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A Level Chemistry > Chapter 7 - Periodicity > Flashcards

Chapter 7 - Periodicity Flashcards

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1
Q

Who was responsible for a more accurate periodic table?

A

Mendeleev

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2
Q

What did Mendeleev do that was different?

A

He left gaps for undiscovered elements so that elements fir in patterns

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3
Q

What are the vertical columns in the periodic table called?

A

Groups

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4
Q

What are the horizontal rows in the periodic table called?

A

A period

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5
Q

How are the elements in the periodic table ordered?

A

By proton (atomic) number

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6
Q

4 Properties which vary across a period

A

Electron configuration
Ionisation energy
Structure
Melting point

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7
Q

How can the periodic table be used to work out electron configuration?

A

Each period signals the start of a new subshell, so you can read along the period and work out how many electrons are in each sub shell

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8
Q

What is the name for the elements in group 1?

A

Alkali metals

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9
Q

What is the name for elements in group 2?

A

Alkaline earth metals

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10
Q

What is the name for elements in groups 3-12?

A

Transition elements

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11
Q

What is the first ionisation energy of an element?

A

The energy needed to remove one mole of electrons from one mole of gaseous atoms

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12
Q

Why is ionisation an endothermic process?

A

It requires energy input

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13
Q

Equation for first ionisation energy

A

X (g) -> X+ +e-

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14
Q

What does a low ionisation energy indicate?

A

It is easier to form an ion

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15
Q

3 factors affecting ionisation energy

A

Nuclear charge
Atomic radius
Shielding

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16
Q

What does a high ionisation energy indicate?

A

There is strong attraction between the electron and the nucleus so more energy is needed to overcome it and remove the electron

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17
Q

How does nuclear charge impact ionisation energy?

A

The more protons there are, the stronger the attraction between the nucleus and electron will be

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18
Q

How does atomic radius impact ionisation energy?

A

The greater the distance between electron and nucleus, the weaker the force of attraction

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19
Q

How does shielding impact ionisation energy?

A

The more electrons shielding the electron from the nucleus, the weaker the force of attraction

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20
Q

Why does ionisation energy decrease down a group?

A

There is more electron shielding and the atomic radius is greater

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21
Q

Why does ionisation energy increase across a period?

A

There is a greater nuclear charge, so a stronger force of attraction

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22
Q

Why does ionisation energy drop between groups two and three?

A

The outer shell in group 3 is in a p-orbital, rather than an s-orbital. This means it is further from the nucleus and there is additional shielding, so ionisation energy drops

23
Q

Why does ionisation energy drop between groups 5 and 6?

A

In group 5 elements, the electron is being removed from a singly occupied orbital, but in group 6, the orbital contains two electrons. The repulsion between the two electrons makes it easier to remove

24
Q

What are successive ionisation energies?

A

Removing additional electrons

25
Q

What can a graph of successive ionisation energies tell us?

A

When different shells are having electrons removed and how many electrons are in each sub shell

26
Q

How are metals held together?

A

Metallic bonding

27
Q

What is the name for the structure of most metals?

A

Giant metallic lattice structure

28
Q

What is metallic bonding?

A

Delocalised electrons are electrostatically attracted to the positively charged nuclei

29
Q

What happens when the number of delocalised electrons increases?

A

Melting point increases

30
Q

Why are metals ductile and malleable?

A

There are no bonds holding the layers together so they can slide past each other

31
Q

Why are metals good thermal and electrical conductors?

A

The delocalised electrons can carry kinetic energy or electrical charge

32
Q

Why are metals insoluble?

A

The metallic bonds are very strong

33
Q

What are giant covalent lattices?

A

Huge networks of covalently bonded atoms

34
Q

What are different forms of the same element called?

A

Allotropes

35
Q

Properties of diamond

A

Very high melting point
Very hard
Insoluble
Can’t conduct electricity

36
Q

Why does diamond have such specific properties?

A

All four of its outer carbons are bonded to another carbon

37
Q

Why is graphite used as a lubricant?

A

Weak forces between layers

38
Q

Why can graphite and graphene conduct electricity?

A

Delocalised electrons to carry charge

39
Q

Why is graphite less dense than diamond?

A

The layers are far apart

40
Q

Why does graphite have a high melting point?

A

It has strong covlent bonds

41
Q

Why is graphite insoluble?

A

The covalent bonds in the sheets are very strong

42
Q

Why is graphene so strong?

A

Delocalised electrons strengthen covalent bonds

43
Q

What is a simple molecular structure?

A

A structure containing only a few atoms

44
Q

Why do simple molecules have low melting points?

A

Weak induced dipole-dipole forces to overcome

45
Q

Why do noble gases have low melting points?

A

They are monatomic, meaning weak dipole-dipole forces

46
Q

Why does melting point change across a period?

A

The type of bond formed changes

47
Q

What will happen to the melting point of metals across a period?

A

It will increase - the number of delocalised electrons increases so the bonds are harder to overcome

48
Q

Why so carbon and silicone have high melting points?

A

They form giant covalent lattices, which require lots of energy to overcome

49
Q

Why are giant covalent lattices insoluble?

A

The covalent bonds holding the atoms in the lattice are far too strong to be broken by solvents

50
Q

Across period 2 and 3, what happens to melting points between groups 1 and 14?

A

There is an increase

51
Q

What happens to melting points of elements in groups 15-18 in periods 2 and 3?

A

They decrease

52
Q

In periods 2 and 3, why does melting point decrease between group 14 and 15?

A

The elements in group 15 onwards form simple molecular structures, not giant, which require less energy to overcome

53
Q

How does the bonding in a simple molecular lattice differ to that in a giant covalent lattice?

A

A simple molecular lattice only has London forces between atoms, whereas a giant covalent lattice has covalent bonds

54
Q

Why does the melting point of metals increase with the number of electrons?

A

The number of delocalised electrons and the charge of the ion increase, which makes the metallic bonds stronger and harder to break

A Level Chemistry (29 decks)

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