topic 3- periodicity Flashcards
(87 cards)
1
Q
what does a periodic table consist of?
A
groups (vertical columns) and periods (horizontal rows)
2
Q
how is the periodic table organised?
A
by order of increasing atomic number
3
Q
elements in the same group…
A
have the same number of valence electrons
4
Q
elements in the same period…
A
have the same number of electron shells
5
Q
location of metals
A
right side
6
Q
location of non-metals
A
left side
7
Q
location of metalloids
A
border between metals and non metals
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)
9
Q
alkali metals
A
group 1
10
Q
halogens
A
group 17
11
Q
noble gases
A
group 18
12
Q
groups 3-12
A
transition metals
13
Q
top row of ‘island’ at bottom
A
lanthanides
14
Q
bottom row of ‘island’ at bottom
A
actinides
15
Q
first 2 blocks
A
outer electron is in s shell
16
Q
middle block
A
outer electron is in d shell
17
Q
far right block
A
outer electron in p shell
18
Q
bottom block
A
outer electron in f block
19
Q
period number (n)=
A
outer energy level occupied by electrons
20
Q
as you go down a group, atomic radius…
A
increases:
- more shells (increased shielding/distance)
- increase in nuclear charge outweighed
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
22
Q
define ionic radius
A
the distance between the nucleus of an ion and the outermost shell of the ion.
23
Q
as you go down a group, ionic radius…
A
increases:
- there is an extra complete inner shell of electrons
24
Q
as you go across a period, ionic radius…
A
decreases:
- nuclear charge increases
- electrons feel a stronger attraction to the nucleus
25
define ionisation energy
The minimum energy required to remove one electron from each atom in a mole of gaseous atoms.
26
as you go down a group, ionisation energy...
decreases:
- shielding and distance increase
- this outweighs the increase in nuclear charge
27
as you go across a period, ionisation energy...
increases:
- all electrons are in the same shell
- nuclear charge increases
28
define electronegativity
a measure of the tendency of an atom to attract a bonding pair of electrons
29
as you go down a group, electronegativity...
decreases:
- shielding and distance increase
- electrons experience less attraction to the nucleus
30
as you go across a period, electronegativity...
increases:
- increased nuclear charge
- shielding is almost the same
31
define electron affinity
the enthalpy change when one mole of gaseous atoms gain an electron each, forming negative ions
32
as you go down a group, electron affinity...
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.
33
as you go across a period, electron affinity
increases:
- atomic radius gets smaller
- nuclear charge gets larger.
34
what are the similarities in the electron configurations of the alkali metals?
they all have 1 electron in their valence shell
35
what are the differences in the electron configurations of the alkali metals?
as you go down the group, there are more electron shells.
36
describe the chemical and physical properties of alkali metals
- they get more reactive as you go down the group
- they lose an electron to become a positive ion
37
Lithium- observations and relative reactivity
- effervescence
- metal disappears
- least reactive
38
Lithium- ease of ignition/flame colour
- least easy
- red
39
Lithium- products formed
lithium hydroxide + hydrogen
40
Sodium- observations and relative reactivity
- effervescence
- metal floats and moves on water
- metal disappears
41
Sodium- ease of ignition/flame colour
- middle easy
- yellow/orange
42
Sodium- products formed
sodium hydroxide + hydrogen
43
Potassium- observations and relative reactivity
- effervescence
- floats and moves quickly on water
- lilac flame
- most reactive
44
Potassium- ease of ignition/flame colour
- lilac flame
- easiest
45
Potassium- products formed
potassium hydroxide + hydrogen
46
why are group 1 metals called alkali metals?
their reactions with water result in the formation of an alkaline solution
47
write balanced equations for the reactions of the alkali metals with water
2Li (s) + 2H2O (l) -> 2LiOH (aq) + H2 (g)
2Na (s) + 2H2O (l) -> 2NaOH (aq) + H2 (g)
2K (s) + 2H2O (l) -> 2KOH (aq) + H2
48
describe and explain the trend in reactivity of group 1 metals
more reactive as you go down the group:
- more electron shells so increased shielding
- easier for outer electron to be lost
49
halogens are -------- oxidising agents
very good (they gain electrons easily)
50
what is the oxidising power of halogens?
the extent to which a halogen atom is able to attract an electron to form a halide ion
51
the halogens become ----- reactive down the group
less
52
oxidising power of halogens ------ down the group
decreases
53
explain why chlorine is less reactive than fluorine
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
why do halogens have a low melting and boiling point?
- diatomic molecules
- London forces
- easily overcome
55
How do we test for halides?
add an aqueous solution of silver nitrate, AgNO3 (aq) then try to dissolve in ammonia, NH3
56
give the symbol equations for the reaction of silver nitrate with halides
AgNO3 (aq) + NaX (aq) -> NaNO3 (aq) + AgX
Ag+ (aq) + X- (aq) -> AgX (s)
57
testing for chlorides result
- white precipitate formed
- fully dissolves in dilute NH3
58
testing for bromides result
- cream precipitate
- partially dissolve in concentrated NH3
59
testing for iodides
- yellow precipitate
- does not dissolve in concentrated NH3
60
chlorine solution
very pale green
61
bromine solution
orange
62
iodine solution
brown
63
why use cyclohexane when determining the reactivity of halogens?
the halogens are very soluble in the cyclohexane layer which will then change colour depending on the dissolved halogen
64
halogens in cyclohexane
chlorine and bromine= same
iodine= purple
65
symbol equations for chlorine displacing bromine
Cl2 (aq) + 2NaBr (aq) -> 2NaCl (aq) + Br2 (aq)
Cl2 (aq) + 2Br- (aq) -> 2Cl- (aq) + Br2 (aq)
66
why are halogen displacement reactions redox?
more reactive halogen oxidises a less reactive halide ion
67
why is chlorine the most reactive in terms of electrons?
it has the strongest affinity for electrons and will remove electrons from bromide ions and iodide ions
68
give the general equation for halogens reacting with alkali metals
2M (s) + X2(g) -> 2MX (s)
69
what is the product of halogens + alkali metals?
- colour
- solubility
salts
- white/colourless
- soluble in water and form colourless, neutral solutions
70
define disproportionation
a reaction in which the same element is both reduced and oxidised.
71
give 2 examples of disproportionation
- chlorine as a disinfectant in water
- chlorine in bleach
Cl2 (aq) + H2O (l) -> HCl (aq) + HClO (aq)
Cl2 (aq) + 2NaOH (aq) -> NaCl (aq) + NaClO (aq) + H2O (l)
72
describe the trend in acid- base nature of the oxide of an element across a period
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
metals are ------, non metals are -------
basic, acidic
74
balanced symbol equation for the reaction of Na2O with water
Na2O + H2O → 2NaOH
75
balanced symbol equation for the reaction of MgO with water
MgO(s) + H2O(l) → Mg(OH)2(aq)
76
balanced symbol equation for the reaction of P4O10 with water
P4O10(s) + 6H2O (l) -> 4H3PO4 (s)
77
NaOH pH and solubility
13/14 (strong base), solubility
78
Mg(OH)2 pH and solubility
9, not very soluble
79
H3PO4 pH
1/2 (H+ ions)
80
balanced symbol equation for the reactions of nitrogen and sulphur oxides with water
2NO2 (g) + H2O (l) → HNO2 (aq) + HNO3 (aq)
SO2 (g) + H2O (l) → H2SO3 (l)
81
use the products of the reactions of nitrogen and sulfur oxides with water to explain the pH of the resulting solution
1. pH 2/3
2. pH 1
82
acid rain- formation of sulphur based acids
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
acid rain- formation from nitrogen oxides
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
incomplete and complete combustion of methane
CH4 + 2O2 → CO2 + 2H2O
2CH4 + 3O2→ 2CO + 4H2O
CH4 + O2→ C + 2H2O
85
catalytic converters
- used in car exhaust boxes to reduce air pollution
- the transition metal catalysts facilitate the conversion of these pollutants into less harmful products
86
steps of a catalytic converter
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
effects of acid rain
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
