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A-Level Chemistry > Bonding > Flashcards

Bonding Flashcards

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1
Q

what is an ionic bond?

A

electrostatic attraction between positive and negative ions

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2
Q

what is a giant ionic lattice?

A
  • a repeating pattern of oppositely charged ions
  • every ion in the lattice is surrounded by oppositely charged ions + are held together by strong electrostatic forces of attraction
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3
Q

list 4 properties of ionic compounds

A
  • very high mp/bp
  • tend to be soluble in polar solvents (e.g water)
  • do not conduct electricity when solid
  • can conduct electricity when molten/in solution
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4
Q

why do ionic compounds have a very high mp/bp?

A

takes a lot of energy to overcome the strong electrostatic forces of attraction

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5
Q

why do ionic compounds tend to be soluble in polar solvents?

A

solvent molecules surround ions, which can overcome electrostatic attraction between ions

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6
Q

how does the charge on ions impact the solubility of an ionic compound? why?

A

increase charge = decrease solubility as the solvent molecules cannot overcome the electrostatic forces of attraction

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7
Q

why don’t ionic compounds conduct electricity when solid but can when molten/in solution?

A

in a solid, ions are locked in place by the electrostatic forces of attraction and so can’t carry charge. when molten/in solution the lattice breaks down and the ions are free to move so they can carry charge

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8
Q

what affects the strength of an ionic bond?

A

ionic charge and ionic radius

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9
Q

how does ionic charge affect the strength of an ionic bond?

A

the higher the charge, the stronger
the bond

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10
Q

what is ionic radius and how does it affect the strength of an ionic bond?

A

ionic radius - the size of the ion
- the larger the radius, the weaker the bond because the charge is spread over a larger surface area

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11
Q

what is a covalent bond?

A

a shared pair of electrons

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12
Q

what holds together a covalent bond?

A

electrostatic attraction between nucleus of each atom + the shared electrons

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13
Q

what is a dative/coordinate bond?

A
  • a covalent bond where both of the electrons in the bond come from just one of the atoms involved
  • donating atom donates a non-bonding pair of electrons (lone pair)
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14
Q

how can a dative/coordinate bond be represented?

A

represented by an arrow coming from the atom that donated the electrons

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15
Q

what are the conditions required for a dative/coordinate bond?

A

acceptor must be electron deficient (must be available orbitals for the electrons to occupy)

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16
Q

how does a dative bond compare to a regular covalent bond?

A

a dative bond is the same as a regular covalent bond.
has the same length and average bond enthalpy (strength of bond)

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17
Q

when representing 3D shapes of molecules, what does a solid line represent?

A

the bond lies on the plane of the page

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18
Q

when representing 3D shapes of molecules, what does a solid wedge represent?

A

the bond is coming out of the plane of the page

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19
Q

when representing 3D shapes of molecules, what does a dotted wedge represent?

A

the bond is projecting back behind the plane of the page

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20
Q

what does electron repulsion theory state?

A
  • shape of a molecule determined by the number of areas of electron density surrounding the central atom
  • based on the fact that pairs of electrons repel all of the other electron pairs
  • the electron pairs now move as far apart as possible to minimise the repulsion
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21
Q

list the general steps to work out the shape of a molecule

A
  1. draw dot + cross diagram
  2. work out number of areas of electron density (1 lone pair, 1 single bond, 1 double bond etc are all ONE area of electron density)
  3. use this to choose the base geometry of the molecule
  4. if lone pairs present, assign actual shape based on number of lone pairs
  5. use shape to assign an appropriate bond angle
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22
Q

why does the angle of a molecule change when there are lone pairs present?

A
  • lone pairs closer to central atom + take up more space than a bonded pair of electrons
  • therefore lone pairs repel more strongly than bonding pairs
  • decreases the other bond angles by ~2.5°
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23
Q

shape and angle(s) of carbon dioxide (CO2) molecule?

A

linear, 180

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24
Q

shape and angle(s) of boron trichloride (BCl3) molecule?

A

trigonal planar, 120

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25
Q

shape and angle(s) of methane (CH4) molecule?

A

tetrahedral, 109.5

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26
Q

shape and angle(s) of ammonia (NH3) molecule?

A

trigonal pyramidal, 107

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27
Q

shape and angle(s) of water (H2O) molecule?

A

bent, 104.5

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28
Q

shape and angle(s) of sulphur hexafluoride (SF6) molecule?

A

octahedral, 90

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29
Q

shape and angle(s) of tetrachloromethane (CCl4) molecule?

A

tetrahedral, 109.5

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30
Q

shape and angle(s) of nitrate ion (NO3-) molecule?

A

trigonal planar, 120

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31
Q

shape and angle(s) of sulphurhexachloride (SCl6) molecule?

A

octahedral, 90

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32
Q

shape and angle(s) of hydrogen sulfide (H2S) molecule?

A

bent, 104.5

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33
Q

shape and angle(s) of nitrogen trichloride (NCl3) molecule?

A

trigonal pyramidal, 107

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34
Q

shape and angle(s) of molecule with 2 areas of electrons density (2 b.p, 0 l.p)

A

linear, 180

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35
Q

shape and angle(s) of molecule with 3 areas of electrons density (3 b.p, 0 l.p)

A

trigonal planar, 120

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36
Q

shape and angle(s) of molecule with 3 areas of electrons density (2 b.p, 1 l.p)

A

bent, 119

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37
Q

shape and angle(s) of molecule with 4 areas of electrons density (4 b.p, 0 l.p)

A

tetrahedral, 109.5

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38
Q

shape and angle(s) of molecule with 4 areas of electrons density (3 b.p, 1 l.p)

A

trigonal pyramid, 107

39
Q

shape and angle(s) of molecule with 4 areas of electrons density (2 b.p, 2 l.p)

A

bent, 104.5

40
Q

shape and angle(s) of molecule with 5 areas of electrons density (5 b.p, 0 l.p)

A

trigonal bipyramidal, 90 and 120

41
Q

shape and angle(s) of molecule with 5 areas of electrons density (4 b.p, 1 l.p)

A

sawhorse - <90 and <120

42
Q

shape and angle(s) of molecule with 5 areas of electrons density (3 b.p, 2 l.p)

A

T-shaped, 90

43
Q

shape and angle(s) of molecule with 5 areas of electrons density (2 b.p, 3 l.p)

A

linear, 180

44
Q

shape and angle(s) of molecule with 6 areas of electrons density (6 b.p, 0 l.p)

A

octahedral, 90

45
Q

shape and angle(s) of molecule with 6 areas of electrons density (5 b.p, 1 l.p)

A

square pyramidal, 90

46
Q

shape and angle(s) of molecule with 6 areas of electrons density (4 b.p, 2 l.p)

A

square planar, 90

47
Q

shape and angle(s) of molecule with 6 areas of electrons density (3 b.p, 3 l.p)

A

T-shaped, 90

48
Q

shape and angle(s) of molecule with 6 areas of electrons density (2 b.p, 4 l.p)

A

linear 180

49
Q

what is electronegativity?

A

the power of an atom to attract the electron density of a covalent bond towards itself

50
Q

on the Pauling scale, what is the largest electronegativity? which element is this?

A

4.0, Fluorine

51
Q

why don’t noble gases have a value on the Pauling scale?

A

they usually don’t form covalent bonds

52
Q

list the 3 factors that affect electronegativity

A
  1. nuclear charge of atom
  2. atomic radius
  3. shielding of nuclear charge by electrons in inner shells
53
Q

how does the nuclear charge of an atom affect electronegativity?

A

increased positive charges increases the attraction between the nucleus and the pair of electrons in the covalent bond

54
Q

how does atomic radius affect electronegativity?

A
  • smaller atomic radius = closer outer electrons will be to nucleus of atom = more attraction
  • increase atomic radius = decrease electronegativity
55
Q

how does the shielding of nuclear charge by electrons in inner shells affect electronegativity?

A

increase shielding = greater number of inner shells = lower electronegativity as there is a lesser attraction between outer electrons and the nucleus

56
Q

why does electronegativity decrease down a group?

A

more shielding + greater atomic radius as you move down group so lesser attraction between outer electrons + nucleus

57
Q

why does electronegativity increase across a period?

A

the number of protons increases across a period, so the nuclear charge increases, which increases the attraction between the nucleus and the pair of electrons in the covalent bond

58
Q

what does polarity mean? what makes a molecule non-polar?

A
  • the unequal sharing of electrons in a bond
  • if the electron sharing is equal, the bond is described as non-polar
59
Q

what is a dipole?

A

molecule/bond that has a separation of charges

60
Q

what is a pure covalent bond?

A

a covalent bond that is non-polar/has no dipole

61
Q

if a molecule has several polar bonds but is symmetrical, does it have an overall polarity? why/why not?

A

because the molecules are symmetrical, there is no overall polarity, and the electron cloud is shared evenly over the whole molecule

62
Q

what is a simple molecular substance?

A

a substance that consists of relatively small molecules with a fixed number of atoms

63
Q

what is an intermolecular force?

A
  • force of attraction that acts between molecules
  • much weaker than covalent bonds + are easily broken
64
Q

list the 3 types of intermolecular forces

A

1) induced dipole-dipole interactions
2) permanent dipole-dipole interactions
3)hydrogen bonding

65
Q

why do simple molecular substances have relatively low bps?

A

when heated the molecules move faster (inc ke). at certain temps, the intermolecular forces break allowing the molecules to move away from each other

66
Q

what is an induced dipole-dipole interaction?

A
  • caused by random electron movement
  • NOT permanent
  • weak and easily broken
  • due to random electron movement, a dipole may form instantaneously as one side of atom has a slightly more positive/negative charge. because the electrons move to one side, the electrons in surrounding atoms are repelled and now they too have a dipole - this is an induced dipole
    all of these dipoles experience a force of attraction
67
Q

which molecules experience induced dipole-dipole interactions?

A

all

68
Q

what does the strength of an induced dipole-dipole interaction depend on?

A
  • size/number of electrons
  • bigger atom = more electrons = stronger induced dipole-dipole interactions
69
Q

why do branched alkanes have a lower bp than their straight counterparts?

A

there is less surface area of contact between the molecules, lessening the induced dipole-dipole interactions

70
Q

Explain why pentane (C5H12) has a higher boiling point than methane (CH4)

A
  • stronger intermolecular forces = higher the boiling point
  • pentane has stronger induced dipole-dipole interactions because it is a larger molecule than methane + has a larger surface area which allows more contact between molecules
71
Q

what is a permanent dipole-dipole interaction?

A
  • the attraction between partially positive and partially negative ends of the molecule
  • exists only in molecules with a positive dipole
  • stronger than induced dipole-dipole interactions
72
Q

why won’t an atom with permanent dipole-dipole interactions always have a higher bp than an atom without?

A
  • permanent dipole-dipole interactions not only force acting as all forces experience induced dipole-dipole interactions
  • size of the induced dipole-dipole interactions depends on the number of electrons present
  • count electrons to determine strength
73
Q

HCl and F2 have the same number of electrons and the same shape. Which will have the higher boiling point F2 or HCl? Why?

A

F2 has only induced dipole-dipole interactions. HCl has induced dipole-dipole and permanent dipole-dipole interactions. Extra energy is needed to break the permanent dipole-dipole interactions in HCl so it has a higher boiling point.

74
Q

what is hydrogen bonding?

A
  • a hydrogen atom bonded to a strongly electronegative element
  • the strongest type of intermolecular force
  • type of permanent dipole-dipole when there is a O-H, N-H or H-F bond
75
Q

when does hydrogen bonding occur?

A

when there is a:
O-H
N-H
H-F bond

76
Q

why does water have a high mp/bp?

A

takes a lot of energy to break the hydrogen bond

77
Q

why is does ice float on water?

A
  • molecules move randomly in liquid water - sometimes they’re closer together + sometimes further apart
  • hydrogen bonds constantly formed + broken
  • molecules in water move more slowly as it cools down
  • as they reach freezing point (0°c), the water molecules arrange themselves in an ordered structure (ice) which is stabilised by network of hydrogen bonds
  • in ice, the water molecules are further apart than in liquid water
  • therefore, ice is less dense than water causing it to float
78
Q

why is it important that ice is less dense than water?

A

ice insulates the water below which prevents the water from freezing completely and prevents the organisms that live in the water from dying

79
Q

why do simple molecular substances have a low mp/bp

A
  • weak intermolecular forces
  • increase number of electrins = increase bp
80
Q

describe the solubility of non-polar simple molecular substances

A
  • dissolve very well in non-polar solvents
  • when dissolved in a non-polar solvent, the solvent molecules form induced dipole-dipole interactions with the solvent molecules
  • generally insoluble in polar solvents
81
Q

describe the solubility of polar simple molecular substances

A
  • dissolve in polar solvents
  • usually insoluble in non-polar solvents
82
Q

describe the electrical conductivity of simple molecular substances

A

simple molecular substances do not contain mobile charged particles and so can’t conduct electricity

83
Q

arrangement of particles in a solid? evidence?

A
  • regular
  • crystals have straight edges
84
Q

arrangement of particles in a liquid? evidence?

A
  • random
  • changes shape to fill the bottom of a container
85
Q

arrangement of particles in a gas? evidence?

A
  • random
  • will fill its container
86
Q

spacing in a solid? evidence?

A
  • close
  • not easily compressed
87
Q

spacing in a liquid? evidence?

A
  • close
  • not easily compressed
88
Q

spacing in a gas? evidence?

A
  • far apart
  • easily compressed
89
Q

movement in a solid? evidence?

A
  • just vibrate
  • slow diffusion, expand on heating
90
Q

movement in a liquid? evidence?

A
  • rapid “jostling”
  • slow diffusion, evaporation
91
Q

movement in a gas? evidence?

A
  • rapid movement
  • exertion of pressure
92
Q

describe the bonding, structure and properties of diamond

A
  • bonding: covalent
  • structure: giant covalent
  • properties:
  • high mp and very strong : held together by very strong covalent bonds which require lots of energy to break.
  • does not conduct electricity: contains no charged particles (electrons or
    ions), to carry current
  • uses: jewellery, blades for cutting hard objects
93
Q

describe the bonding, structure and properties of graphite

A

*bonding: covalent
*structure: giant covalent
*properties:
- high mp but soft : held together by very strong covalent bond, but weak intermolecular forces between the layers
- conducts electricity: delocalised electrons are free to move and transfer the electrical current
*uses: electrodes and pencils

94
Q

how does a hydrogen bond arise?

A
  • large difference in electronegativity between H and other atom
  • dipole created where H is slightly positive and other atom is slightly negative
  • bond formed between slightly positive hydrogen and the lone pair on other atom

A-Level Chemistry (23 decks)

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