Module 3.1 Flashcards

1
Q

How is the periodic table arranged

A

In groups and periods

In order of increasing atomic number
Metals are on the left and non -metals are on the right

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2
Q

What is a group (periodic table)

A

A vertical column in the periodic table

Each element in the same group has the same number of electrons in their outer shell
Elements in the same group have similar properties

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3
Q

What are periods (periodic tables)

A

A horizontal row in the periodic table

There are general trends in the periods

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4
Q

What is periodicity

A

The trend in properties that is repeated across each period

It can be used to predict the properties of elements

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5
Q

What do all elements in a group have

A

The same type of orbitals
The same number of electrons in the outer shell

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6
Q

How do you use the periodic table to tell the final subshell of an element

A

Depending on the ‘block’ of the periodic table

S block = s orbital
D block = d orbital
P block = p orbital
F block = f orbittal

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7
Q

How does atomic radius affect nuclear attraction

A

The larger the atomic radius, the weaker the nuclear attraction

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8
Q

How does nuclear charge effect nuclear attraction

A

The greater the charge, the greater the attraction

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9
Q

What is electron shielding

A

Electrons in inner shells repel outer shell electrons

The more inner shells the larger the shielding effect

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10
Q

How does electron shielding effect nuclear attraction

A

The larger the shielding effect, the smaller the nuclear attraction is on the outside ekectrons

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11
Q

How does nuclear attraction affect ionisation energy
(Forming positive ions)

A

The smaller the nuclear attraction, the smaller ionisation energy required

Electrons in the outer shell are removed first as they experience the lowest nuclear attraction

They require the lowest amount of ionisation energy to remove

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12
Q

What is ionisation energy

A

The energy needed to form POSITIVE ions
It is a measure of how easy it is to lose electrons

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13
Q

What is ionisation

A

The process where atoms gain or lose electrons to form an ion

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14
Q

Using ionisation energies how can you tell when there is a new shell

A

Electrons in the same shell require a similar amount of energy to remove

When the electrons between two shells are removed there is a large spike in ionisation energy

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15
Q

How does ionisation energy generally change across periods

A

Generally ionisation energy increases moving towards the non metals (as nuclear attraction is greater)

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16
Q

Why is the ionisation energy in group 3 lower than group 2

A

E.g Boron

Beryllium has a full stable orbital
Boron onky has 1 electron in the p orbital, which is easier to remove

(Easier to remove 1 unstable electron to obtain a stable orbital)

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17
Q

Why is the ionisation energy in group 6 lower than group 5

A

E.g Oxygen

Nitrogen has 3 electrons in the p orbital, it is half full and therefore stable

Oxygen has 2 electrons in the Px orbital
These electrons repel each other, so it is slightly easier to remove one of the Px electrons

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18
Q

How does atomic radius change across the period

A

It decreases, as there are more protons in the nucleus, so nuclear attraction is stronger, and the shells are pulled inwards

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19
Q

How does ionisation energy change moving down groups

A

Ionisation energy decreases as:
The number of shells increases (weaker force on outer electrons)
The shielding effect is greater

20
Q

How does nuclear attraction change going down a group

A

There are more electron shells so it decreases

(Greater shielding effect)

21
Q

What is the basic giant metallic structure

A

A lattice of cations surrounded by a sea of delocalised electrons

22
Q

Why do metals have a high melting and boiling point

A

Electrons are free to move but the cations are stationary
The attraction between the cations and delocalised electrons is strong
A high temperature is needed to overcome the metallic bonds

23
Q

Why can metals conduct electricity

A

The delocalised electrons can move freely within the lattice

24
Q

Why are metals ductile and malleable

A

The sea of electrons allows the layers of cations to slide over each other

25
Q

How do melting points change moving across the period (up to group 4)

A

The melting point increases steadily

This is because the elements have a giant structure
-the nuclear charge increases (electrons in the atom increase)
- There is a stronger attraction

26
Q

How do melting points change after group 54 (going across the periods)

A

It decreases sharply between group 4 and 5 as the elements go from giant covalent structures, to simple molecular structures

From group 5+ the melting points remain low, as the elements have simple molecular structures

27
Q

What are the physical properties of group 2 elements

A

High melting + boiling points
Light metals with low density
Form colourless compounds

28
Q

How does reactivity change going down group 2

A

Reactivity increases
Going down the group nuclear attraction decreases, so the outer electrons require less ionisation energy to lose

29
Q

How do group 2 elements react with water

A

They all lose electrons and form a hydroxide

Moving down the group each metal reacts more vigorusly with water

30
Q

How do group 2 elements react with acids

A

They form salt and hydrogen gas

The reaction becomes more vigorous moving down the group

31
Q

How does the stability of group 2 hydroxides change going down the groupb

A

Solubility increases going down the group (creating more alkaline solutions)

32
Q

What group is the halogens

A

Group 7

33
Q

How does reactivity change going down the halogens

A

Reactivity decreases going down the group
The number of electron shells increases so the shielding effect is higher - nuclear attraction is weaker

Going down the group it is harder for them to attract an additional electron

34
Q

What change would be observed if chlorine displaces bromine

A

The solution would become the colour of bromine

It would turn orange / brown

35
Q

What is the test for sulphate ions (So4 2-)

A

Add dilute HCL and barium chloride to the sulphate

A white precipitate (barium sulphate) should be produced

36
Q

How is carbonate (CO3 2-) tetswd for

A

Add a dilute strong acid to the carbonate
Effervescence / colourless gas is produced

Pass the gas through limewater
The gas turns the limewater cloudy

37
Q

How are halid ions tested for

A

Add nitric acid to the halid solution (to react with other elements)

Next add silver nitrate to the solution, to form a silver halide

38
Q

Why is silver nitrate used to test for halide ions

A

As silver halides are insoluble, so the colour of the precipitate can be used to identify the halide

39
Q

What precipitates do chlorine, bromine and iodine make after reacting with silver nitrate

A

Chlorine - white (silver chloride)

Bromine - cream (silver bromide)

Iodine - yellow (silver bromide)

40
Q

If the precipitate colour when testing for halid ions is unclear, what should you do

A

Add nitric acid

Silver chloride + dilute nitric acid = colourless solution

Silver bromide + concentrated nitric acid = colourless solution

Silver Iodide + ANY nitric acid = no change

41
Q

What colour / state is fluorine naturally

A

A pale yellow gas

42
Q

What colour / state is chlorine naturally

A

A faint green gas

43
Q

What colour / state is bromine naturally

A

A brown liquid

44
Q

What colour / state is fluorine naturally

A

A dark grey solid

45
Q

How does the stability of group 2 carbonates change going down the group

A

Going down the group, the carbonates become more stable

This is because the charge density decreases (as atomic radi increases), so the carbonate ion is less distorted (its electrons are not attracted strongly)