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CHEMISTRY (OCR A-LEVEL) > 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
  • the number of electron pairs surrounding the central atom determine the shape of the molecule
  • electron pairs repel as far as possible so that they are arranged equally
  • the arrangement of electron pairs minimise repulsion and hold the shape of the molecule
  • lone pairs repel other pairs more than bonded pairs
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2
Q

What factors determine the shape of a compound or ion?

A
  • the number of electron pairs around the central atom

- the nature of the pairs (bonding / lone)

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3
Q

What are the features of bonding pairs?

A
  • they repel equally

- involved in bonding

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4
Q

What are the features of lone pairs?

A
  • more compact so repel other pairs more than bonding pairs
  • not involved in bonding
  • for each lone pair, bond angle decreases by 2.5
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5
Q

Explain what the different wedges represent when drawing molecules.

A
  • normal : bond in plane of paper
  • dotted : bond going into paper
  • bold : bond coming out of paper
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6
Q

What is the octet rule?

A
  • 8 electrons in pairs maximum around one atom
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7
Q

Explain shrinking octet.

A
  • not enough electrons to pair and form an octet so….
  • unpaired electrons pair up
  • octet is not achieved
  • e.g. boron trifluoride
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8
Q

Explain expanding octet.

A
  • bonding atoms have 8+ electrons in outer shell
  • occurs period 3 down only in groups 15, 16, 17
  • e.g. SF6-sulfur hexafluoride
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9
Q

Describe a linear shaped molecule and give an example.

A
  • one or two bonding pair
  • bond angle 180
  • e.g. H2 or CO2
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10
Q

Describe a trigonal planar shaped molecule and give an example.

A
  • 3 bonding pairs
  • bond angle 120
  • e.g. BF3
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11
Q

Describe a pyramidal shaped molecule and give an example.

A
  • 3 bonding pairs
  • 1 lone pair
  • bond angle 107
  • e.g. NH3
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12
Q

Describe a non-linear shaped molecule and give an example.

A
  • 2 bonding pairs
  • 2 lone pairs
  • bond angle 104.5
  • e.g. H2O
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13
Q

Describe a tetrahedral shaped molecule and give an example.

A
  • 4 bonding pairs
  • bond angle 109.5
  • e.g. CH4
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14
Q

Describe a trigonal bypyramidal shaped molecule and give an example.

A
  • 5 bonding pairs
  • bond angle 90 (on normal wedges) and 120 between (solid and dotted wedges)
  • e.g. PCl5
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15
Q

Describe an octahedral shaped molecule and give an example.

A
  • 6 bonding pairs
  • bond angle 90
  • e.g. SF6
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16
Q

How do you work out the shape of a molecule (4 steps) ?

A
  1. Write number of electrons on outer shell of central atom
  2. Write the number of atoms bonded to the central atom
  3. Work out number of lone pairs —> 1/2 (step 1- step 2)
  4. Find shape and bond angle with number of bonding pairs and lone pairs of electrons
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17
Q

What are the 5 steps used to explain the shape of a molecule?

A
  1. State the number of bonding and lone pairs of electrons
  2. State: ‘electron pairs repel and try to get as far apart as possible to a position of minimal repulsion’
  3. If there are NO lone pairs, state: ‘the bonded pairs repel equally’
  4. If there are lone pairs, state that: ‘the lone pairs repel more than bonded pairs’
  5. State the actual shape and bond angle.
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18
Q

Define BOND ANGLE

A

the angle between two bonds in a molecule

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19
Q

Define BONDED PAIR

A

A pair of electrons shared between two atoms to make a covalent bond

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20
Q

Define DIPOLE

A

A separation in electrical charge so that one end of a polar molecule has a slightly positive charge and the other end has a slightly negative charge

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21
Q

Define ELECTRONEGATIVITY

A

The ability of an atom to attract electrons in a covalent bond

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22
Q

Define COVALENT BOND

A

the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

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23
Q

Define INTERMOLECULAR FORCE

A

An attractive force between molecules (e.g. London forces, permanent dipole-dipole interactions, hydrogen bonding)

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24
Q

Define HYDROGEN BOND

A

A strong dipole-dipole attraction between an electron- deficient hydrogen atom on one molecule and a lone pair of electrons (N,O or F) on a different molecule

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25
Q

Define INDUCED DIPOLE-DIPOLE INTERACTION (London forces)

A

Attractive forces between induced dipoles in different molecules

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26
Q

Define LONE PAIR

A

An outer shell pair of electrons that is not involved in chemical bonding

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27
Q

Define PAULING ELECTRONEGATIVITY VALUE

A

A value assigned as a measure of the relative attraction of a bonded atom for the pair of electrons in a covalent bond

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28
Q

Define NON-POLAR

A

With no charge separation across a bond or in a molecule

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29
Q

Define PERMANENT DIPOLE

A

A small charge difference that does not change across a bond with partial charges ( slightly + and slightly -) on the bonded atoms —> this because bonding atoms have different electronegativities

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30
Q

Define PERMANENT DIPOLE-DIPOLE INTERACTION

A

An attractive force between permanent dipoles in neighbouring polar molecules

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31
Q

Define POLAR

A

With slightly positive and slightly negative charges at different ends of the molecule

32
Q

Define POLAR COVALENT BOND

A

A bond with a permanent dipole, having slightly positive and negative partial charges on the bonded atoms

33
Q

Define POLAR MOLECULE

A

A molecule with an overall dipole, having taken into account any dipoles across bonds and the shape of the molecule

34
Q

Define SHIELDING EFFECT

A

The repulsion between electrons in different inner shells. Shielding reduces the net attractive force between the positive nucleus and the outer shell electrons

35
Q

What is the electronegativity trend?

A
  • increase up and to the right of the periodic table
  • noble gases not included
  • fluorine the most electronegative
36
Q

Why does Electronegativity increase across a period?

A
  • charge on the nucleus increases
  • number of protons increase
  • increased attraction for outer electrons
  • bonding pair of electrons attracted more strongly
37
Q

Why does electronegativity increase up a group?

A

As you go down group:

  • bonding pairs held increasingly further away from nucleus
  • number of shells increase
  • distance of outer electrons from nucleus increase
  • shielding of inner shell electrons increase
  • bonding pair attracted less strongly
38
Q

What is electron density?

A

The probability of finding electrons at a particular position in space.

39
Q

Explain non-polar bonds.

A
  • bonding atoms are identical
  • the attraction for the shared pair of electrons is equal
  • the electrons are equally distributed
  • the bond is perfectly covalent
  • atoms have the same or similar electronegativity
40
Q

Explain polar bonds.

A
  • the two bonding atoms are different
  • the atom with greater attraction is more electronegative
  • the bond is polarised
41
Q

Explain dipoles and permanent dipoles

A

Dipoles- differing attraction for pair of electrons allows small charge difference between atoms
Permanent dipoles- slight charge difference always present

42
Q

How is electronegativity measured?

A

On the Pauling scale.

43
Q

State the features of a non-symmetrical molecule.

A
  • difference in charge across molecule
  • there is an overall dipole
  • molecule is polar
  • e.g. H2O
44
Q

State the features of a symmetrical molecule.

A
  • permanent dipole charges cancel out
  • all atoms attached to centre identical
  • no difference in charge
  • molecule is non-polar
45
Q

Explain the spectrum of bonds.

A
  1. ionic : difference in electronegativity very big/ one atom takes the electrons from the other atom
  2. polar-covalent : difference in electronegativity is quite small/ atoms share electrons unequally
  3. covalent : no difference in electronegativity/ molecule symmetrical
46
Q

What are the 3 types of intermolecular forces?

A
  • London forces (induced dipole-dipole forces)
  • permanent dipole-dipole forces
  • hydrogen bonds
47
Q

What are intermolecular interactions?

A
  • forces of attraction between molecules
  • non-bonded interactions
  • do not involve the transfer of electrons
  • the result of constant and random movement of electrons within shells of atoms
48
Q

State the intermolecular forces in order of strength (strongest to weakest)

A
  1. Hydrogen bonding
  2. Permanent dipole-dipole interactions
  3. London forces
49
Q

Explain how London forces are created.

A
  • electrons are constantly moving
  • in an instant, more electrons on one side of the molecule than the other
  • charged unequally distributed temporarily and causes an instantaneous dipole
  • instantaneous dipole induces neighbouring atoms and so forth to form induced dipoles
  • these dipoles are attracted to each by London forces
  • dipoles are constantly being formed and destroyed
50
Q

What factor effects the size and strength of London forces and why?

A
  • number of electrons
  • more electrons increases chance of instantaneous dipole forming / increases strength of the force so melting and boiling point increases
51
Q

Where are London forces found?

A
  • in all molecules (sometimes along with other forces)
52
Q

Where are permanent dipole-dipole interactions found?

A
  • between molecules with permanent dipoles
53
Q

Explain how permanent dipole-dipole interactions are formed?

A
  • if the molecules are correctly aligned, than they will attract each other and form interactions
  • BUT as molecules are constantly and randomly moving, molecules do not always align
54
Q

What are the Van def Waals’ forces

A
  • London forces

- dipole-dipole interactions

55
Q

What are hydrogen bonds?

A
  • strong permanent dipole-dipole forces of attraction
56
Q

Where are hydrogen bonds found?

A
  • between a hydrogen atom and the lone pair from either oxygen, nitrogen, or sulfur (more electronegative element)
57
Q

What bonds are in water molecules?

A
  • hydrogen bonds
58
Q

What are 3 properties that water has because of its hydrogen bonding?

A
  1. Ice less dense than water: H bonds in ice hold water molecules apart in open lattice structure of rings which makes it less dense
  2. High melting and boiling point: H bonds much stronger than other intermolecular forces so more energy needed to overcome them
  3. High surface tension: lattice strong and flexible- allows insects to walk on water
59
Q

Why do noble gases increase in boiling point down their group?

A
  • weak Van der waals’ attractions exist
  • attractions increase as number of electrons increase
  • boiling point increases
60
Q

Why do H2O, HF, NH3 have higher boiling points than expected?

A
  • all have hydrogen bonding
  • hydrogen bonds stronger than other intermolecular forces
  • boiling point increases
61
Q

Explain why simple molecular substances have low melting/boiling points.

A
  • have weak intermolecular forces (easy to break)
62
Q

Explain why NON-POLAR simple molecular substances are soluble in NON-POLAR solvents.

A
  • intermolecular forces form between molecule and solvent
  • weakens intermolecular forces in simple lattice
  • so they break and compound dissolves
63
Q

Explain why NON-POLAR simple molecular substances are insoluble in POLAR solvents.

A
  • little interaction between molecules in lattice and solvent
  • intermolecular bonding in solvent is too strong to break
64
Q

Explain why the solubility of POLAR simple molecular substances is hard to predict.

A
  • depends on dipole strength
  • may dissolve as polar solute and solvent molecules attract each other
  • some with part polar and part non-polar dissolve
65
Q

Explain why simple molecular substances do not conduct electricity.

A
  • no mobile charged particles
66
Q

State the features of long, straight chained hydrocarbons.

A
  • large SA
  • more induced dipole-dipole interactions
  • more energy needed to overcome
  • higher boiling point
67
Q

State the features of branched hydrocarbons.

A
  • can’t pack closely together
  • weakens induced dipole-dipole forces between chains
  • lowers boiling point
68
Q

Why do branched hydrocarbons decrease the strength of London forces?

A
  • strength of London forces depend on how close the particles get to each other
  • branching prevents molecules from getting near each other
  • reduces forces and boiling point
69
Q

How do you identify a permanent dipole?

A

A large difference in electronegativity between the atoms

70
Q

Explain the trend of boiling points of hydrogen halides?

A
  • strength of London forces decreases going down group
  • due to decreasing electronegativity
  • because of increased number of electrons
  • London forces increase
  • boiling point increases
71
Q

State the 4 types of bonds and their bond enthalpy.

A

London forces —> 1 - 10
Permanent dipole-dipole —> 3 - 25
Hydrogen bonding —> 10 - 40
Single covalent —> 150 - 500

72
Q

State the type of bond and their electronegativity difference.

A

Covalent —> 0
Polar covalent —> 0 to 1.8
Ionic —> greater than 8

73
Q

Order types of pairs from decreasing to increasing repulsion.

A
  • bonded pair / bonded pair
  • bonded pair / lone pair
  • lone pair / lone pair
74
Q

Explain the shape and bonding in water.

A
  • 2 bonding and 2 lone pairs
  • electron pairs repel as much as possible
  • lone pairs repel more than bonding pairs
  • bond angle is 104.5 (non-linear)
75
Q

Explain the shape and bonding in ice.

A
  • two lone pairs involved in hydrogen bonds
  • so all four electron pairs are bonding
  • shape tetrahedral with bond angle 109.5

CHEMISTRY (OCR A-LEVEL) (8 decks)

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