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OCR A-level chemistry! :( > module 3 - the periodic table > Flashcards

module 3 - the periodic table Flashcards

giving Mendeleev the credit he deserves

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1
Q

How did Mendeleev arrange the elements across a period?

A

In order of atomic mass!

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2
Q

How did Mendeleev arrange the elements in groups?

A

He put the elements in groups with similar properties

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3
Q

what does the number of the period tell you?

A

the highest energy shell number, in an element’s atoms

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4
Q

What is periodicity??

A

A repeating trend in properties of elements across each period

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5
Q

How is the modern periodic table arranged?

A

In order of increasing atomic number

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6
Q

What is the chemistry of each element determined by?

A

Its electron configuration

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7
Q

What does each period start with?

A

An electron in a new highest energy shell

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8
Q

What are the 5 properties that show periodicity across the periods?

A
  • electron configuration
  • metal to non-metal
  • ionisation energies
  • structure
  • melting point
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9
Q

What is the trend in electron configuration for each period?

A

The s and p sub-shells are filled in the same way

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10
Q

What is the trend in electron configuration in each group?

A

Elements in each group have atoms with the same number of electrons in each sub-shell

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11
Q

What des the similarity in each group give to those elements in that group?

A

Similar chemistry/ chemical reactions

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12
Q

What name is given to group 1 elements?

A

Alkali metals

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13
Q

What name is given to group 2 elements?

A

Alkaline earth metals

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14
Q

Which group did Mendeleev miss out in the periodic table? why was he unaware of it?

A

Group 18 (noble gases). They’re very unreactive and no elements from that group had been discovered at the time.

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15
Q

What does ionisation energy measure?

A

How easily an atom loses electrons to form positive ions

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16
Q

What is the first ionisation energy?

A

The energy required to remove 1 electron from each atom in 1 mole of gaseous atoms of an element to form 1 mole of gaseous 1+ ions

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17
Q

Write the equation for the 1st ionisation energy of sodium.

A

Na (g) —> Na+ (g) + e -

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18
Q

what energy level is the first electron lost from? What’s its attraction to the nucleus like?

A
  • highest energy level

- experiences the least attraction from the nucleus

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19
Q

what 3 factors affect the attraction between the nucleus and the outer electrons of an atom (and the ionisation energy)?

A
  • atomic radius
  • nuclear charge
  • electron shielding
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20
Q

how does atomic radius affect nuclear attraction with the outermost electrons?

A
  • greater distance between nucleus and outer electrons = less nuclear attraction
  • atomic radius has a large effect
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21
Q

how does nuclear charge affect nuclear attraction with the outer electrons?

A

more protons in nucleus = greater attraction between nucleus and outer electrons

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22
Q

how does electron shielding affect nuclear attraction with the outer electrons?

A
  • inner shell electrons repel outer shell electrons = shielding effect
  • reduces attraction between nucleus and outer electrons
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23
Q

In helium, why is the 2nd ionisation energy greater?

A
  • after the 1st electron is lost, the last electron is pulled closer to the nucleus
  • the nuclear attraction on the remaining electron increases and more ionisation energy is needed to remove it
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24
Q

define the second ionisation energy?

A

the energy required to remove 1 electron from each atom in 1 mole of gaseous 1+ ions of an element to form 1 mole of gaseous 2+ ions

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25
Q

when looking at an equation that shows ionisation energy, how can you tell which ionisation energy it is?

A

it’s the same as the charge on the ion produced

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26
Q

what do successive ionisation energies provide evidence of?

A

of the different electron energy levels in an atom

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27
Q

on a graph showing successive ionisation energy, what goes on the x and y axis?

A

x axis = ionisation number

y axis= ionisation energy

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28
Q

what trend is usually shown on a graph showing the successive ionisation energies of an atom of an element?

A

increase in ionisation energy

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29
Q

what does the large increase between 2 successive ionisation energies show?

A
  • electron with the higher successive ionisation energy must be removed from a different shell
  • that’s closer to the nucleus and with less shielding
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30
Q

what do periodic trends in first ionisation energies provide important evidence of?

A

the existence of shells and sub-shells

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31
Q

when looking at a graph that shows the trend in 1st ionisation energy across periods, what 2 key patterns can be seen?

A
  • a general increase in 1st ionisation energy ACROSS EACH PERIOD
  • a sharp decrease in 1st ionisation energy between the end of one period and the start of the next period
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32
Q

What happens to 1st ionisation energy when you go down a group?

A

it decreases

33
Q

1- Down a group, what are the 2 main factors for the decrease in 1st ionisation energy?

2- which outweighs the nuclear charge more?

A

1- increased atomic radius
1- more inner shells, so increased shielding

2- increased radius

34
Q

what happens to the number of shells as atomic radius increases?

A

there’s more of the inner shells

35
Q

how does the increased atomic radius and shielding make the 1st ionisation energy decrease down a group?

A

because the nuclear attraction on outer electrons decreases

36
Q

what goes on the x and y axis of a graph showing the trend in 1st ionisation energy across periods?

A

x axis = atomic number

y axis = 1st ionisation energy

37
Q

what is the most important factor across a period for the general increase in first ionisation energy?

A
  • the increased nuclear charge
38
Q

List 4 reasons why the 1st ionisation energy across a period increases?

A

across a period:

  • nuclear charge increases (number of protons)
  • all the atoms have the same number of shells = they all have similar shielding
  • nuclear attraction increases
  • atomic radius decreases as outer electrons are pulled closer
39
Q

across period 2, which two places does the ionisation energy fall?

A

from Be - B and N-O

40
Q

in period 2, why is there a drop between Be and B?

A
  • from beryllium to boron, the 2p sub-shells begin to fill.
  • the 2p sub-shell in boron has a higher energy than the 2s sub-shell in beryllium
  • so, in boron, the 2p electron is easier to remove than a 2s electron in beryllium
  • the 1st ionisation energy of boron is less than that of beryllium
41
Q

in which places does the ionisation energy drop slightly in period 3 and why?

A
  • from Mg to Al
  • from P to S
  • for the same reasons as the ones in period 2
42
Q

what does the fall in 1st IE from nitrogen to oxygen mark?

A
  • the start of electron pairing in the p-orbitals of the 2p sub-shell
43
Q

why is there a fall in 1st IE from nitrogen to oxygen?

A
  • in oxygen, the paired electrons in one of the 2p orbitals repel each other
  • this makes it easier to remove an electron from an oxygen atom rather than a nitrogen atom
  • so, the 1st IE of oxygen is less than of Nitrogen
44
Q

explain why successive ionisation energies always increase?

A
  • As each electron is removed, the outer shell is drawn closer to the nucleus.
  • Nuclear attraction is greater and more energy is needed to remove the next electron
45
Q

Explain the sharp drop in 1st ionisation energy between Ne and Na ?

A
  • sharp drop because of the addition of a new shell with a resulting increase in distance and shielding
  • this decreases the nuclear attraction on the outer electrons and decreases the IE
46
Q

The first six ionisation energies of an element in period 3 are 787, 1577, 3232, 4356, 16091, 19805 kj mole^-1. identify the element and explain your reasoning.

A
  • Silicon
  • because there’s a large increase between 4th and 5th ionisation energy, so the 5th electron was removed from an inner shell
  • so there’s 4 electrons in the outer shell, so it’s in group 4
47
Q

Explain why Al has a lower 1st IE than Mg ?

A

the 3p sub-shell in Al has a higher energy level than the 3s sub-shell in Mg
- the 3p electron is easier to remove

48
Q

Explain why Sulfur has a lower first ionisation energy than Phosphorus ?

A
  • Phosphorus has 3 electrons in the 3p sub-shell with 1 electron in each 3p orbital.
  • sulfur has 4 electrons in the 3p sub-shell, with 2 electrons paired in one 3p orbital and 1 in each of the others
  • The paired electrons in sulfur repel each other, making it easier to remove 1 of them rather than an unpaired electron.
49
Q

which elements are metalloids and why are they called that?

A
  • elements near to the metal-non metal divide e.g. B, Si, Ge, As e.t.c.
  • they can show in-between properties
50
Q

at the divide between metal and non-metal, what is the trend down the groups?

A
  • non-metal to metal
51
Q

are germanium and silicon metals or non-metals?

A

semi-metals!!

52
Q

when are all metals in a solid state, and what is the exception?

A
  • at room temp

- except mercury

53
Q

what property do al metals have?

A

they can conduct electricity

54
Q

in a solid metal structure, what has each atom donated and where to?

A
  • its negative outer shell electrons to a shared pool of electrons which are delocalised throughout the whole structure
55
Q

in a metal structure, what is left behind when the atoms donate their electrons?

A

the cations (consisting of nucleus and inner electron shells of the metal atoms)

56
Q

define metallic bonding?

A

the strong electrostatic attraction between cations and delocalise electrons

57
Q

in a metallic structure, can the cations move?

A

no, they’re fixed in position, to maintain the structure of the metal

58
Q

what word describes what the delocalised electrons can do?

A
  • they’re mobile

- can move throughout the structure

59
Q

in a metal structure, what are the metal atoms held together in and what by?

A
  • a giant metallic lattice

- by metallic bonds

60
Q

what are metals’ mp and bp like?

A

HIGH!!

61
Q

in what states can metals conduct electricity?

A
  • solid and liquid states
62
Q

How can metals conduct electricity?

A
  • when voltage applied across a metal

- delocalised electrons move through structure, carrying a charge

63
Q

which metal has the highest melting point?

A

tungsten (w)

64
Q

which group of metals all have lower melting points below 200 degrees?

A

group 1 metals

65
Q

what does the meting point of metals depend on?

A

the strength of the metallic bonds holding together the atoms in the giant metallic lattice

66
Q

what is the solubility of metals like?

A

metals don’t dissolve

67
Q

what structure do non-metallic molecules form? what does this lead to?

A

a simple molecular lattice held together by weak intermolecular forces
- therefore, they have low mp and bp

68
Q

what structure does c, si and b form?

what holds it together?

A
  • giant covalent lattice

- strong covalent bonds

69
Q

what’s the name of the structure for diamond?

A

tetrahedral

70
Q

what are the mp and bp of giant covalent lattices like? why?

A
  • high

- because high temp is necessary to give the large amount of energy that’s needed to break the strong covalent bonds

71
Q

what’s the solubility of giant covalent lattices like? why?

A
  • they’re insoluble in almost all solvent

- the covalent bonds in the lattice are to strong to be broken by interaction with solvents

72
Q

can giant covalent lattices conduct electricity? what are the 2 exceptions?

A

no!

- graphene and graphite

73
Q

why can’t diamond conduct electricity but graphene can?

A
  • in diamond, all 4 outer shell electrons are involved in covalent bonding, so none are available for conducting electricity
  • in graphene, a different structure is formed where 1 of the electrons is free to conduct electricity
74
Q

in graphene, what does carbon form and what happens to the remaining free electrons?

A
  • planar hexagonal layers
  • the remaining electrons are released into pool of delocalised electrons between the layers
  • so can conduct electricity
75
Q

which 2 carbon structures have layers ?

whats the bond angles?

A
  • graphene and graphite

- 120

76
Q

what is graphene?

A

a single layer of graphite

77
Q

in graphite, what are the layers of graphene bonded by?

A
  • weak London forces
78
Q

How is carbon fibre linked to the structures of graphite and graphene?

A

the structure in carbon fibre is the hexagonal arrangement as in graphene and graphite

79
Q

what are the 8 molecules with simple molecular structure? what forces do they have between their molecules?

A
  • N2, O2, F2, Ne,p4, s8, CL2, Ar

- weak London forces!

OCR A-level chemistry! :( (12 decks)

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