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A Level Chemistry > Module 2: Chapter 5 - Electrons and Bonding > Flashcards

Module 2: Chapter 5 - Electrons and Bonding Flashcards

1
Q

What are electrons arranged into?

A

Electron shells (energy levels), which contain subshells (sub-levels), which contain orbitals

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2
Q

How many electrons can 1 orbital hold?

A

2

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3
Q

Describe the 2 electrons in an orbital.

A

The electrons spin in opposite directions, 1 electron spins up and the other spins down

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4
Q

What is an electron orbital?

A

A region of space in which electrons are most likely to be in. It can hold 2 electrons with opposite spins

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5
Q

What is the shape of an s orbital?

A
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6
Q

What is the shape of a P orbital?

A
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7
Q

How many orbitals are there in subshell S?

A

1

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8
Q

How many orbitals are there in subshell P?

A

3

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9
Q

How many orbitals are there in subshell D?

A

5

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10
Q

How many orbitals are there in subshell F?

A

7

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11
Q

How many electrons can be in subshell S?

A

2

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12
Q

How many electrons can be in subshell P?

A

6

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13
Q

How many electrons can be in subshell D?

A

10

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14
Q

How many electrons can be in subshell F?

A

14

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15
Q

What is the order of subshells in ascending energy (order in which they are filled)?

A

1s
2s
2p
3s
3p
4s
3d
4p

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16
Q

What subshells does shell 1 contain?

A

S

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17
Q

What subshells does shell 2 contain?

A

S and P

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18
Q

What subshells does shell 3 contain?

A

S, P, and D

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19
Q

What subshells does shell 4 contain?

A

S, P, D, and F

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20
Q

What is the Aufbau Principle?

A

Electrons enter the lowest energy orbital available

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21
Q

What is Hund’s rule?

A

Electrons prefer to occupy orbitals on their own, and only pair up when no empty orbitals of the same energy are available - Electrons singly occupy each orbital in a subshell before pairing starts

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22
Q

In which order are electrons lost when an ion is formed?

A

The highest energy electrons are lost first (4s electrons are lost before 3d despite being filled after)

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23
Q

What are the 2 exceptions to the electronic structures?

A

Copper (Cu) and Chromium (Cr)

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24
Q

What is the full electronic structure of Copper?

A

1s2 2s2 2p6 3s2 3p6 4s1 3d10 (fills 3d with paired electrons before beginning to fill 4s)

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25
Q

What is the full electronic structure of chromium?

A

1s2 2s2 2p6 3s2 3p6 4s1 3d5 (fills 3d with unpaired electrons before beginning to fill 4s)

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26
Q

How does subshell energy change with distance from the nucleus

A

It increases

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27
Q

Draw the electron configuration table (highest energy subshell table)

A
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28
Q

What is the electronic configuration of S?

A

1s² 2s² 2p⁶ 3s² 3p⁴

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29
Q

What is the electronic configuration of Br?

A

1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵

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30
Q

What is the shorthand electronic configuration of Ni?

A

[Ar] 4s² 3d⁸

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31
Q

What is the shorthand electronic configuration of Ca²⁺?

A

[Ar]

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32
Q

What is the electronic configuration diagram of Fe?

A
33
Q

What is the electronic configuration diagram of Fe³⁺?

A
34
Q

How many electrons can fit in shell 1?

A

2

35
Q

How many electrons can fit in shell 2?

A

8

36
Q

How many electrons can fit in shell 3?

A

18

37
Q

How many electrons can fit in shell 4?

A

32

38
Q

What is the shell number / energy level number known as?

A

The principal quantum number, n

39
Q

How does energy change as the principal quantum number increases?

A

It increases

40
Q

Why can an orbital only hold 2 electrons with opposite spins?

A

Electrons are both negatively charged and therefore repel each other. By having opposite spins (up or down), it helps to counteract the repulsion between the negative charges of the 2 electrons

41
Q

What is ionic bonding?

A

The electrostatic attraction between positive and negative ions

42
Q

In which direction does ionic bonding act?

A

All directions

43
Q

Describe the structure of ionic compounds

A

Giant ionic lattices resulting from oppositely charges ions strongly attracted in all directions

44
Q

Why are ionic compounds solid at room temperature?

A

Ionic compounds contain many strong ionic bonds. It requires high temperatures to provide the large quantity of energy needed to overcome the strong electrostatic attraction between the ions. However, at room temperature there is insufficient energy to overcome the strong electrostatic forces of attraction and therefore ionic compounds are solid at room temperature

45
Q

How does ionic charges affect melting and boiling points of ionic compounds?

A

The melting points are higher for lattices containing ions with greater ionic charges as there is a stronger attraction between the ions

46
Q

How does size of ions affect melting and boiling points of ionic compounds?

A

The melting points are higher for lattices containing physically smaller ions as the ions can pack much more tightly within the lattice and therefore there is a stronger attraction between the ions

47
Q

What is the solubility of ionic compounds?

A

Many ionic compounds are soluble in polar solvents, such as water

48
Q

How does ionic charge affect solubility?

A

The greater the ionic charge, the less soluble the substance as the ionic attraction may be too strong for the solvent to break down the lattice structure

49
Q

What processes does solubility require?

A

Solubility requires 2 main processes:
*The ionic lattice must be broken down
*Solvent molecules must attract and surround the ions

50
Q

Do ionic substances conduct as a solid and why?

A

No as the ions are in a fixed position in the giant ionic lattice and therefore there are no mobile charge carriers. Ionic substances are non-conductors in the solid state

51
Q

Do ionic substances conduct when molten/aqueous solution and why?

A

Yes as the solid ionic lattice breaks down meaning that the ions are now free to move as mobile charge carriers. Ionic substances are conductors of electricity when molten/aqueous solution

52
Q

Explain why ionic compounds dissolve in water

A

Polar water molecules are attracted towards the ions on the surface of the ionic lattice. Water molecules bond to the ions, weakening the ionic bonding. The ionic bonds are broken and the ions become surrounded by water molecules, breaking away from the lattice

53
Q

What is covalent bonding?

A

The strong electrostatic force of attraction between a shared pair of electrons and the nuclei of the bonded atoms

53
Q

Between which type of atoms does covalent bonding occur?

A

Between non-metals

54
Q

What does “covalent bonding is localised” mean?

A

The attraction is solely between the shared pair of electrons in the covalent bond and the nuclei of the bonding atoms, it does not act in all directions (as ionic bonding does)

55
Q

Which nuclei are the bonding electrons attracted to?

A

The shared pair of electrons is attracted to the nuclei of both the bonding atoms

56
Q

Explain how covalent bonding occurs:

A

A covalent bond is the overlap of atomic orbitals, each containing one electron, to give a shared pair of electrons. The shared pair of electrons is attracted to the nuclei of both the bonding atoms.

57
Q

What is a molecule?

A

A small unit made of 2 or more atoms covalently bonded together

58
Q

What is a lone pair?

A

Paired electrons that are not shared

59
Q

When do you draw electrons on the displayed formula?

A

When there is a lone pair

60
Q

What is an example of an electron deficient compound?

A

Boron Trifluoride, BF₃

61
Q

Explain why Boron Trifluoride is electron deficient

A

Boron has the electron configuration 1s2 2s2 2p1, therefore it only has 3 valence electrons which can form covalent bonds. Therefore, once all 3 valence electrons have formed covalent bonds, boron will still only have 6 electrons causing it to be “electron deficient”

62
Q

Explain what “expansion of the octet” is

A

Elements in period 3 of the periodic table have the principal quantum number 3, meaning they contain electrons in shell 3. Therefore the atom can accommodate 18 electrons (rather than just 8) as the 3d subshell is available for expansion.

63
Q

Why is expansion of the octet only available from period 3?

A

The d subshell is only available from shell 3, which is only available from elements in the third period of the periodic table.

64
Q

How many covalent bonds can carbon form?

A

4

65
Q

How many covalent bonds can nitrogen form?

A

3

66
Q

How many covalent bonds can oxygen form?

A

2

67
Q

How many covalent bonds can hydrogen form?

A

1

68
Q

What elements often perform expansion of the octet?

A

Phosphorus, Sulfur, Chlorine

69
Q

Explain how SF₂, SF₄, and SF₆ are all possible

A

Different numbers of unpaired electrons lead to different possibilities for covalent compounds. SF₄ and SF₆ are possible as sulfur is in period 3 and the d subshell is therefore available for expansion of the octet.

70
Q

Why is the noble gas electron structure not always correct when describing covalently bonded atoms?

A

*Electron deficiency
*expansion of the octet

71
Q

Why can period 2 elements not expand the octet?

A

There is too much energy required to expand form the 2p subshell to the 3s subshell as it is the next principal quantum shell. Therefore elements in period 2 must obey the octet rule

72
Q

What is a double covalent bond?

A

The electrostatic attraction is between 2 shared pairs of electrons and the nuclei of the bonding atoms

73
Q

What is a dative/coordinate bond?

A

a dative/coordinate bond is a covalent bond in which the both electrons in the shared pair have been supplied by a single bonding atom only.

74
Q

How does a dative bond form?

A

A dative bond is initially a lone pair on one of the bonded atoms, this lone pair is then shared with another atom/ion to form a dative covalent bond

75
Q

What is a dative covalent bond represented by in the displayed formula?

A

An arrowhead going from the atom containing the lone pair, showing how it provides both electrons

76
Q

What is average bond enthalpy?

A

A measurement of covalent bond strength. The larger the value the stronger the covalent bond

77
Q

How do you write electronic configurations?

A

(principal quantum number) (shell letter) ^ (number of electrons)

Number of electrons must be SUPERSCRIPT

78
Q

What are delocalised electrons?

A

Electrons that are not attached to a single atom or covalent bond, they are shared between more than 2 atoms

A Level Chemistry (31 decks)

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