q

far right block

A

outer electron in p shell

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18
Q

bottom block

A

outer electron in f block

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19
Q

period number (n)=

A

outer energy level occupied by electrons

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20
Q

as you go down a group, atomic radius…

A

increases:
- more shells (increased shielding/distance)
- increase in nuclear charge outweighed

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21
Q

as you go across a period, atomic radius…

A

decreases:
- shielding stays the same; nuclear charge increases
- outer electrons experience stronger attraction to the nucleus

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22
Q

define ionic radius

A

the distance between the nucleus of an ion and the outermost shell of the ion.

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23
Q

as you go down a group, ionic radius…

A

increases:
- there is an extra complete inner shell of electrons

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24
Q

as you go across a period, ionic radius…

A

decreases:
- nuclear charge increases
- electrons feel a stronger attraction to the nucleus

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25
Q

define ionisation energy

A

The minimum energy required to remove one electron from each atom in a mole of gaseous atoms.

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26
Q

as you go down a group, ionisation energy…

A

decreases:
- shielding and distance increase
- this outweighs the increase in nuclear charge

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27
Q

as you go across a period, ionisation energy…

A

increases:
- all electrons are in the same shell
- nuclear charge increases

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28
Q

define electronegativity

A

a measure of the tendency of an atom to attract a bonding pair of electrons

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29
Q

as you go down a group, electronegativity…

A

decreases:
- shielding and distance increase
- electrons experience less attraction to the nucleus

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30
Q

as you go across a period, electronegativity…

A

increases:
- increased nuclear charge
- shielding is almost the same

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31
Q

define electron affinity

A

the enthalpy change when one mole of gaseous atoms gain an electron each, forming negative ions

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32
Q

as you go down a group, electron affinity…

A

decreases:
- nuclear charge decreases
- as a result, it becomes more difficult to add an electron - resulting in a decrease in the amount of energy released.

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33
Q

as you go across a period, electron affinity

A

increases:
- atomic radius gets smaller
- nuclear charge gets larger.

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34
Q

what are the similarities in the electron configurations of the alkali metals?

A

they all have 1 electron in their valence shell

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35
Q

what are the differences in the electron configurations of the alkali metals?

A

as you go down the group, there are more electron shells.

36
Q

describe the chemical and physical properties of alkali metals

A
  • they get more reactive as you go down the group
  • they lose an electron to become a positive ion
37
Q

Lithium- observations and relative reactivity

A
  • effervescence
  • metal disappears
  • least reactive
38
Q

Lithium- ease of ignition/flame colour

A
  • least easy
  • red
39
Q

Lithium- products formed

A

lithium hydroxide + hydrogen

40
Q

Sodium- observations and relative reactivity

A
  • effervescence
  • metal floats and moves on water
  • metal disappears
41
Q

Sodium- ease of ignition/flame colour

A
  • middle easy
  • yellow/orange
42
Q

Sodium- products formed

A

sodium hydroxide + hydrogen

43
Q

Potassium- observations and relative reactivity

A
  • effervescence
  • floats and moves quickly on water
  • lilac flame
  • most reactive
44
Q

Potassium- ease of ignition/flame colour

A
  • lilac flame
  • easiest
45
Q

Potassium- products formed

A

potassium hydroxide + hydrogen

46
Q

why are group 1 metals called alkali metals?

A

their reactions with water result in the formation of an alkaline solution

47
Q

write balanced equations for the reactions of the alkali metals with water

A

2Li (s) + 2H2O (l) -> 2LiOH (aq) + H2 (g)
2Na (s) + 2H2O (l) -> 2NaOH (aq) + H2 (g)
2K (s) + 2H2O (l) -> 2KOH (aq) + H2

48
Q

describe and explain the trend in reactivity of group 1 metals

A

more reactive as you go down the group:
- more electron shells so increased shielding
- easier for outer electron to be lost

49
Q

halogens are ——– oxidising agents

A

very good (they gain electrons easily)

50
Q

what is the oxidising power of halogens?

A

the extent to which a halogen atom is able to attract an electron to form a halide ion

51
Q

the halogens become —– reactive down the group

A

less

52
Q

oxidising power of halogens —— down the group

A

decreases

53
Q

explain why chlorine is less reactive than fluorine

A

chlorine has a greater atomic radius than fluorine (more shielding);
electrons experience less attraction to nucleus;
fluorine gains electron more easily;
less oxidising power

54
Q

why do halogens have a low melting and boiling point?

A
  • diatomic molecules
  • London forces
  • easily overcome
55
Q

How do we test for halides?

A

add an aqueous solution of silver nitrate, AgNO3 (aq) then try to dissolve in ammonia, NH3

56
Q

give the symbol equations for the reaction of silver nitrate with halides

A

AgNO3 (aq) + NaX (aq) -> NaNO3 (aq) + AgX
Ag+ (aq) + X- (aq) -> AgX (s)

57
Q

testing for chlorides result

A
  • white precipitate formed
  • fully dissolves in dilute NH3
58
Q

testing for bromides result

A
  • cream precipitate
  • partially dissolve in concentrated NH3
59
Q

testing for iodides

A
  • yellow precipitate
  • does not dissolve in concentrated NH3
60
Q

chlorine solution

A

very pale green

61
Q

bromine solution

A

orange

62
Q

iodine solution

A

brown

63
Q

why use cyclohexane when determining the reactivity of halogens?

A

the halogens are very soluble in the cyclohexane layer which will then change colour depending on the dissolved halogen

64
Q

halogens in cyclohexane

A

chlorine and bromine= same
iodine= purple

65
Q

symbol equations for chlorine displacing bromine

A

Cl2 (aq) + 2NaBr (aq) -> 2NaCl (aq) + Br2 (aq)
Cl2 (aq) + 2Br- (aq) -> 2Cl- (aq) + Br2 (aq)

66
Q

why are halogen displacement reactions redox?

A

more reactive halogen oxidises a less reactive halide ion

67
Q

why is chlorine the most reactive in terms of electrons?

A

it has the strongest affinity for electrons and will remove electrons from bromide ions and iodide ions

68
Q

give the general equation for halogens reacting with alkali metals

A

2M (s) + X2(g) -> 2MX (s)

69
Q

what is the product of halogens + alkali metals?
- colour
- solubility

A

salts
- white/colourless
- soluble in water and form colourless, neutral solutions

70
Q

define disproportionation

A

a reaction in which the same element is both reduced and oxidised.

71
Q

give 2 examples of disproportionation
- chlorine as a disinfectant in water
- chlorine in bleach

A

Cl2 (aq) + H2O (l) -> HCl (aq) + HClO (aq)
Cl2 (aq) + 2NaOH (aq) -> NaCl (aq) + NaClO (aq) + H2O (l)

72
Q

describe the trend in acid- base nature of the oxide of an element across a period

A

Acidity increases from left to right, ranging from strongly basic oxides on the left to strongly acidic ones on the right, with an amphoteric oxide (aluminum oxide) in the middle

73
Q

metals are ——, non metals are ——-

A

basic, acidic

74
Q

balanced symbol equation for the reaction of Na2O with water

A

Na2O + H2O → 2NaOH

75
Q

balanced symbol equation for the reaction of MgO with water

A

MgO(s) + H2O(l) → Mg(OH)2(aq)

76
Q

balanced symbol equation for the reaction of P4O10 with water

A

P4O10(s) + 6H2O (l) -> 4H3PO4 (s)

77
Q

NaOH pH and solubility

A

13/14 (strong base), solubility

78
Q

Mg(OH)2 pH and solubility

A

9, not very soluble

79
Q

H3PO4 pH

A

1/2 (H+ ions)

80
Q

balanced symbol equation for the reactions of nitrogen and sulphur oxides with water

A

2NO2 (g) + H2O (l) → HNO2 (aq) + HNO3 (aq)
SO2 (g) + H2O (l) → H2SO3 (l)

81
Q

use the products of the reactions of nitrogen and sulfur oxides with water to explain the pH of the resulting solution

A
  1. pH 2/3
  2. pH 1
82
Q

acid rain- formation of sulphur based acids

A

Fossil fuels are often contaminated with small amounts of sulfur impurities:
- When these contaminated fossil fuels are combusted, the sulfur in the fuels get oxidised to sulfur dioxide
- S (s) + O2 (g) → SO2 (g)
- Sulfur dioxide may be further oxidised to sulfur trioxide
- 2SO2 (g) + O2 (g) ⇌ 2SO3 (g)
The sulfur dioxide and sulfur trioxide then dissolve in rainwater droplets to form sulfurous acid and sulfuric acid
- SO2(g) + H2O (l) → H2SO3 (aq)
- SO3 (g) + H2O (l) → H2SO4 (aq)

83
Q

acid rain- formation from nitrogen oxides

A

The temperature in an internal combustion engine can reach over 2000 °C
Here, nitrogen and oxygen, which at normal temperatures don’t react, combine to form nitrogen monoxide:
N2 (g)+ O2 (g) ⇌ 2NO (g)
Nitrogen monoxide reacts further forming nitrogen dioxide:
2NO (g) + O2 (g) ⇌ 2NO2 (g)

Nitrogen dioxide gas reacts with rain water to form a mixture of nitrous and nitric acids, which contribute to acid rain:
2NO2 (g) + H2O (l) → HNO2 (aq) + HNO3 (aq)

Lightning strikes can also trigger the formation of nitrogen monoxide and nitrogen dioxides in air
Nitrogen dioxide gas reacts with rain water and more oxygen to form nitric acid
4NO2 (g) + 2H2O (l) + O2 (g)→ 4HNO3 (aq)

  • When the clouds rise, the temperature decreases, and the droplets get larger
  • When the droplet containing these acids are heavy enough, they will fall down as acid rain
84
Q

incomplete and complete combustion of methane

A

CH4 + 2O2 → CO2 + 2H2O

2CH4 + 3O2→ 2CO + 4H2O
CH4 + O2→ C + 2H2O

85
Q

catalytic converters

A
  • used in car exhaust boxes to reduce air pollution
  • the transition metal catalysts facilitate the conversion of these pollutants into less harmful products
86
Q

steps of a catalytic converter

A
  1. Molecules of carbon monoxide and nitrogen monoxide are absorbed onto the surface
  2. The bonds in both molecules are weakened causing them to react together to form carbon dioxide and nitrogen
  3. The products are then desorbed from the surface of the catalyst

2NO (g) + 2CO (g) → N2 (g) + 2CO2 (g)

CH3CH2CH3 (g) + 5O2 (g) → 3CO2 (g) + 4H2O (g)

87
Q

effects of acid rain

A

Acid deposition can react with metals and rocks (such as limestone) causing buildings and statues to get damaged

Apart from acid deposition directly falling on leaves and killing plants, acid particulates can block stomata ( plant pores) and prevent gaseous exchange

When acid rain falls on rivers and lakes the pH can fall to levels that are unable to support life

IB chemistry (29 decks)

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