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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

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
Define INDUCED DIPOLE-DIPOLE INTERACTION (London forces)
Attractive forces between induced dipoles in different molecules
26
Define LONE PAIR
An outer shell pair of electrons that is not involved in chemical bonding
27
Define PAULING ELECTRONEGATIVITY VALUE
A value assigned as a measure of the relative attraction of a bonded atom for the pair of electrons in a covalent bond
28
Define NON-POLAR
With no charge separation across a bond or in a molecule
29
Define PERMANENT DIPOLE
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
30
Define PERMANENT DIPOLE-DIPOLE INTERACTION
An attractive force between permanent dipoles in neighbouring polar molecules
31
Define POLAR
With slightly positive and slightly negative charges at different ends of the molecule
32
Define POLAR COVALENT BOND
A bond with a permanent dipole, having slightly positive and negative partial charges on the bonded atoms
33
Define POLAR MOLECULE
A molecule with an overall dipole, having taken into account any dipoles across bonds and the shape of the molecule
34
Define SHIELDING EFFECT
The repulsion between electrons in different inner shells. Shielding reduces the net attractive force between the positive nucleus and the outer shell electrons
35
What is the electronegativity trend?
- increase up and to the right of the periodic table - noble gases not included - fluorine the most electronegative
36
Why does Electronegativity increase across a period?
- charge on the nucleus increases - number of protons increase - increased attraction for outer electrons - bonding pair of electrons attracted more strongly
37
Why does electronegativity increase up a group?
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
What is electron density?
The probability of finding electrons at a particular position in space.
39
Explain non-polar bonds.
- 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
Explain polar bonds.
- the two bonding atoms are different - the atom with greater attraction is more electronegative - the bond is polarised
41
Explain dipoles and permanent dipoles
Dipoles- differing attraction for pair of electrons allows small charge difference between atoms Permanent dipoles- slight charge difference always present
42
How is electronegativity measured?
On the Pauling scale.
43
State the features of a non-symmetrical molecule.
- difference in charge across molecule - there is an overall dipole - molecule is polar - e.g. H2O
44
State the features of a symmetrical molecule.
- permanent dipole charges cancel out - all atoms attached to centre identical - no difference in charge - molecule is non-polar
45
Explain the spectrum of bonds.
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
What are the 3 types of intermolecular forces?
- London forces (induced dipole-dipole forces) - permanent dipole-dipole forces - hydrogen bonds
47
What are intermolecular interactions?
- 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
State the intermolecular forces in order of strength (strongest to weakest)
1. Hydrogen bonding 2. Permanent dipole-dipole interactions 3. London forces
49
Explain how London forces are created.
- 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
What factor effects the size and strength of London forces and why?
- number of electrons - more electrons increases chance of instantaneous dipole forming / increases strength of the force so melting and boiling point increases
51
Where are London forces found?
- in all molecules (sometimes along with other forces)
52
Where are permanent dipole-dipole interactions found?
- between molecules with permanent dipoles
53
Explain how permanent dipole-dipole interactions are formed?
- 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
What are the Van def Waals’ forces
- London forces | - dipole-dipole interactions
55
What are hydrogen bonds?
- strong permanent dipole-dipole forces of attraction
56
Where are hydrogen bonds found?
- between a hydrogen atom and the lone pair from either oxygen, nitrogen, or sulfur (more electronegative element)
57
What bonds are in water molecules?
- hydrogen bonds
58
What are 3 properties that water has because of its hydrogen bonding?
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
Why do noble gases increase in boiling point down their group?
- weak Van der waals’ attractions exist - attractions increase as number of electrons increase - boiling point increases
60
Why do H2O, HF, NH3 have higher boiling points than expected?
- all have hydrogen bonding - hydrogen bonds stronger than other intermolecular forces - boiling point increases
61
Explain why simple molecular substances have low melting/boiling points.
- have weak intermolecular forces (easy to break)
62
Explain why NON-POLAR simple molecular substances are soluble in NON-POLAR solvents.
- intermolecular forces form between molecule and solvent - weakens intermolecular forces in simple lattice - so they break and compound dissolves
63
Explain why NON-POLAR simple molecular substances are insoluble in POLAR solvents.
- little interaction between molecules in lattice and solvent - intermolecular bonding in solvent is too strong to break
64
Explain why the solubility of POLAR simple molecular substances is hard to predict.
- 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
Explain why simple molecular substances do not conduct electricity.
- no mobile charged particles
66
State the features of long, straight chained hydrocarbons.
- large SA - more induced dipole-dipole interactions - more energy needed to overcome - higher boiling point
67
State the features of branched hydrocarbons.
- can’t pack closely together - weakens induced dipole-dipole forces between chains - lowers boiling point
68
Why do branched hydrocarbons decrease the strength of London forces?
- 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
How do you identify a permanent dipole?
A large difference in electronegativity between the atoms
70
Explain the trend of boiling points of hydrogen halides?
- 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
State the 4 types of bonds and their bond enthalpy.
London forces —> 1 - 10 Permanent dipole-dipole —> 3 - 25 Hydrogen bonding —> 10 - 40 Single covalent —> 150 - 500
72
State the type of bond and their electronegativity difference.
Covalent —> 0 Polar covalent —> 0 to 1.8 Ionic —> greater than 8
73
Order types of pairs from decreasing to increasing repulsion.
- bonded pair / bonded pair - bonded pair / lone pair - lone pair / lone pair
74
Explain the shape and bonding in water.
- 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
Explain the shape and bonding in ice.
- 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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