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Chemistry OCR A-Level > Chapter 7 > Flashcards

Chapter 7 Flashcards

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1
Q

How did Mendeleev arrange elements?

A

1) in order of atomic mass

2) lining up elements in groups with similar properties

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2
Q

If the group properties did not fit, what did Mendeleev do with the elements?

A

Swapped them around and left gaps, assuming that the atomic mass measurements were incorrect and that some elements were yet to be discovered

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3
Q

What does the arrangement, pattern, and shape of the periodic table reveal?

A

trends among the elements

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4
Q

What are the positions of the elements in the periodic table linked to?

A

their physical and chemical properties

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5
Q

Why is the periodic table essential?

A

for predicting the properties of elements and their compounds

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6
Q

How are elements arranged?

A

in order of increasing atomic number

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7
Q

What is the same in each group of the periodic table?

A

Each element in a group has atoms with the same number of outer-shell electrons and similar properties

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8
Q

What does the number of the period give?

A

the number of the highest energy electron shell in an element’s atoms

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9
Q

What is periodicity?

A

A repeating trend in properties of the elements across each element

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10
Q

What is the most obvious periodicity in properties?

A

the trend from metals to non-metals

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11
Q

What properties of periodicity are there?

A

1) electron configuration
2) ionisation energy
3) structure
4) melting points

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12
Q

What is the chemistry of each element determined by?

A

its electron configuration, particularly the outer, highest energy electron shell

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13
Q

What does each period start with?

A

an electron in a new highest energy shell

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14
Q

What do the blocks in the periodic table correspond to?

A

their highest energy sub-shell

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15
Q

What are the 4 blocks in the periodic table?

A

2) s
2) p
3) f
4) d

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16
Q

What does ionisation energy measure?

A

how easily an atom loses electrons to form positive ions

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17
Q

What is the first ionisation energy?

A

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

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18
Q

What are the factors involving ionisation energy?

A

1) atomic radius
2) nuclear charge
3) electron shielding

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19
Q

How are electrons held in their shells?

A

by attraction from the nucleus

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20
Q

Where will the first electron be lost from?

A

the highest energy level

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21
Q

How does atomic radius affect ionisation energy?

A

1) the greater the distance between the nucleus and outer electrons, the less the nuclear attraction
2) force of attraction decreases with increasing distance

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22
Q

How does nuclear charge affect ionisation energy?

A

The more protons in the nucleus, the greater the attraction between the nucleus and outer electrons

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23
Q

How does electron shielding affect ionisation energy?

A

inner-shell electrons repel outer-shell electron and reduces the attraction between nucleus and outer electrons

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24
Q

What is the shielding effect?

A

Electrons are negatively charged and so inner-shell electrons repel outer-shell electrons

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25
Q

How many ionisation energies does does an element have?

A

As many as there are electrons

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26
Q

What is the first ionisation energy of helium?

A

He(g) –> He⁺(g) + e⁻

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27
Q

What is the second ionisation energy of helium?

A

He⁺(g) –> He²⁺(g) + e⁻

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28
Q

Is the second ionisation energy of helium greater or smaller than the first ionisation energy?

A

greater

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29
Q

Why is the second ionisation energy of helium greater than the first ionisation energy?

A

1) after the 1st electron is lost, the single electron is pulled closer to the nucleus
2) nuclear attraction on the remaining electron increases
3) more ionisation energy will be needed to remove 2nd electron

30
Q

What is the second ionisation energy?

A

The energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions

31
Q

What do successive ionisation energies provide evidence for?

A

the different electron energy levels in an atom

32
Q

What do successive ionisation energies allow predictions to be made about?

A

1) number of electrons in the outer shell
2) group of the element in the periodic table
3) identity of an element

33
Q

What do periodic trends in first ionisation energies provide evidence for?

A

the existence of shells and sub-shells

34
Q

What are the patterns in the first ionisation energies for the first 20 elements in the periodic table?

A

1) a general increase in first ionisation energy across each period
2) a sharp decrease in first ionisation energy between the end of one period and the start of the next period

35
Q

Do first ionisation energies increase or decrease down a group?

A

decrease

36
Q

Why do first ionisation energies decrease down a group?

A

1) atomic radius increases
2) more inner shells so shielding increases
3) nuclear attraction on outer electrons decreases

37
Q

Do first ionisation energies increase or decrease across a period?

A

increase

38
Q

Why do first ionisation energies increase across a period?

A

1) nuclear charge increases
2) same shell: similar shielding
3) nuclear attraction increases
4) atomic radius decreases

39
Q

What does the fall in first ionisation energy from beryllium to boron mark?

A

the start of filling the 2p sub-shell

40
Q

Why is the first ionisation energy of boron less than beryllium?

A

1) 2p sub-shell in boron has a higher energy than 2s sub-shell in beryllium
2) 2p electron in boron is easier to remove than one of the 2s electrons in beryllium

41
Q

What does the fall in first ionisation energy from nitrogen to oxygen mark?

A

the start of electron pairing in the p-orbitals of the 2p sub-shell

42
Q

Why is the first ionisation energy of oxygen less than nitrogen?

A

1) highest energy electrons are in a 2p sub-shell
2) in oxygen, paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from oxygen than nitrogen

43
Q

What are the elements near to the metal/non-metal divide called?

A

semi-metals/metalloids

44
Q

At room temperature, what state are all metals except mercury?

A

solid

45
Q

What state is mercury at room temperature?

A

liquid

46
Q

What is the one constant property of all metals?

A

ability to conduct electricity

47
Q

What is metallic bonding?

A

the strong electrostatic attraction between cations (positive ions) and delocalised electrons

48
Q

In a solid metal structure, what does each atom do?

A

donate its negative outer-shell electrons to a shared pool of electrons, which are delocalised throughout the whole structure

49
Q

What is the structure of a metal?

A

giant metallic lattice

50
Q

What are the properties of metals?

A

1) stong metallic bonds
2) high electrical conductivity
3) high melting and boiling points

51
Q

How do electrons conduct electricity?

A

when a voltage is applied, delocalised electrons move through the structure, carrying charge

52
Q

Can ionic compounds conduct electricity in the solid state?

A

no

53
Q

What does the melting point of metals depend on?

A

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

54
Q

Why are high temperatures needed to melt metals?

A

to provide the large amount of energy needed to overcome the strong electrostatic attraction between the cations and electrons

55
Q

Do metals dissolve?

A

no

56
Q

What is the structure of many non-metallic elements?

A

simple molecular lattice

57
Q

How are simple molecular lattice structures held together?

A

weak intermolecular forces

58
Q

Do simple molecules have high or low melting and boiling points?

A

low

59
Q

What is the structure of the non-metals boron, carbon, and silicon?

A

giant covalent lattice

60
Q

How are giant covalent lattice structures held together?

A

strong covalent bonds

61
Q

How many electrons do atoms of carbon and silicon have in their outer shells?

A

4

62
Q

What is the shape of the structure carbon and silicon form when bonding to other carbon or silicon atoms?

A

tetrahedral

63
Q

What are the bond angles in a tetrahedral structure?

A

109.5°

64
Q

What are the properties of substances with a giant covalent lattice structure?

A

1) high melting and boiling points
2) insolubility
3) non-conductivity

65
Q

Why do giant covalent lattices have high melting and boiling points?

A

1) covalent bonds are strong

2) high temperatures necessary to provide the large quantity of energy needed to break strong covalent bonds

66
Q

Why are giant covalent lattices insoluble?

A

covalent bonds holding together atoms in the lattice are too strong to be broken by interaction with solvents

67
Q

What giant covalent lattices can conduct electricity?

A

1) graphene

2) graphite

68
Q

Why can giant covalent lattices not conduct electricity?

A

all four outer-shell electrons are involved in covalent bonding, none are available for conducting electricity

69
Q

Why can graphene and graphite conduct electricity?

A

1) only three of four outer-shell electrons are used in covalent bonding
2) remaining electron is released into pool of delocalised electrons shared by all atoms in the structure

70
Q

What is graphene?

A

a single layer of graphite

71
Q

Does melting point increase/decrease from Group 14 to Group 15?

A

increase

72
Q

Why is there a sharp decrease in melting point from Group 15 to Group 18

A

change from giant to simple molecular structures - weak forces to overcome)

Chemistry OCR A-Level (28 decks)

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