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Chemistry > Chapter 6 Shapes of Molecules and Intermolecular Forces > Flashcards

Chapter 6 Shapes of Molecules and Intermolecular Forces Flashcards

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1
Q

What is Electron-Pair repulsion theory

A

Pairs of electrons around a central atom repel each other, so they move as far apart as possible to minimise this repulsion. This means molecules containing 3 or more atoms have shapes

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2
Q

What is the molecular shape and bonding angles for:

2 pairs of electrons (or 2 regions of electron density) repelling around a central atom

A

Linear
180°

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3
Q

When working out shapes, how many regions of electron density do double and triple bonds count as

A

1

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4
Q

What is the molecular shape and bonding angles for:

3 pairs of electrons around a central atom

A

Trigonal Planar
120°

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5
Q

What is the molecular shape and bonding angles for:

4 bonding pairs around a central atom with no lone pairs

A

Tetrahedral
109.5°

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6
Q

What is the molecular shape and bonding angles for:

4 pairs of electrons around a central atom with 1 lone pair

A

Pyramidal
107°

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7
Q

What is the molecular shape and bonding angles for:

4 pairs of electrons around a central atom with 2 lone pairs

A

Non Linear
180°

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8
Q

What is the molecular shape and bonding angles for:

5 pairs of electrons around a central atom

A

Trional Bipyramid
90°
120°

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9
Q

What is the molecular shape and bonding angles for:

6 pairs of electrons around a central atom

A

Octahedral
90°

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10
Q

By how many degrees is the bond angle reduced for each lone pair

A

2.5°

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11
Q

Define Electronegativity

A

Electronegativity is a measure of the ability of an atom in a molecule to attract a pair of electrons in a covalent bond to itself

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12
Q

Finish this sentence:

The higher the electronegativity value…

A

The more strongly the bonding electrons are attracted to that atom

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13
Q

Going up a group does electronegativity increase or decrease

A

Electronegativity increases up a group

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14
Q

Across a period, does electronegativity increase or decrease

A

Electronegativity increases across a period (up to group 7)

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15
Q

What is the most electronegative atom

A

Fluorine

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16
Q

What group on the periodic table has the least electronegative atoms

A

1

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17
Q

Which elements have the most electronegative atoms

A
  • Nitrogen
  • Oxygen
  • Fluorine
  • Chlorine
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18
Q

What happens to the pull on bonding electrons as nuclear charge increases

A

It increases

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19
Q

Why does the pull on bonding electrons decrese further from the nucleus

A

Pull falls rapidly as electrons become further from the nucleus and the number of shielding inner shells increases

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20
Q

What is a non-polar bond

A

In a non-polar bond, the bonded electron pair is shared equally between the bonded atoms

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21
Q

When is a bond non polar

A
  • The bonded atoms are the same
  • The bonded atoms have the same, or similar, electronegativity values.
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22
Q

What is a polar bond

A

In a polar bond, the bonded electron pair is shared unequally between the bonded atoms

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23
Q

When is a bond polar

A

A bond will be polar when the bonded atoms are different and have different electronegativity values, resulting in a polar covalent bond

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24
Q

Why is H–F a polar bond

A

In H-F, F has a higher electronegativity than H so the bonding electrons are more attracted to the F end of the bond. The bond is polarised

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25
Q

What is the separation of opposite charges called

A

A dipole

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26
Q

Finish this sentence:

The bigger the difference in electronegativity between the bonding atoms…

A

The more polar the bond and the greater the ionic character

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27
Q

What difference in electronegativity will result in the bond being ionic instead of Polar Covalent

(Give a value)

A

1.8+

28
Q

How can a molecule with polar bonds be non-polar

A

The molecule is symmetrical and the dipoles cancel out

29
Q

What has to be true for a molecule with polar bonds to be polar

A

The molecule has to be unsymmetrical so the dipoles do not cancel out

30
Q

Can polar molecules (eg water) dissolve ions

A

Yes, polar molecules (particularly water) can form strong enough attractions to ions to permit dissolving in some cases

31
Q

When is dissolving likely

A

If the energy needed to break the lattice is similar to that released, it will dissolve

32
Q

What is hydration

A

When attractions form between water moleculesand ions releasing energy

33
Q

Are Nitrates soluble

A

Yes, all nitrates are soluble

34
Q

Are Chlorides soluble

A

Yes, all chlorides are soluble, except for AgCl and PbCl₂

35
Q

Are Sulfates soluble

A

Yes, most sulfates are soluble, except for BaSO₄ PbSO₄ and SrSO₄

36
Q

Are Carbonates soluble

A

No, all carbonates are insoluble, except (NH₄)₂CO₃ and those of the Group 1 elements

37
Q

Are sodium, potassium, and ammonium salts soluble

A

Yes all sodium, potassium and ammonium salts are soluble

38
Q

What do the solubility rules depend on

A

Amount of solute and/or solvent and temperature

39
Q

Define intermolecular forces

A

Weak interactions between dipoles of different molecules

40
Q

Which atoms or molecules London forces (induced dipole dipole forces)

A

All atoms and molecules

41
Q

Explain London Forces

A
  • Electrons in an atom or molecule are constantly moving
  • A temporary uneven distribution of electrons can occur for an instant
  • This creates an instantaneous dipole
  • The instantaneous dipole induces a dipole in its neighbours
  • leading to an attraction
  • Each dipole can induce many others
42
Q

How long do london forces last

A

Induced dipoles are only temporary. In the next instant of time, the induced dipoles may disappear, only for the whole process to take place amongst other molecules

43
Q

Which intermolecular forces do non-polar molecules have

A

London Forces Only

44
Q

What happens to the london force as the molecule gets larger

A

Number of electrons increases so size of the induced dipoles also get larger, giving stronger forces

45
Q

Why do the melting and boiling points increase as the strength of the London force increases

A

As the strength of the London forces increases, more energy is needed to overcome them, so the melting and boiling point increases

46
Q

Define Permanent Dipole-Dipole interactions

A

Dipole-Dipole interactionshese are formed between molecules which are polar (molecules with both polar bonds and a shape that does not allow the dipoles to cancel each other out)

47
Q

How long do permanent dipole-dipole interactions last

A

They are permanent and exert stronger forces. Opposites attract

48
Q

What are simple molecular substances

A

Simple molecular substances are made up of discrete molecules containing a specific number of atoms and having a consistent molecular formula e.g CH₄, N₂

49
Q

What is the regular structure of solid simple molecular substances called

A

Simple Molecular Lattice

50
Q

Are the intermolecular forces holding simple molecular substances together strong or weak

A

Weak

51
Q

What are the properties of simple molecular substances

Summarised

A
  • Low melting and boiling points
  • Non-polar molecules tend to dissolve in non-polar solvents, whereas polar molecules tend to dissolve in polar solvents
  • Do not conduct electricity
52
Q

Why do simple molecular substances have low melting and boiling points

A

Intermolecular forces usually do not need a lot of energy to overcome so most simple molecular substances have low melting and boiling temperatures with many being gases or liquids at room temperature

53
Q

Describe the solubility of simple molecular substances

A

Like dissolves like
Non-polar molecules tend to dissolve in non-polar solvents.
Polar molecules tend to dissolve in polar solvents

54
Q

Why can’t simple molecular substances conduct electricity

A

There are no mobile charged particles

55
Q

What is Hydrogen bonding

A

A special type of permanent dipole-dipole interaction containing:
* An electronegative atom with a lone pair of electrons, e.g. oxygen, nitrogen, or fluorine
* A hydrogen atom attached to an electronegative atom, e.g. H-O, H-N, or H-F

56
Q

What is the strongest type of intermolecular force

A

Hydrogen Bonding

57
Q

What are the anomalous properties of water

A
  • Solid (ice) is less dense than liquid (water)
  • Water has a relatively high melting point and boiling point
  • Water has a relatively high surface tension
58
Q

Why is ice less dense than water

A
  • Hydrogen bonds hold water molecules apart in an open lattice structure
  • The water molecules in ice are further apart than in liquid water
  • Solid ice is less dense than liquid water and floats
59
Q

How many hydrogen bonds can each water molecule form

A

4

60
Q

The hydrogen bonds extend outwards, holding water molecules slightly apart in an open tetrahedral arrangement

A

The hydrogen bonds extend outwards, holding water molecules slightly apart in an open tetrahedral arrangement

61
Q

What is roughly the bond angle for each Hydrogen bond

A

180⁰

62
Q

DO General trend – boiling point increases. As molecule size (and number of electrons) increases, the strength of the London forces between molecules increases and more energy is needed to overcome them.

A
63
Q

FINISH H2O has the highest boiling point as it can form more extensive hydrogen bonding between its molecules

A
64
Q

FINISH H2O, NH3, and HF do not follow the trend. They have stronger hydrogen bonds between their molecules which require more energy to break

A
65
Q

FINISH urface tension is caused by molecules on the surface experiencing unbalanced hydrogen bonding forces pulling them in.
Molecules in the bulk experience balanced forces in every direction.

A
66
Q

FINISH The double helix structure of DNA is held together by hydrogen bonds between base pairs.

A

Chemistry (16 decks)

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