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A-level Chemistry > Topic 2: Bonding > Flashcards

Topic 2: Bonding Flashcards

Ionic, Covalent, Metallic, shapes of molecules, properties of water

1
Q

What is an ionic bond

A

The electrostatic force pf attraction between oppositely charged ions

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2
Q

What is the effect of ionic radius on the strength of ionic bonds?

A

The smaller the ion, the stronger the electrostatic attraction between ions. Smaller ions can be more closely packed together, so more energy is required to overcome stronger forces.

Smaller ionic radius means the ion has a higher charge density.

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3
Q

What is the effect of ionic charge on the strength of ionic bonds?

A

The bigger the charge on an ion, the stronger the electrostatic attraction between ions, and the higher the melting point.
e.g. K+ and Cl- = 770 C
vs Ca2+ an O2- = 2572 C

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4
Q

Solubility of ionic compounds

A

Most dissolve in water.
Water molecules are polar, so attract the positive and negative ions in the ionic compound

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5
Q

Conduction of ionic compounds

A

Can conduct when molten / aqueous as the ions are free to move around

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6
Q

Melting points of ionic bonds

A

Very high melting points.
Many strong electrostatic forces of attraction between oppositely charged ions - acting in all directions.

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7
Q

Qualities of ionic compounds

A

Hard + Brittle - layers slide past each other and positive ions align with each other and repel

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8
Q

What are isoelectric ions?

A

Different ions with the same number of electrons

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9
Q

Trend in ionic radius of isoelectric ions

A

Ionic radius decreases as the atomic number increases.
e.g.
N2- O2- F- Na+ Mg2+ Al3+
Larger —————-> Smaller
The nuclear charge increases, and so has stronger attraction to the outer electron shell, pulling it in more, decreasing the size of ionic radius.

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10
Q

What is the evidence for charged particles?

A

Electrolysis of Copper(II) Chromate(VI)
1. Green copper chromate is dropped on to wet filter paper.
2. Electricity is passed through - ions begin to separate
3. Positive Cu 2+ ions move to negative electrode - blue solution
4. Negative CrO4 2- ions move towards positive electrode - yellow solution

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11
Q

Covalent bonds

A

The strong electrostatic attraction between the two nuclei and the shared pair of electrons between them.

The sharing of outer electrons in order for atoms to obtain a full outer shell

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12
Q

What is a dative / co-ordinate bond?

A

Where one atom donates both electrons to another atom / ion to form a bond
e.g. NH4 +
(Shown by an arrow)

Carbon monoxide - has a double covalent bond as well as dative bond

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13
Q

Effect of bond length on bond enthalpy

A

The shorter the bond, the higher the bond enthalpy.
This is because a higher electron density between atoms leads to a stronger attractive force. Atoms are puled further towards each other. Forming a shorter bond + higher bond enthalpy

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14
Q

Effect of bond number on bond enthalpy

A

A triple covalent bond ha a much higher electron density than a single covalent bond. Therefore a triple bond has a shorter bond length and higher bond enthalpy

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15
Q

What does the shape of a molecule depend on?

A

electron pairs arrange themselves around the central atom to:
1. Minimise repulsion
2. Maximise separation

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16
Q

What shape and angle does a molecule with 2 bonding pairs have

A

Linear
180

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17
Q

What shape and angle does a molecule with 3 bonding pairs have

A

Trigonal planar
120

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18
Q

What shape and angle does a molecule with 4 bonding pairs have

A

Tetrahedral
109.5

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19
Q

What shape and angle does a molecule with 5 bonding pairs have

A

Trigonal Bipyramidal
90 + 120

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20
Q

What shape and angle does a molecule with 6 bonding pairs have

A

Octahedral
90

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21
Q

What shape and angle does a molecule with 3 bonding pairs and 1 lone pair have

A

Trigonal Pyramidal
107

22
Q

What shape and angle does a molecule with 2 bonding pairs and 2 lone pairs have

A

Bent
104.5

23
Q

What shape and angle does a molecule with 3 bonding pairs and 2 lone pairs have

A

T-shape / trigonal planar
87.5

24
Q

What shape and angle does a molecule with 4 bonding pairs and 1 lone pair have

A

Seesaw
87 + 102

25
Q

What shape and angle does a molecule with 5 bonding pairs and 1 lone pair have

A

Square Pyramidal
90 + 81.9

26
Q

What shape and angle does a molecule with 4 bonding pairs and 2 lone pairs have

A

Square Planar
90

27
Q

What is electronegativity?

A

The ability of an atom to attracts the bonding pair of electrons from a covalent bond.

28
Q

What does the difference in electronegativity of atoms mean?

A

A small to 0 difference in En values means the compound is covalently bonded.
A larger difference in En values shows some ionic character.

29
Q

Trend in electronegativity across a period

A

Increases
Nuclear charge increases
Shielding remains relatively constant across the period as no new shells are being added
Nucleus attracts outer electrons
This results in smaller atomic radii
Therefore greater electronegativity

30
Q

Trend in electronegativity down a group

A

Decreases
More shells are added
Atomic radius increases

31
Q

What is a non-polar bond

A

When two atoms in a bond have the same electronegativities e.g. Cl2

32
Q

What electronegativity value implies an ionic bond

A

A difference of more than 1.7

33
Q

What electronegativity value implies a polar bond

A

A difference between 0.3 - 1.7

34
Q

Explain how molecules with polar bonds are not overall a polar molecule

A

Some molecules have polar bonds but are overall not polar because the polar bonds in the molecule are arranged in such way that the individual dipole moments cancel each other out e.g. CCl4 or CO2

35
Q

What are intramolecular forces?

A

Forces within a molecule e.g. covalent bonds

36
Q

What are intermolecular forces?

A

Forces between molecules e.g. Hydrogen bonds, London forces, Permanent dipoles

37
Q

How are London Dispersion forces formed?

A

Fluctuations in the electron cloud causes one side of a bond to have more or less electrons.
This causes a temporary dipole to arise
The temporary dipole induces a dipole on neighbouring molecules
The δ+ end of one dipole is attracted to the δ- end of the next dipole, forming a London force between them

38
Q

What are permanent dipole - dipole forces?

A

Forces between two molecules that have permanent dipoles are called permanent dipole - dipole forces
The Forces between two molecules that have permanent dipoles are called permanent dipole - dipole forces
The δ+ end of the dipole in one molecule and the δ- end of the dipole in a neighbouring molecule are attracted towards each other

39
Q

What are Hydrogen bonds?

A

Hydrogen bonding is a type of permanent dipole – permanent dipole bonding
When hydrogen is covalently bonded to an O, N or F (very electronegative) the bond becomes highly polarised
The H becomes so δ+ charged that it can form a bond with the lone pair of an O, N or F atom in another molecule e.g. Water

40
Q

What is the strength of London forces?

A

Small molecules with the same number of electrons, permanent dipole interactions are stronger than induced dipoles

41
Q

Explain the melting and boiling points of water

A

Both are very high.
Strong intermolecular forces of hydrogen bonding between the molecules in both ice (solid H2O) and water (liquid H2O)
A lot of energy is therefore required to separate the water molecules and melt or boil them

42
Q

Why does water have a high surface tension?

A

Surface tension is the ability of a liquid surface to resist any external forces (i.e. to stay unaffected by forces acting on the surface)
The water molecules at the surface of liquid are bonded to other water molecules through hydrogen bonds
These molecules pull downwards the surface molecules causing the surface of them to become compressed and more tightly together at the surface
This increases water’s surface tension

43
Q

Properties of water: What is the relationship between the densities of ice and water?

A

The water molecules are packed into an open lattice in water.
This way of packing the molecules and the relatively long bond lengths of the hydrogen bonds means that the water molecules are slightly further apart than in the liquid form

44
Q

How does an increasing carbon chain effect the boiling point of alkanes?

A

As the length increases, the boiling point increases. This is because there are more London forces between molecules with longer chains.

45
Q

How does increased branching effect the boiling point of alkanes?

A

Boiling point decreases as branching increases.
Branching prevents molecules from being close to each other and reduces the strength of London forces

46
Q

Explain the volatility of alcohols and alkanes of similar electron number

A
47
Q

Explain the trend in boiling points of hydrogen halides

A

HF: Highest boiling point as it has hydrogen bonds which are the strongest type of intermolecular force.
HCl - HBr - HI: Increases slightly from HCl. This is because there is an increase in electron number, so the strength of London forces between molecules increases.

48
Q

What is a metallic bond?

A

Strong electrostatic force of attraction between positively charged metal ions and delocalised electrons

49
Q

Explain the solubility of compounds with hydrogen bonding or organic compounds containing N or O, with water

A

Compounds with hydrogen bonding can form hydrogen bonds with water.
e.g. Alcohols
When hydrogen bonds are formed between alcohols and water, enough energy is released to break H bonds between water molecules and between alcohol molecules

50
Q

Explain the solubility of non-polar substances with water

A

They cannot form hydrogen bonds with water, so do not release energy to disrupt the bonding between H2O molecules, so do not dissolve

A-level Chemistry (26 decks)

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