Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

Chemistry AS OCR > Continue 2.1 > Flashcards

Continue 2.1 Flashcards

You may prefer our related Brainscape-certified flashcards:
Periodic Table flashcards Chemistry 101 flashcards AP® Chemistry flashcards Organic Chemistry 101 flashcards
1
Q

Give some examples of some elements with oxidation numbers of -1.

A

Combined oxygen in peroxide a, for example H2O2, combined hydrogen in metal halides for example LiH and combined fluorine for example NaF, CaF2, AlF3

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

What must the sum of oxidation numbers in a compound match?

A

The overall charge

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

Does an element always have the same oxidation number?

A

Some elements can have more than one stable oxidation number, transition elements form ions with different oxidation numbers.

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

What are oxyanions?

A

Oxyanions are negative ions that contain an element along with oxygen. SO42-, NO3-, CO32-.
A element may form oxyanions in which the element has different oxidation numbers.

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

What is the oxidation numbers of NO2- nitrate (III) and NO3- nitrate (V)?

A

Oxidation numbers +3 and +5

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

What is a redox reaction?

A

A reaction in which both reduction and oxidation take place is called a redox reaction.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

What is oxidation and reduction?

A

Oxidation is the loss of electrons, reduction is the gain of electrons, reduction must always be accompanied by oxygen

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

What tends to be oxidised and what tends to be reduced?

A

Metals tend to be oxidised, loosing electrons to form positive ions and non- metals tend to be reduced, gaining electrons to form negative ions

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Describe redox reactions of metals with acids.

A

The metal is oxidised, forming positive metal ions. The hydrogen in the acid is reduced, forming the element hydrogen as a gas.

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

What is the word and ionic equation for the redox reactions of a metal and an acid?

A

Metal + acid -> salt + hydrogen

Mg(s) + 2H+ (aq) -> Mg2+(aq) + H2

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

What is the oxidation number of combined oxygen?

A

-2, for example H2O and CaO

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

How do you use oxidation numbers with equations?

A

Assign oxidation numbers to each atom in any equation to identify whether a redox reaction has taken place and what has been oxidised and what has reduced.

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

When talking about shells and orbitals what does 3n stand for?

A

The principle quantum number, n, indicates the shell that the electrons occupy, the larger the value of n the further the shell is from the nucleus, fir example the 3rd shell is 3n and can hold 18 electrons.

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

Describe atomic orbitals.

A

Each atomic orbital is made up of a number of atomic orbitals.
Each atomic orbital can hold a maximum of two electrons with opposite spins.
There are four different types of orbital- S, p, d and F

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

Describe S-orbitals,

A

An S-orbitals has a spherical shape,

From n=1 upwards, each she’ll contain done S-orbital.

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

Describe P-orbitals

A

A p-orbital has a three dimensional dumb-bell shape.
From n=z upwards each shell contains three p-orbitals, Px, Py and Pz at right angles to each other.
Each P orbital will hold 2 electrons, 3 x 2 = 6. Px, Py and Pz can therefore hold 6 electrons.

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

Describe d-orbitals and-orbitals

A

From n = 3 upwards each shell contains 5 d-orbitals

From n = 4 upwards each shell contains 7 f-orbitals

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

Describe box diagrams to represent electrons in orbitals.

A

As orbitals often has different types and shapes, chemists use box diagrams to represent electrons in orbitals.
Each box represents an individual orbital and can hold 2 electrons.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

Describe the electrons property called spin, and how this is represented on box diagrams.

A

Each electron has a negative charge but does not spin, this is because they have a property called spin. The two electrons must have opposite spins, this is represented using arrows either up it down.

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

What is an electron?

A

In some ways an electro behaves like a particle and in other ways it behaves like a wave.
We can’t even be certain where an electron is within an orbital.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

You can work out the electron configuration of an atom by following a set of rules called?

A

The Aufbau principle

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

What is the set of rules called the Aufbau principle?

A

Electrons are added, one at a time to build up an atom.
The lowest available energy level is filled first, you can consider this level as being the closest to the nucleus.
Each energy level must be full before the next, higher energy level starts to fill.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

What do sub- shells consist of?

A

Sub-shells are made of several orbitals, each with the same energy level.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

How are sub-shells filled?

A

When a sub-shell is built up with electrons, each orbitals filled singly before pairing starts.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

Is there any anomalies to the filling of sub-shells?

A

The 4s orbital is at a slightly lower energy level than the 3d orbital. This means 4s will fill before 3d.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

Electron configuration, what form is each occupied shell written in?

A

Each occupied shell is written in the form nXy where;
n is the shell number
X is the type if orbital
y is the number of electrons in the orbitals making up the sub shells.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

When are ions formed?

A

Ions are formed when atoms loose or gain electrons.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

If an atom is ionised to form a positive ion which electrons are affected?

A

The electrons found in the highest energy level list first.

The electron configuration will therefore show fewer electrons in the highest energy levels.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

If an atom is ionised to form a negative ion which electrons are affected?

A

When atoms gain electrons to become a negative ion, the extra electrons will continue to fill the sub shells.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

Give the electronic configuration of lithium and a lithium ion.

A

Li 1s2 2s1 Li+ 1s2

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
31
Q

When does ionic bonding occur?

A

Between a metal and a non metal

Electrons transferred from metal atom to non metals, to form oppositely charged ions that attract each other.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
32
Q

Describe a covalent bond.

A

Consists of two non-metals
Electrons shared between the atoms and attracted to the nuclei of both bonded atoms.
Elements involved achieve Nobel gas configurations.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
33
Q

Describe metallic bonding.

A

Occurs in metals electrons shared between all the atoms

Includes all metals and their alloys.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
34
Q

What do all ionic compounds exist as?

A

All ionic compounds exist as giant ionic lattices in the solid state, and when they are in the liquid state all the ions are free to move around.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
35
Q

Describe the properties of ionic compounds.

A

High melting and boiling points.
Ionic compounds are soiled at room temperature, a large amount of energy is needed to break the strong electrostatic bonds that hold oppositely charged ions together in the solid lattice.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
36
Q

Why would the melting points of MgO be higher than the melting point of NaCl?

A

The charges of Mg2+ and O2- ions are greater than those on Na+ and Cl- . The greater the charge, the stronger the electrostatic forces between the ions.

37
Q

Describe the solubility of a ionic lattice.

A

An ionic lattice dissolves in polar solvents such as water.
Polar molecules break down an ionic lattice by surrounding each ion to form a solution.
The slight changes within a polar substance are able to attract the charged ions in the giant ionic lattice.
This means that lattice is disrupted and ions are pulled out of it.

38
Q

What allows a covalent bond to form?

A

The attraction of the electron pair and the nuclei overcomes the repulsion between to positively charged nuclei.

39
Q

When does a dative covalent bond form?

A

When all the shared electrons come from one atom. For example NH3 may react with a H+ ion to form a dative bond and become NH4.

40
Q

What is the oxonium ion?

A

When acid is added to water, water molecules form oxonium ions.
HCl (g) + H2O -> H3O+ (aq) + Cl-(aq)

41
Q

What are oxonium ions responsible for?

A

H3O+ ions are responsible for reactions of acids. In equations , the oxonium ion is often simplified to H+.

42
Q

What is the octet rule?

A

When a covalent bond forms, unpaired electrons often pair up so that the bonded atoms obtain a Nobel gas electron configuration obeying the Octet rule.

43
Q

Why might the octet rule not be possible?

A

Not enough electrons for an octet

Expansion of the octet

44
Q

Describe why the octet rule may not be able to be achieved because there are not enough electrons for an octet.

A

Within period 2, the elements beryllium, Be, and Boron, B, both form compounds with covalent bonds. Neither of these elements has enough unpaired electrons to reach a novel gas electron configuration.

45
Q

Describe expansion of the octet.

A

As we move down the periodic table more of the outer shell electrons are able to take part in bonding.
In the resulting molecules, one of the bonding atoms may finish up with more than 8 electrons in its outer shell.
This breaks the octet rule and is often called expansion of the octet. (Sulphur)

46
Q

How does expansion of the octet modify the octet rule?

A

Unpaired electrons pair up, the maximum number of electrons that pair up is equivalent to the number of electrons in the outer shell.

47
Q

How are simple covalent structures held together?

A

By strong covalent bonds,

Different molecules held together by weak intermolecular forces such as London forces - low melting and boiling points.

48
Q

What are the properties of simple molecular structures?

A

No free particles to move - non conductors of electricity.
Simple molecules are usually soluble in non-polar solvents, such as hexane -> weak London forces are able to form between covalent molecules and these solvents.

49
Q

What are the properties of giant covalent structures (diamond, graphite and SiO2)?

A

Strong covalent bonds within the lattice -> high melting and boiling points.
Giant covalent structures are non-conductors of electricity because there are no free charged particles apart from graphite - delocalised electrics between layers.
Giant covalent structures are insoluble in both polar and non polar solvents because the covalent bonds in the lattice are too strong to be broken by either polar or non polar solvents.

50
Q

What are the shapes of molecules and ions determined by?

A

Determined by the number of electron pairs in the outer shell surrounding the central atom.
As electrons all have a negative charge, each electron pair repels other electron pairs.
The shape adopted will be the shape that allows electron pairs to be as far apart as possible.

51
Q

Give and example of, and the bond angle of a linear molecule.

A

CO2, 180 degrees

52
Q

Give and example of, and the bond angle of a trigonal planar.

A

BF3, 120 degrees.

53
Q

Give and example of, and the bond angle of a tetrahedral molecule.

A

CH4, 109.5 degrees.

54
Q

Give and example of, and the bond angle of a trigonal bipyramidal molecule.

A

PCl5, 90 degrees + 120 degrees.

55
Q

Give and example of, and the bond angle of a octahedral molecule.

A

SF6, 90 degrees

56
Q

Why do molecules with lone pairs affect the bond angles?

A

A lone pair of electrons is slightly more electron-dense than a bonded pair -> a lone pair repels more than a bonded pair.

57
Q

What are the relative strengths of repulsion between electron pairs?

A

Lone pair / lone pair > bonded pair / line pair > bonded pair / bonded pair

58
Q

What are the bond shapes of molecules with line pairs.

A

Tetrahedral with bond angles of 109.5 - methane
Pyramidal - ammonia molecule -107
Non linear - water - 104.5

59
Q

What is electronegativity?

A

In 1932, the US chemist Linus Pauling invented the Pauling scale - to measure electronegativity of an atom.

60
Q

Describe polarity in a 100% polar bond.

A

Each hydrogen atom has an equal share of the pair of electrons in the bond, resulting in a perfect 100% covalent bond.
The nucleus of each bonded atom is equally attracted to the bonded electron pair.

61
Q

In what situation does a non-polar, 100% covalent bond occur?

A

If the bonding atoms are different, one if the atoms is likely to be attract the bonding electrons more.
If the bonding atoms are different, one if the atoms is likely to attract the bonding electrons more.
The bonding atom with a greater attraction for the electron pair is said to be more electronegative than the other atom.

62
Q

Describe the polarity if a HCl- molecule .

A

Cl is more electronegative than H.
Cl has greater attraction for electrons, (bonding pair)
The bonding electrons held closer to Cl causing a permanent dipole.
HCl is non-symmetrical and has a charge difference across the whole molecule.

63
Q

How does the shape of a molecule affect polarity?

A

CCl4 and CO2 are symmetrical and are non polar because the pull of any one electronegative Cl- is cancelled out by the pull if the others.
H2O is non symmetrical and has a charge difference across the whole molecule making it polar.

64
Q

What are the different intermolecular forces and their strength?

A

Ionic and covalent - 1000
Hydrogen - 50
Permanent - 10
London (dispersion) forces - 1

65
Q

What types of dipole-dipole interactions are there?

A

Permanent dipole - induced dipole interactions and permanent dipole - permanent dipole interactions.

66
Q

What causes London forces?

A

London forces are caused by the constant random movement of electrons in atoms shells. This movement unbalances the distribution of charge within electron shells.
At any moment, there will be an instantaneous dipole across the molecule.
The instantaneous dipole induces a dipole in neighbouring molecules.
The small induced dipoles attract one another, causing weak intermolecular forces known as London (dispersion) forces or instantaneous dipole - induced dipole forces.

67
Q

What may change the size of London forces?

A

The size of London forces increases with increasing numbers of electrons.

68
Q

What is hydrogen bonding?

A

Molecules containing O-H, N-H and F-H bonds are permanent with permanent dipoles. The electro positive hydrogen atoms attract the electronegative nitrogen, oxygen or fluorine atoms.

69
Q

Why is water able to form hydrogen bonds?

A

Water is able to form hydrogen bonds because the O has one lone pair on it. These can form a hydrogen bond with the H atom on another water molecule, as it is electron deficient.

70
Q

How are hydrogen bonds formed between ammonia and water?

A

The lone pairs on the electronegative O atom are attracting the H on the molecule of ammonia. This H atom is very electron deficient because it is attached to N, which is highly electronegative and has love pairs on it.

71
Q

What is the effect of hydrogen bonding on the properties of water?

A

Ice is less dense than water.

Water has an unusually high boiling point compared to other group 16 hydrides.

72
Q

Why is ice less dense than water?

A

When I’ve forms, water molecules arrange themselves into an orderly patten and hydrogen bond form between the molecules. (This will happen in the liquid phase but not as often as liquid molecules can move past each other and overcome the bonds)
Ice has an open lattice with hydrogen bonds holding water molecules apart.
When ice melts, the rigid hydrogen bonds collapse, allowing the H2O molecules to move closer together.

73
Q

Why does water have an unusually high boiling point compared to other group 16 hydrides?

A

Water is able to form hydrogen bonds, which are much stronger than other intermolecular forces.
Water has higher surface tension and viscosity.

74
Q

What is periodicity?

A

Periodicity is the trend in properties that is repeated across each period.

75
Q

Describe periodicity across the periodic table.

A

Across each period elements change from metals to non metals.
Similar elements are placed in vertical groups, they are similar due to the same number of elements being in the outer shell.

76
Q

What is an exception to the filling of energy levels in the periodic table?

A

The 4s every level is lower than the 3rd energy level.
The 4s orbital fills before the 3d orbital.
The 4s orbital would be emptied before the 3rd orbital during ionisation.

77
Q

Evidence for election shells;

Give an example of where we can see the effect of electron shells.

A

Plasma TV
Ne(g) -> Ne+(g) + e-
Xe(g) -> Xe- (g) + e-

78
Q

What is the first ionisation energy a measure of?

A

The first ionisation energy is a measure of how easily an atom looses an electron to form a 1+ ion.
(Na -> Na+(g) + e-)

79
Q

What is the first ionisation energy?

A

The firs ionisation energy of an element is the energy required to remove one electron from each atom in one molecule of the gaseous element to form one mole of gaseous 1+ ions.

80
Q

How does an atom form a positive ion?

A

To form a positive ion, energy must be supplied to an electron to overcome this attraction.

81
Q

When an atom forms a positive ion, which atoms are ionised first?

A

Electrons in the outer shell are removed first because they experience the smallest nuclear attraction. (the outer shell electrons are furthest away from the nucleus and require the least ionisation energy.

82
Q

What does the nuclear attraction experienced by an electron depend on?

A

Atomic radius
Nuclear charge - the higher the nuclear charge, the larger the attractive force on the outermost electrons.
Electron shielding or screening

83
Q

What is electron electron shielding?

A

The more inner shells there are, the larger the shielding effect and the smaller the nuclear attraction experienced by the outer electrons.

84
Q

What are successive ionisation energies?

A

Successive ionisation energies is a measure of the amount of energy required to remove each electron in turn.

85
Q

Why is each successive ionisation energy higher than the one before?

A

As each electron is removed there is less repulsion between remaining electrons, and each shell will be drawn slightly closer to the nucleus.
The positive nuclear charge will outweigh the negative charge every time an electron is removed.
As the distance of each electron from the nucleus decreases slightly, the nuclear attraction increases. More energy is needed to remove each successive electron.

86
Q

How did the Bohr model define energy levels?

A

Successive ionisation energies are reflecting the leap from one shell to another getting closer to the nucleus. On a ionisation energy graph there is a sharp energy as electrons are withdrawn from the level below.

87
Q

Periodicity; what are the trends across a period?

A

The number of protons in the nucleus increases, so there is a higher attraction on the electrons.
electrons are added to the same shell, so the outer shell is drawn slightly inward.
There is the same number of inner shells so electron shielding hardly changes.
Moving across a period the attraction between the nucleus and the outer electrons increases, so more energy is needed to remove an electron. (first ionisation energy increases across a period).

88
Q

Between which groups is there a small decrease in ionisation energies.

A

Decrease between groups 2 to 13 and 15 to 16.

Chemistry AS OCR (7 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions