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A Level Chemistry > 2.2 Foundations In Chemistry 5.1-6.4 > Flashcards

2.2 Foundations In Chemistry 5.1-6.4 Flashcards

1
Q

Ionisation

A

A process in which atoms lose or gain electrons and become ions

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2
Q

First ionization energy definition

A

The energy required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous cations

M(g) —> M^+ (g) + e^- (g)

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3
Q

Second ionization energy definition

A

Involved the removal of a second electron

M^+ (g) —> M^2+ (g) + e^- (g)

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4
Q

Evidence for energy levels

A

Plotting the successive ionisation energies of an element clearly shows the existence of different energy levels and the number of electron in each

Successive ionisation energies increase as more electrons are removed
Large jumps in the ionisation energy reveal where electrons are being removed from the next principal energy level

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5
Q

What are the 4 sub levels

A

S p d f

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6
Q

Max number of electrons for each sub level

A

S- 2
P- 6
D- 10
F- 14

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7
Q

Principal energy levels’ sub levels and max electrons

A

1- 1s -2
2- 2s 2p -8
3- 3s 3p 3d -18
4- 4s 4p 4d 4f -32

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8
Q

Order of energy sub levels

A 
1s
2s
2p
3s
3p
4s
3d 
4p
4d
4f
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9
Q

The Aufbau principle

A

the lowest energy sub levels are occupied first

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10
Q

Who developed the Aufbau principle

A

Niels Bohr

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11
Q

Why does the 4s sub level fill up before the 3d

A

It’s lower in energy

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12
Q

Hydroxide

A

OH-

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13
Q

Sulphate

A

SO4^2-

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14
Q

Carbonate

A

CO3^2-

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15
Q

Nitrate

A

NO3^-

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16
Q

Ammonium

A

NH4^+

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17
Q

Silver

A

Ag+

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18
Q

Electron orbitals

A

It is impossible to exactly locate the position of an electron within an energy sub level. By measuring the electron density around the nucleus, is possible to define regions where electrons are most likely to be found at any one time.
Each energy sublevel has one or more orbitals, each one can contain a maximum of two electrons

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19
Q

Name the sublevels, the number of orbitals and the maximum number of electrons within them

A

S – 1– 2
P – 3–6
D – 5–10
F – 7–14

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20
Q

The Pauli exclusion principle

A

States that each orbital may contain no more than two electrons

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21
Q

Spin

A

The Pauli exclusion principle introduces a property of electrons called spin, which has two states: up and down. The spins of electrons in the same orbital must be opposite

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22
Q

Spin diagram

A

Shows how the orbitals are filled. Orbitals are represented by squares, and electrons by arrows pointing up or down

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23
Q

Rules for filling electrons in spin diagrams

A

When two electrons occupy a p-sub level, they could either completely fill the same p orbital or half fill two different P orbitals
Hund’s rule states that the single electrons occupy all empty orbitals within a sub level before they start to form pairs in orbitals

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24
Q

What happens when two electrons are in the same orbital

A

There is repulsion between them due to their negative charges
The most stable configuration is with single electrons in different orbitals

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25
Q

Hund’s rule

A

States that single electrons occupy all empty orbitals within a sub level before they start to form pairs in orbitals

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26
Q

Why is the electron configuration for chromium and copper different

A

In each case the 4s orbital contains one electron
This is because the 4s and 3d sub levels lie very close together in energy, and the 3d being either half full or completely full is a lower energy arrangement

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27
Q

S orbital shape

A

Spherical

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28
Q

P orbital shape

A

Dumbbell

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29
Q

Electron relative mass

A

1/2000

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30
Q

Ionic bond definition

A

The electrostatic attraction between pos and neg ions

Holds tog cations (pos ions) and anions (neg ions) in ionic compounds

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31
Q

Ionic compounds structure and properties

A

Results in giant ionic lattice structure
High mp and bp
Strong electrostatic forces between opp charged ions
Higher for lattices w ions w larger ionic charges
Conducts elec when when melted or aqueous solution
Dissolved in polar solvents (e.g. water)

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32
Q

Covalent bonds of compound ions examples

A

Hydroxide ion
Sulphate ion
Nitrate ion
Carbonate ion

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33
Q

Covalent bonding diagram of hydroxide ion

A

OH-
Single bond
Oxygen has extra electron from unknown metal

34
Q

Covalent bonding diagram of sulphate ion

A

SO4^2-
Two oxygen atoms two oxygen ions
Double binds with oxygen atoms. Single with ions
O ions have extra electron from Unknown metal

35
Q

Covalent bonding diagram of nitrate ion

A

NO3^-
Double bond with one Oxygen
Single bind with oxygen that has extra electron from unknown metal
Single (dative) bond with oxygen (both electrons from nitrogen)

36
Q

Covalent bonding diagram of carbonate ion

A

CO3^2-
Double bond with O
Single bind with two O that both have extra electron from unknown metal

37
Q

Covalent bond definition

A

The electrostatic force if attraction between the shared pair of electrons and the 2 adjacent nuclei of the atoms

38
Q

Covalent bond- orbital overlap

A

A covalent bond is the overlap of atomic orbitals, each containing one electron to give a shared pair of electrons

39
Q

Sigma bond

A

Single covalent bond
S and S
S and P
P and P

40
Q

Dative covalent bond

A

(Coordinate bond)
Covalent bond in which the shared pair of electrons has been supplied by one of the bonding atoms only. The shared electron pair was originally a lone pair of electrons on one of the bonded atoms

41
Q

Average bond enthalpy

A

A measurement of covalent bond strength

42
Q

Electron pair repulsion theory

A

The electron pairs surrounding a central atom determine the shape of the molecule or ion
The electron pairs repel one another so that they’re arranged as far apart as possible
The arrangement of electron pairs minimises repulsion and thus holds the blinded atoms in a definite shape
Different numbers of electron pairs result in different shapes

43
Q

Shape of molecules and ions with 2 electron pairs around central atom

A

Linear

180

44
Q

Shape of molecules and ions with 3 electron pairs around central atom

A 
Trigonal planar (2D)
120
45
Q

Shape of molecules and ions with 4 electron pairs around central atom

A

Tetrahedral

109.5

46
Q

Shape of molecules and ions with 5 electron pairs around central atom

A

Trigonol bipyramidal
90
120

47
Q

Shape of molecules and ions with 6 electron pairs around central atom

A

Octahedral

90

48
Q

Distortion of basic shapes due to lone pairs- 3 electron pairs and 1 lone pair

A

Trigonometrische pyramidal

107

49
Q

Distortion of basic shapes due to lone pairs- 2 electron pairs and 2 lone pairs

A

Bent

104.5

50
Q

Distortion of basic shapes due to lone pairs- 4 electron pairs and 2 lone pairs

A

Square planar

90

51
Q

Bonded pair and lone pair repulsion

A

A lone pair of electrons is slightly closer to the central atom and occupies more space than a bonded pair
Resulting in a lone pair repelling more strongly than a bonding pair

52
Q

What’s the bond angle reduced by for each lone pair

A

2.5

53
Q

Electronegativity

A

The ability of an atom to attract a bonding or non bonding pair of electrons

54
Q

Electronegativity- what changes when the bonded atoms are different elements?

A

The nuclear charges are different
Atoms may be different sizes
The shared pair of electrons may be closer to one nucleus than the other

55
Q

Electronegativity- ionic or covalent?

A

If the electroneg difference is large, one bonded atom will have will have a much greater attraction for the shared pair than the other bonded atom
The more electroneg atom will have gained control of the electrons and the bond will now be ionic rather than covalent

56
Q

Non polar bonds

A

The bonded electron pair is shared equally between the bonded atoms
A bond will appear polar when:
- the bonded atoms are the same or
- the bonded atoms have the same or similar electronegitivity

57
Q

Polar bonds

A

The bonded electron pair is shared unequally between the bonded atoms
A bond will be polar when the bonded atoms are different and have different electroneg values, resulting in a polar covalent bond

58
Q

Dipole

A

The separation of opposite charges

59
Q

Dipole in a polar covalent bond

A

Doesn’t change

Permanent dipole

60
Q

Polar molecules

A

Asymmetrical

Dipoles don’t even each other out

61
Q

3 intermolecular forces

A

Induced dipole-dipole interactions (London forces)
Permanent dipole-dipole interactions
Hydrogen bonding

62
Q

Strongest force

A

Covalent bond

63
Q

Induced dipole-dipole interactions (London forces)

A

Weak intermolecular forces that exist between all molecules, whether polar or non polar
Act between induced dipoles in diff molecules

Only temporary

64
Q

Origin of induced dipole-dipole interactions (London forces)

A

Movement of electrons produces a changing dipole in a molecule
At any instance, an instantaneous dipole will exist, but it’s position is constantly shifting
The instantaneous dipole induces a dipole on a neighbouring molecule
The induced dipole induces further dipoles on neighbouring molecules, which then attract one another

65
Q

Permanent dipole-dipole interactions

A

Act between the permanent dipoles in diff polar molecules

66
Q

Properties of simple molecular substance

A

Low mp
Low bp
Weak intermolecular forces break when melted
Covalent bonds are strong and don’t break

67
Q

Hydrogen bonds

A

Special type of permanent d-d interaction found between molecules containing:

  • an electroneg atom with a lone pair of electrons (O, N, F)
  • a hydrogen atom attached to an electroneg atom (H-O, H-N, H-F)

H bonds acts between a lone pair of electrons on an electroneg atom in one molecule and a H atom in a diff molecule

68
Q

Strongest type of intermolecular attraction

A

Hydrogen bonds

69
Q

Anomalous properties of water

A 
The solid (ice) is less dense than the liquid (water)
Water has relatively high mp and bp
70
Q

Anomalous prop of H2O- ice less dense than water

A

H bonds hold water molecules apart in open lattice struc
The water molecules in ice are further apart than in water
Solid ice is less dense and floats

71
Q

Anomalous prop of water- high mp and bp

A

Water has London forces between molecules (like all molecules, not just water)
H bonds are extra forces over and above the London forces
An appreciable quantity of energy is needed to break the hydrogen bonds in water, so water has much higher melting and bp than would be expected from just London forces

72
Q

Without H bonds, water would have a bp of …

A

-75oC

73
Q

Strength of London forces increases as..

A

Molecular size increases
Number of protons and electrons increase
Branches decrease

74
Q

Why does the strength of London forces increase as molecular size increases

A

As atomic radius increases the outer electrons become further from the nucleus
They are attracted less strongly by the nucleus and so temporary dipoles are easier to induce

75
Q

How do the number of protons and electrons effect the strength of London forces

A

The more pros and elecs the stronger

The greater the charges of the pos nucleus and neg elecs are so stronger force of attraction

76
Q

How does branching in molecules effect London forces

A

The more branching the weaker the forces

As The less the molecules can pack together so there is less surface contact between molecules

77
Q

If the molecule is polar then there are…
“ non polar “ …
Exception…

A

Permanent d-d attractions

London forces

Linear molecules with polar bonds (not polar molecule as symmetrical)

78
Q

Strength of metallic bonds depends on

A

Charge of metal ions

Size of ions

79
Q

Charge of metal ions relative to strength

A

Greater the charge the greater the attraction between the ions and the delocalised elecs and the stronger the metallic bonds

80
Q

Size of metal ions relative to strength of metallic bonds

A

Smaller the metal ion the closer the positive nucleus is to the delocalised elecs
Meaning there is a greater attraction between the two, which creates a stronger metallic bond

A Level Chemistry (16 decks)

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