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A Level Chemistry > Module 2: Chapter 6 - Shapes of Molecules and Intermolecular Forces > Flashcards

Module 2: Chapter 6 - Shapes of Molecules and Intermolecular Forces Flashcards

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1
Q

What is Electron-pair repulsion theory?

A

EPRT is a model used in explaining and predicting the shapes of molecules and polyatomic ions

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2
Q

Explain Electron-pair repulsion theory:

A

The electron pairs surrounding a central atom of a molecule determine the shape of the molecule or ion. As all electrons are negatively charged, the electron pairs repel one another so that they are arranged as far apart as possible. The arrangement of electron pairs minimies repulsion and thus holds the bonded atoms in a definite shape. Different numbers of electron pairs result in different shapes.

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3
Q

How is a bond in the plane of the paper represented when drawing a 3d molecule?

A

A solid line

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3
Q

How is a bond going into the plane of the paper represented when drawing a 3d molecule?

A

A dotted wedge

(insert image once brainscape decides to work)

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4
Q

Why does a lone pair repel more strongly than a bonding pair?

A

A lone pair of electrons is closer the the central atom and occupies more space than a bonded pair. This results in a lone pair repeling more strongly than a bonding pair

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5
Q

What is the order of relative electron pair respulsions? (from highest repulsion to lowest)

A
  1. Lone-pair/Lone-pair
  2. Lone-pair/Bonded-pair
  3. Bonded-pair/Bonded-pair
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6
Q

What are the 6 type of molecule shape?

A
  • Linear
  • Non-linear
  • Trigonal planar
  • Pyramidal
  • Tetrahedral
  • Octahedral
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7
Q

What is the molecular shape of methane?

A

Tetrahedral

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8
Q

What is the bond angle of methane?

A

109.5°

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9
Q

What is the moleular shape of ammonia?

A

pyramidal

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10
Q

What is the bond angle of ammonia?

A

107°

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11
Q

What is the molecular shape of water?

A

non-linear

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12
Q

What is the bond angle of water?

A

104.5°

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13
Q

What does lone pairs repeling more strongly cause?

A

Lone pairs repel more strongly than bonded pairs, therefore, lone pairs repel bonded pairs slightly closer together, decreasing the bond angle.

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14
Q

What is the bond angle?

A

The angle between the bonded pairs of electrons

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15
Q

How much does each lone pair reduce the bond angle by?

A

2.5°

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16
Q

How are multiple bonds treated when determining the molecular shape?

A

They are treated as “bonding regions” which are treated the same as single bonds

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17
Q

What happens to the bond angle if there is a greater number of electron pairs?

A

It decreases

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18
Q

What is the bond angle in linear molecules?

A

180°

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19
Q

What is the bond angle in trigonal planar molecules?

A

120°

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20
Q

What is the bond angle in tetrahedral molecules?

A

109.5°

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21
Q

What is the bond angle in (trigonal) pyramidal molecules?

A

107°

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22
Q

What is the bond angle in octahedral molecules?

A

90°

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23
Q

What is the molecular shape of a molecule with 2 electron pairs/regions around a central atom and no lone pairs?

A

Linear

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24
Q

What is the molecular shape of a molecule with 3 electron pairs/regions around a central atom and no lone pairs?

A

Trigonal Planar

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25
Q

What is the molecular shape of a molecule with 4 electron pairs/regions around a central atom and no lone pairs?

A

tetrahedral

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26
Q

What is the molecular shape of a molecule with 4 electron pairs/regions around a central atom of which 1 is a lone pair?

A

(trigonal) pyramidal

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27
Q

What is the molecular shape of a molecule with 4 electron pairs/regions around a central atom of which 2 are lone pairs?

A

nonlinear

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28
Q

What is the molecular shape of a molecule with 6 electron pairs/regions around a central atom and no lone pairs?

A

octahedral

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29
Q

How is a bond coming out of the plane of the paper represented when drawing a 3d molecule?

A

A solid wedge

(insert image once brainscape decides to work)

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30
Q

SKIP

A

SKIP

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31
Q

What is the molecular shape of a molecule with 3 electron pairs/regions around a central atom of which 1 is a lone pair?

A

non-linear

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32
Q

What is the molecular shape and bond angle of SiH₄?

A

Tetrahedral, 109.5°

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33
Q

What is the molecular shape and bond angle of H₂CO?

A

Trigonal Planar, 120°

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34
Q

What is the molecular shape and bond angle of CS₂?

A

Linear, 180°

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35
Q

What is the molecular shape and bond angle of SO₂?

A

non-linear, 117.5°

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36
Q

What is the molecular shape and bond angle of H₃O⁺?

A

Pyramidal, 107°

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37
Q

What is the bond angle in most non-linear molecules?

A

104.5° (4 electron regions) or 117.5° (3 electron regions) (exception to both when there are double/triple bonds)

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38
Q

What is the shape of a linear molecule?

A
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39
Q

What is the shape of a trigonal planar molecule?

A
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40
Q

What is the shape of a tetrahedral molecule?

A
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41
Q

What is the shape of an octahedral molecule?

A
42
Q

What is the shape of a (trigonal) pyramidal molecule?

A
43
Q

What is the shape of a nonlinear molecule?

A
44
Q

What is the molecular shape of a molecule with 3 electron pairs/regions around a central atom of which 1 is a lone pair?

A

nonlinear

45
Q

What is electronegativity?

A

The ability of an atom to attract electron density in a covalent bond

46
Q

What is the range of difference in electronegativity for a non-polar covalent bond to be formed?

A

≤ 0.4

47
Q

What is the range of difference in electronegativity for a polar covalent bond to be formed?

A

0.4 - 1.8

48
Q

What is the range of difference in electronegativity for an ionic bond to be formed?

A

> 1.8

49
Q

Is there any charge separation in a non-polar covalent bond?

A

No charge separation

50
Q

Is there any charge separation in a polar covalent bond?

A

Yes, charge separation with permanent dipoles

51
Q

What are dipoles?

A

A separation in electrical charge so that one atom of a polar covalent bond, or one end of a polar molecule, has a small positive charge δ+ and the other a small negative charge δ-

52
Q

Why do noble gases have no electronegativity values?

A

They cannot form covalent bonds

53
Q

What factors affect electronegativity?

A
  • Nuclear Charge
  • Atomic Radius
  • Electron Shielding
54
Q

How does Nuclear charge affect electronegativity?

A

Increased nuclear charge (more protons in the nucleus) causes a stronger attraction between nucleus and bonding pair of electrons

55
Q

How does atomic radius affect electronegativity?

A

Decreased atomic radius causes electrons to be closer to the nucleus and therefore stronger attraction between nucleus and bonding pair of electrons

56
Q

How does electron shielding affect electronegativity?

A

Fewer shells of electrons between the nucleus and the electrons, causes less shielding (repulsion) against the bonding electrons, therefore stronger attraction beween the nucleus and bonding pair of electrons

57
Q

What is the trend in electronegativity across a period?

A

Electronegativity increases as:
* Atomic Radius decreases
* Greater Nuclear Charge
∴ stronger attraction between the nucleus and bonding pair of electrons

58
Q

What is the trend in electronegativity down a group?

A

Electronegativity decreases as:
* Atomic Radius Increases
* Electron Shielding Increases
∴ less attraction between nucleus and bonding pair of electrons

59
Q

What is the trend in atomic radius?

A
  • Atomic radius decreases across a period
  • Atomic Increases down a group
60
Q

What is a bond dipole moment?

A

A measure of the strength and direction of the polarity in a polar bond

61
Q

What type of bonds can a non-polar molecule contain?

A
  • non-polar bond
  • polar bonds where the dipole moments cancel each other out
62
Q

What intermolecular forces do non-polar molecules contain?

A

Induced dipole-dipole interactions (London Forces)

63
Q

What type of bonds does a polar molecule contain?

A

Polar Bonds where the dipole moments do not cancel out

64
Q

What intermolecular forces can polar molecules contain?

A
  • Induced dipole-dipole interactions (London Forces)
  • Permanent dipole-dipole interactions
  • Hydrogen Bonds
65
Q

How do you determine whether a molecule is polar or not?

A
  1. Draw a 3d sketch of the shape of the molecule
  2. Determine whether is contains polar bonds
  3. Determine whether it is a polar molecule (do the polar bonds cancel each other out?)
66
Q

What is electronegativity measured using?

A

The Pauling electronegativity values

67
Q

What is a pure covalent bond?

A

When the electron pair is shared equally (non-polar bond)

68
Q

Is PF3 polar or non-polar?

A

Polar

69
Q

Is BCl3 Polar or Non-polar?

A

non-polar

70
Q

Is Br2 Polar or non-polar?

A

non-polar

71
Q

Why does atomic radius decreases across a period?

A

As you travel across a period, effective nuclear charge increases, causing a greater attraction on the valence electrons, pulling the electron cloud closer to the nucleus, resulting in a decerased atomic radius

72
Q

What is effective nuclear charge?

A

The net charge that valence electrons experience

73
Q

Why does effective nuclear charge increase across a period?

A

As you travel across a period, the electrons shielding remains constant, however the nuclear charge increases. Therefore the effective nuclear charge increases

74
Q

What are the 3 types of intermolecular forces?

A
  • Induced dipole-dipole interactions (London Forces)
  • Permanent dipole-dipole interactions
  • Hydrogen Bonding
75
Q

What is the order of strength of intermolecular forces (from strongest to weakest)?

A
  1. Hydrogen Bonding
  2. Permanent dipole-dipole interactions
  3. London forces
76
Q

Explain how induced dipole-dipole interactions work

A

The random motion of electrons within a molecule produces a changing instantaneous dipole in a molecule. At any instant, an instantaneous dipole will exist, but its position will be constantly changing. This instantaneous dipole induces a dipole on a neighbouring molecule. The induced dipole induces further dipoles on neightbouring molecules, which then attract one another as opposite dipoles attract.

green = instantaneous dipole, blue = induced dipole
77
Q

What does the strength of London forces depend on?

A

The number of electrons in the molecules

78
Q

How does the number of electrons affect the strength of London forces?

A

The more electrons in each molecule:
* The larger the instantaneous and induced dipoles
* The greater the induced dipole-dipole interactions
* The stronger the attractive forces between molecules

79
Q

What term is used to refer to both induced and permanent dipole-dipole interactions?

A

Van der Waals’ forces

80
Q

Explain how permanent dipole-dipole interactions work

A

Permanent dipoles in different polar molecules interact with each other

81
Q

Where do permanent dipole-dipole interactions exist?

A

Between the permanent dipoles in different polar molecules

82
Q

Where do induced dipole-dipole interactions exist?

A

Between all molecules

But not in giant ionic or covalent structures!

83
Q

What are the steps to determine the relative boiling points of molecules?

A
  1. Does it have hydrogen bonding?
  2. Does it have permanent dipole-dipole interactions?
  3. What is the number of electrons in the molecule?
84
Q

What is a simple molecular substance?

A

A substance made up of simple molecules - small units containing definite number of atoms with a definite molecular formula

85
Q

What is the structure of simple molecular substances when solid?

A

Simple molecular lattice

86
Q

What are the boiling and melting points of simple substances?

A

In a simple molecular lattice, the weak intermolecular forces can be easily broken by the energy present at low temperatures. Theerefore they have low melting and boiling points

87
Q

What is hydrogen bonding?

A

A special type of permanent dipole-dipole interaction found between molecules containing a lone pair on a nitrogen, oxygen, or fluorine atom which is bonded to a hydrogen atom. The hydrogen bond acts between the lone pair on one molecule and a hydrogen atom in another molecule

88
Q

Whst is the shape around the hydrogen atom involved in a hydrogen bond?

A

Linear

89
Q

What are 2 anomalous properties of water?

A
  • Solid water (ice) is less dense than liquid water
  • Water has relatively high melting and boiling points
90
Q

Why is solid water (ice) is less dense than liquid water?

A

Each water molecule can form 4 hydrogen bonds, these hydrogen bonds extend outwards between each molecule, holding water molecules slightly apart and forming an open tetrahedral lattice full of holes. These holes in the open lattice structure decrease the density of water on freezing. When the ice melts, the lattice collapses and the molecules move closer together. Therefore, liquid water is more dense than solid ice

91
Q

Why does water have relatively high melting and boiling points

A

Water contains London forces, however it also contains strong hydrogen bonds. An appreciable quantity of energy is required to break these hydrogen bonds, therefore water has much higher melting and boiling points than would be expected from just London forces

91
Q

Why does water have relatively high melting and boiling points

A

Water contains London forces, however it also contains strong hydrogen bonds. An appreciable quantity of energy is required to break these hydrogen bonds, therefore water has much higher melting and boiling points than would be expected from just London forces

92
Q

What is the solubility rule?

A

Like dissolves Like

93
Q

What is the solubility of non-polar simple molecular substances:

A

Soluble in non-polar solvents

93
Q

Why are non-polar simple molecular substances soluble in non-polar solvents?

A

When added to a non-polar solvent, intermolecular forces form between the molecules and the solvent. The interactions weaken the intermolecular forces in the simple molecular lattice, therefore the intermolecular forces break and the compound dissolves

94
Q

Why are non-polar simple molecular substances insoluble in polar solvents?

A

When added to a polar solvent there is little interaction between the molecules in the lattice and the solvent molecules. The intermolecular forces in the polar solvent are too strong to be broken, therefore non-polar substances are insoulble in polar solvents

95
Q

What is the solubility of polar simple molecular substances:

A

Usually soluble in polar solvents

96
Q

Explain why polar solvents usually dissolve polar substances:

A

Polar covalent substances may dissolve in polar solvents as the polar solute molecules and the polar solvent molecules can interact with each other and attract each other. They can form intermolecular forces between each other causing the solute to dissolve

97
Q

What is the solubility of ionic substances?

A

Soluble in polar solvents

98
Q

Explain why ionic substances are soluble in polar solvents:

A

The polar solvent molecules interact with and attract the charged ions. This causes the ionic lattice to break down. The solvent molecules then surround the ions causing the solute to dissolve

99
Q

What substance can dissolve both polar and non-polar substances and why?

A

Ethanol. Ethanol contains both polar (-OH) and nonpolar (Carbon chain) in their structure and can therefore dissolve both polar and non-polar substances

100
Q

How do molecules containing hydrophobic and hydrophilic parts work?

A

Some biological molecules contain hydrophobic and hydrophilic parts. the hydrophilic part will be polar and can interact with the water, whereas the hydrophobic part is non poalr comprised of a carbon chain

101
Q

Are most simple molecualr substances polar or non-polar?

A

non-polar

A Level Chemistry (31 decks)

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