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EDEXCEL CHEMISTRY AS LEVEL > TOPIC 2 BONDING AND STRUCUTRE > Flashcards

TOPIC 2 BONDING AND STRUCUTRE Flashcards

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1
Q

Define ionic bonding.

A

The bonding between oppositely charged ions held together by electrostatic attractions.

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2
Q

What is the structure of an ionic compound?

A

Giant lattice

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3
Q

What are the properties of ionic compounds?

A
  • Dissolve in water or aqueous.
  • High melting point
  • Brittle
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4
Q

Why are ionic compounds brittle?

A

When struck with a hammer, the layers slide and the alike ions become close to each other, hence repel and the structure breaks apart.

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5
Q

Why do ionic compounds have high melting points?

A

They have strong electrostatic forces between oppositely charged ions. Thus lots of energy is needed to overcome the forces.

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6
Q

Why do ionic compounds conduct electricity when molten/aqueous?

A

Ions are free to move around.

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7
Q

Why do ionic compounds dissolve in water?

A

Water is polar so attract the anions and the cations which then break up the structure.

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8
Q

How does the size of the charge on the ion affect the strength of an ionic bond?

A

The bigger the charge on an ion, the stronger the electrostatic attraction between ions. Thus more energy is required to overcome these forces and hence having high melting and boiling points.

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9
Q

Which of the 2 ionic compounds has a higher melting point? CaO or KCl?

A

CaO because it has a larger charge difference==> Ca2+ O2-

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10
Q

How does the ionic radii affect the strength of the ionic bond?

A

The smaller the size of the ion, the stronger the electrostatic attraction between ions, as they can pack closely together hence more energy is required to overcome those forces.

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11
Q

What is a feature we see if an ionic compound has a high charge density?

A

Smaller the ionic radii higher the charge the stronger the electrostatic attraction.

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12
Q

Define an isoelectronic ion.

A

Different atoms that have the same number of electrons.

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13
Q

What happens to the radius of isoelectronic ions as the atomic number increases? Why?

A

Decreases because the attractive force from the nucleus increases as it has more protons which pulls the outer shell electron in more.

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14
Q

What is the evidence for charged particles? Describe the experiment.

A

The electrolysis of Copper (II)Chromate(VI) on wet filter paper.
When a wet filter paper on a glass slide is placed the positive and negative electrodes are connected, a drop of CrCuO4 is added and electricity passes through the ions causing it to separate. The cations moves towards the anode and the cations move towards the cathode. Now you can see the a yellow and blue solution, thus showing the existence of charged particles.

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15
Q

Define covalent bonding.

A

Sharing of outer electrons in order form atoms to obtain a full shell.

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16
Q

What is a covalent bond?

A

The strong electrostatic attraction between 2 nuclei and the shared pair of electrons between them.

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17
Q

What is a dative bond?

A

Where one atom donates 2 electrons to an atom or an ion to form a bond.

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18
Q

Draw the dot cross diagram of ammonia and carbon monoxide.

A

https: //cdn.savemyexams.co.uk/wp-content/uploads/2020/11/1.3-Chemical-Bonding-Dative-Covalent-Bonding-Ammonium-ion.png
https: //www.chemguide.co.uk/atoms/bonding/co.GIF

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19
Q

In covalent molecules what are the forces of attraction between?

A

Positive nuclei and negative electrons being shared.

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20
Q

What is bond enthalpy linked to? How?

A

The length of a bond, so the shorter the bond, the higher the bond enthalpy.

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21
Q

Where are the repulsive forces in a covalent bond?

A

Between the 2 positive nuclei and between the electrons involved and not involved in the bonding.

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22
Q

Define bond length.

A

The balance between the attractive and repulsive forces in a covalent bond.

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23
Q

How can covalent bond have a high bond enthalpy?

A

The greater the density between the atoms, the stronger the attractive force meaning that the atoms are pulled in further towards each other hence leading to a shorter bond which leads to a higher bond enthalpy.

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24
Q

Between a single, double and triple bond, which one has the highest bond enthalpy and why?

A

The triple bond because it has a higher electron density meaning that there is a stronger attractive force thus resulting in a shorter bond and a higher bond enthalpy.

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25
Q

Why do molecules have specific shapes and angles?

A

Bonds repel each other equally as they contain electrons which wants to stay as far apart as possible.

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26
Q

Describe the repulsion of a lone next to a bond pair compared to that of 2 lone pairs.

A

A lone pair next to a bond pair would repel further than that of 2 bond pairs, however, 2 lone pairs would repel even further.

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27
Q

How is the shape of a simple molecule determined?

A

The repulsion between the electron pairs that surround a central atom.

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28
Q

What do lone pairs do? What is the rule?

A

Lone pairs change the shape and bond angles by pushing bonding pairs closer together. For every lone pair reduce the bond angle by 2.5 degrees.

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29
Q

Linear shape.

Draw the diagram.

A

2 BP
0 LP
180°
BeCl2

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30
Q

Trigonal planar shape.

Draw the diagram.

A

3 BP
0 LP
120°
BF3

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31
Q

Tetrahedral shape.

Draw the diagram.

A

4BP
0LP
109.5°
CH4

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32
Q

Trigonal bipyramidal shape.

Draw the diagram.

A

5BP
0LP
120° and 90°
PCl5

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33
Q

Octahedral shape.

Draw the diagram.

A

6BP
0LP
90°
SF6

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34
Q

Trigonal Pyramidal shape.

Draw the diagram.

A

3BP
1LP
107°
NH3

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35
Q

Bent shape.

Draw the diagram.

A

2BP
2LP
104.5°
H2O

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36
Q

Distorted T shape.

Draw the diagram.

A

3BP
2LP
87.5°
ClF3

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37
Q

Seesaw shape.

Draw the diagram.

A

4BP
1LP
87° and 102°
SF4

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38
Q

Square Pyramidal shape.

Draw the diagram.

A

5BP
1LP
81.9° and 90°
IF5

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39
Q

Square Planar shape.

Draw the diagram.

A

4BP
2LP
90°
XeF4

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40
Q

What are some examples of giant covalent structures?

A

Graphite, Diamond, Graphene and Silicon Dioxide.

41
Q

Describe the bonding of graphite.

A

Each carbon is bonded 3 times and the 4th electron is delocalised which forms layers.

42
Q

Why does graphite have a very high melting point.

A

Many strong covalent bonds.

43
Q

Why does graphite conduct electricity?

A

There are delocalised electrons between the layers which carry a charge.

44
Q

Why is graphite a good lubricant?

A

Its layers slide easily as there are weak forces between the layers.

45
Q

Why does graphite have a low density?

A

The layers are far apart in comparison to the covalent bond length.

46
Q

Why is graphite insoluble?

A

The covalent bonds are too strong to break.

47
Q

Describe the bonding of diamond.

A

Each carbon is bonded 4 times in a tetrahedral shape.

48
Q

Why is diamond a good conductor?

A

It has a tightly packed rigid arrangement.

49
Q

Why does diamond have a high melting point?

A

It has many strong covalent bonds so lots of energy is needed to overcome those bonds.

50
Q

Why doesn’t diamond conduct electricity?

A

It doesn’t have any delocalised electrons.

51
Q

Why is diamond insoluble?

A

The covalent bonds are too strong to break.

52
Q

Describe the structure of graphene.

A

1 layer thick as it is one atom thick and made up of hexagonal carbon rings.

53
Q

Why is graphene lightweight and transparent?

A

Only one cell thick.

54
Q

Why is graphene a good conductor of electricity?

A

Has delocalised electrons which carry charge.

55
Q

What strengthens the covalent bonds of graphene?

A

The delocalised electrons hence it is very strong.

56
Q

What are 3 uses of graphene?

A

Aircraft shells
Super computers
Smart phone screens.

57
Q

What is the structure of metals?

A

Giant metallic lattice structure.

58
Q

Describe metallic bonding.

A

Metal atoms which donate electron to form a sea of delocalised electrons with positive metal ions.

59
Q

Where is the electrostatic attraction in metals?

A

Between the positive metal ions and the negative delocalised electrons.

60
Q

Why does magnesium have a higher melting point that sodium?

A

Sodium is a 1+ ion so can only donate 1 electron whereas magnesium is a 2+ ion so can donate 2 electrons.

61
Q

Why are metals good thermal conductors?

A

Delocalised electrons can transfer kinetic energy.

62
Q

Why are metals good electrical conductors?

A

Delocalised electrons are mobile and can carry a current.

63
Q

Why do metals have high melting points?

A

They have strong electrostatic attractions.

64
Q

Why are metals insoluble?

A

They have strong electrostatic attractions , thus the metallic bonds are too hard to break.

65
Q

Why metals malleable and ductile?

A

The ion layers can slide when hit with a hammer but still retain an attraction between the ions and delocalised electrons.

66
Q

Define elctronegativity.

A

The ability for an atom to attract electrons towards itself in a covalent bond.

67
Q

From the periodic table, how do you know if an element is more or less electronegative?

A

The further up and right you go excluding the noble gases, the more electronegative the element.

68
Q

What is the most electronegative element?

A

Fluorine.

69
Q

What happens if a compound has a big difference in electronegativity?

A

The compound will be more ionic.

70
Q

What happens if a compound has an electronegativity difference of 2?

A

It is purely covalent.

71
Q

How can covalent bonds become polar?

A

If the atoms attached to it have a difference in electronegativity.

72
Q

How do you know if a bond is more or less polar?

A

The bigger the difference in electronegativity the more polar the bond will be.

73
Q

What can lead to polar molecules?

A

The uneven distribution of charge.

74
Q

What 2 factors determine if a molecule is NOT polar?

A

Atoms that are bonded with the same or similar electronegativity value and if the polar bonds are arranged symmetrically.

75
Q

Where do london forces exist between?

A

Atoms and molecules.

76
Q

What is the weakest intermolecular force?

A

London forces

77
Q

What type of electrons can form a dipole?

A

Any molecule or atoms with electrons.

78
Q

How is a temporary dipole created?

A

When one molecule goes closer to another molecule and the electron clouds will repel each other, thus forming a temporary dipole.

79
Q

When does a temporary dipole only exist?

A

When 2 molecules or atoms are near each other. When they move away the dipole interaction is destroyed.

80
Q

How are london forces created?

A

Caused by an uneven distribution of electrons within an atom thus resulting in a δ+ and δ- charge on either side of the atom. A temporary dipole is created and can now induce a temporary dipole on a neighbouring atom. The london forces are the electrostatic attractions set up between the δ+ end of one atom and the δ- of the other.

81
Q

How do you know if you will have more or less london forces in an atom or a molecule?

A

The bigger the molecule the forces as you have more electron clouds.

82
Q

What happens when we boil a liquid?

A

The weak london forces are breaking.

83
Q

Why do branched hydrocarbons have low boiling points?

A

They can’t pack close together hence weakening the london forces between the chains.

84
Q

Where do permanent dipole-dipole interactions exist?

A

Exist in molecules with a polarity.

85
Q

Describe the permanent dipole-dipole forces.

A

The δ- on one molecule is attracted to the δ+ in another, but unlike london forces they are not temporary so this force is stronger than that of london.

86
Q

How do we test for polar molecules?

A

Place a charged rod near a stream of polar liquid and the liquid should bend towards the rod as the molecules align to face the oppositely charged rod.

87
Q

What is the strongest type of intermolecular force?

A

Hydrogen bonding.

88
Q

When does hydrogen bonding occur?

A

When you have very electronegative elements.

89
Q

Which 3 elements does hydrogen bonding occur?

A

Nitrogen, Oxygen and fluorine.

90
Q

Draw the hydrogen bonding of water.

A

http://www.chem.ucalgary.ca/courses/350/Carey5th/Ch27/ch27-4-4.html

91
Q

Why is water less dense as a solid?

A

As water freezes, hydrogen bonds start to form, pushing the molecules further apart, so there is less water molecules in a given volume therefore less dense as a solid.

92
Q

Why are alcohol not as volatile as alkanes?

A

Alcohol has strong hydrogen bonds where as alkanes has weak london forces so less energy is needed for the alkanes to turn into a gas.

93
Q

From HCl to HI why does the boiling points increase?

A

It has a increased mass therefore, it has a bigger electron cloud so more london forces.

94
Q

Why do most ionic compounds dissolve in polar solvents?

A

Because the δ+ of the solvent attracts the negative ions of the compound and the δ- attracts the positive ions and the structure starts to break down.

95
Q

What is hydration?

A

When a molecule dissolves in water the ions are surrounded by the water molecules.

96
Q

How does hydration occur?

A

The new bonds formed must be the same strength or greater than those that are broken.

97
Q

Why can’t Al2O3 not dissolve in water?

A

Its ionic bonds are stronger than that of water molecules.

98
Q

Why can alcohols dissolve in polar substances? Draw the diagram of methanol and water.

A

They can form a hydrogen bond with the water molecules. But the hydrogen part cannot.
44:23

99
Q

Why do alkanes dissolve best in non polar molecules?

A

They can form london forces between the molecules.

EDEXCEL CHEMISTRY AS LEVEL (2 decks)

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