Brainscape's Knowledge GenomeTM


Top Flashcards (Certified)
All Flashcards

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

Second ionization energy definition

A

Involved the removal of a second electron

M^+ (g) —> M^2+ (g) + e^- (g)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

What are the 4 sub levels

A

S p d f

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Max number of electrons for each sub level

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Order of energy sub levels

A 
1s
2s
2p
3s
3p
4s
3d 
4p
4d
4f
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

The Aufbau principle

A

the lowest energy sub levels are occupied first

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Who developed the Aufbau principle

A

Niels Bohr

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

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

A

It’s lower in energy

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Hydroxide

A

OH-

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Sulphate

A

SO4^2-

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Carbonate

A

CO3^2-

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Nitrate

A

NO3^-

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

Ammonium

A

NH4^+

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

Silver

A

Ag+

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

The Pauli exclusion principle

A

States that each orbital may contain no more than two electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

Spin diagram

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Hund’s rule
States that single electrons occupy all empty orbitals within a sub level before they start to form pairs in orbitals
26
Why is the electron configuration for chromium and copper different
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
27
S orbital shape
Spherical
28
P orbital shape
Dumbbell
29
Electron relative mass
1/2000
30
Ionic bond definition
The electrostatic attraction between pos and neg ions | Holds tog cations (pos ions) and anions (neg ions) in ionic compounds
31
Ionic compounds structure and properties
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)
32
Covalent bonds of compound ions examples
Hydroxide ion Sulphate ion Nitrate ion Carbonate ion
33
Covalent bonding diagram of hydroxide ion
OH- Single bond Oxygen has extra electron from unknown metal
34
Covalent bonding diagram of sulphate ion
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
Covalent bonding diagram of nitrate ion
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
Covalent bonding diagram of carbonate ion
CO3^2- Double bond with O Single bind with two O that both have extra electron from unknown metal
37
Covalent bond definition
The electrostatic force if attraction between the shared pair of electrons and the 2 adjacent nuclei of the atoms
38
Covalent bond- orbital overlap
A covalent bond is the overlap of atomic orbitals, each containing one electron to give a shared pair of electrons
39
Sigma bond
Single covalent bond S and S S and P P and P
40
Dative covalent bond
(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
Average bond enthalpy
A measurement of covalent bond strength
42
Electron pair repulsion theory
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
Shape of molecules and ions with 2 electron pairs around central atom
Linear | 180
44
Shape of molecules and ions with 3 electron pairs around central atom
``` Trigonal planar (2D) 120 ```
45
Shape of molecules and ions with 4 electron pairs around central atom
Tetrahedral | 109.5
46
Shape of molecules and ions with 5 electron pairs around central atom
Trigonol bipyramidal 90 120
47
Shape of molecules and ions with 6 electron pairs around central atom
Octahedral | 90
48
Distortion of basic shapes due to lone pairs- 3 electron pairs and 1 lone pair
Trigonometrische pyramidal | 107
49
Distortion of basic shapes due to lone pairs- 2 electron pairs and 2 lone pairs
Bent | 104.5
50
Distortion of basic shapes due to lone pairs- 4 electron pairs and 2 lone pairs
Square planar | 90
51
Bonded pair and lone pair repulsion
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
What’s the bond angle reduced by for each lone pair
2.5
53
Electronegativity
The ability of an atom to attract a bonding or non bonding pair of electrons
54
Electronegativity- what changes when the bonded atoms are different elements?
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
Electronegativity- ionic or covalent?
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
Non polar bonds
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
Polar bonds
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
Dipole
The separation of opposite charges
59
Dipole in a polar covalent bond
Doesn’t change | Permanent dipole
60
Polar molecules
Asymmetrical | Dipoles don’t even each other out
61
3 intermolecular forces
Induced dipole-dipole interactions (London forces) Permanent dipole-dipole interactions Hydrogen bonding
62
Strongest force
Covalent bond
63
Induced dipole-dipole interactions (London forces)
Weak intermolecular forces that exist between all molecules, whether polar or non polar Act between induced dipoles in diff molecules Only temporary
64
Origin of induced dipole-dipole interactions (London forces)
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
Permanent dipole-dipole interactions
Act between the permanent dipoles in diff polar molecules
66
Properties of simple molecular substance
Low mp Low bp Weak intermolecular forces break when melted Covalent bonds are strong and don’t break
67
Hydrogen bonds
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
Strongest type of intermolecular attraction
Hydrogen bonds
69
Anomalous properties of water
``` The solid (ice) is less dense than the liquid (water) Water has relatively high mp and bp ```
70
Anomalous prop of H2O- ice less dense than water
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
Anomalous prop of water- high mp and bp
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
Without H bonds, water would have a bp of ...
-75oC
73
Strength of London forces increases as..
Molecular size increases Number of protons and electrons increase Branches decrease
74
Why does the strength of London forces increase as molecular size increases
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
How do the number of protons and electrons effect the strength of London forces
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
How does branching in molecules effect London forces
The more branching the weaker the forces | As The less the molecules can pack together so there is less surface contact between molecules
77
If the molecule is polar then there are... “ non polar “ ... Exception...
Permanent d-d attractions London forces Linear molecules with polar bonds (not polar molecule as symmetrical)
78
Strength of metallic bonds depends on
Charge of metal ions | Size of ions
79
Charge of metal ions relative to strength
Greater the charge the greater the attraction between the ions and the delocalised elecs and the stronger the metallic bonds
80
Size of metal ions relative to strength of metallic bonds
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)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2025 Bold Learning Solutions. Terms and Conditions