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A-level Chemistry > 3.1.3 Bonding > Flashcards

3.1.3 Bonding Flashcards

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1
Q

Define ionic bonding

A

Electrostatic force of attraction between oppositely charged ions formed by electron transfer

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2
Q

Ionic Bonding

Metal atoms ___ electrons to form ___ ions

A

Metal atoms lose electrons to form +ve ions

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3
Q

Ionic Bonding

Non-metal atoms ____ electrons to form ___ ions

A

Non-metal atoms gain electrons to form -ve ions

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4
Q

Simplest ions are…

A

single atoms which have lost/gained electrons to make full outer shell

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5
Q

What are compound ions?

A

Ions that are made up of groups of atoms with an overall charge

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6
Q

State the formula for a sulfate ion

A

SO42-

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7
Q

State the formula for a hydroxide ion

A

OH-

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8
Q

State the formula for a nitrate ion

A

NO3-

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9
Q

State the formula for a carbonate ion

A

CO32-

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10
Q

State the formula for an ammonium ion

A

NH4+

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11
Q

Name the structure of ionic crystals

A

Giant Ionic Lattice

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12
Q

Sodium nitrate contains Na+ (1+) and NO3- (1-) ions. State the fomula of the sodium nitrate.

A
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13
Q

Magnesium chloride contains Mg2+ (2+) and Cl- (1-) ions. State the fomula of the magnesium chloride.

A
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14
Q

Name 3 physical properties of ionic compounds

A
  1. Conduct electricity only when they’re molten or dissolved
  2. High melting points
  3. Tend to dissolve in water
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15
Q

Why can ions conduct electricity when they’re molten or dissolved?

A

∵ ions in liquid are free to move and carry a charge

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16
Q

Why can’t ions conduct electricity when they’re in a solid?

A

∵ ions are in fixed position by strong ionic bonds

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17
Q

Why do ionic compounds have high melting points?

A
  • Giant ionic lattices
  • Strong electrostatic forces of attraction between oppositely charged ions
  • Takes a lot of energy to overcome these forces
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18
Q

Why do ionic compounds tend to dissolve in water?

A
  • Water molecules are polar
    • Part of molecule has a small negative charge and other bits have small positive charges
  • Charged parts pull ions away from lattice = causing it dissolve
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19
Q

Ionic bonding is stronger and melting points are higher when ions are… (2x)

A

smaller and/ or have higher charges

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20
Q

When do molecules form and how are they held together?

A
  • Form when 2 or more atoms bond together
  • Held together by strong covalent bonds
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21
Q

What do single covalent bonds contain?

A

Shared pair of electrons

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22
Q

Describe covalent bonding

A
  1. Two atoms share electrons so they’ve both got full outer shells
  2. Both postive nuclei are attracted electrostatically to shared electrons
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23
Q

Multiple covalent bonds contain…

A

multiple shared pairs of electrons

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24
Q

Draw methane, represent the covalent bonds by drawing lines

A
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25
Q

Why can carbon form giant covalent structures?

A

∵ they can form 4 covalent bonds

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26
Q

Describe the structure of graphite

A
  • Carbon atoms are arranged in sheets of flat hexagons covalently bonded with 3 bonds each
  • 4th outer electron of each carbon atom is delocalised
  • Sheets of hexagons are boned together by weak van der Waal forces
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27
Q

Name 5 properties of graphite

A
  1. Low density
  2. Dry lubricant/slippy
  3. Electrical conductor
  4. Insoluble in any solvent
  5. Very high melting point
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28
Q

Explain why graphite has a low density

A

Layers are quite far apart compared to the length of covalent bonds

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29
Q

Explain why graphite is an electrical conductor

A

‘Delocalised’ electrons aren’t attached to any particular carbon atoms & free to move along sheets carrying a charge

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30
Q

Explain why graphite is a dry lubricant/slippy

A

Weak bonds between layers in graphite = easily broken ∴ sheets can slide over each other

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31
Q

Explain why graphite has a very high melting point

A

Covalent bonds are very strong and require lots of energy to break

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32
Q

Explain why graphite is insoluble in any solvent

A

Covalent bonds in sheets are too strong to break

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33
Q

Describe the structure of diamond

A
  • Each carbon atom is covalently bonding to 4 other carbon atoms (giant covalent structure)
  • Tetrahedral shape
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34
Q

Name 5 properties of diamond

A
  1. Very high melting point
  2. Extremely hard
  3. Good thermal conductor
  4. Can’t conduct electricity
  5. Won’t dissolve in any solvent
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35
Q

Why is diamond a good thermal conductor?

A

Vibrations travel easily through stiff lattice

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36
Q

Why can’t diamond conduct electricity?

A

Outer electrons held in localised bonds

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37
Q

Why do diamond gemstones sparkle a lot?

A

Its structure makes it refract light a lot

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38
Q

What is dative covalent bonding (or co-ordinate bonding)?

A

When shared pair of electrons in covalent bond come from only one of the bonding atoms

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39
Q

Name an example of dative covalent bonding & explain how it is an example of this bonding

A

Ammonium ion (NH4+)

Forms when nitrogen atom in ammonia molecule donates a pair electrons to proton (H+)

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40
Q

Illustrate dative covalent bonding in an ammonium ion (NH4+)

A
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41
Q

Define metallic bonding

A

Metallic bonding is the electrostatic force of attraction between the positive metal ions and the delocalised electrons

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42
Q

Metals elements exist as…

A

giant metallic lattice structures

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43
Q

Describe metallic bonding

A
  1. Outer shell electrons of metal are delocalised
    1. Electrons free to move
    2. Leaves positive metal ion
  2. Positive metal ions attracted to delocalised negative electrons
    1. Form lattice of closely packed positively ions in sea of delocalised electrons
    2. This is metallic bonding
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44
Q

Name 4 properties of metals

A
  • High melting points
  • Good thermal conductors
  • Good electrical conductors
  • Insoluble (expect in liquid metals)
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45
Q

Why do metals have high melting points?

A

Strong electrostatic attraction between positive metal ions and delocalised sea of electrons

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46
Q

Why are metals good thermal conductors?

A

Delocalised electrons can pass kinetic energy to each other

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47
Q

Why are metals good electrical conductors?

A

Delocalised electrons can move and carry current

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48
Q

Why are metals insoluble?

A

Strong metallic bonds

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49
Q

Name 3 factors that affect the strength of metallic bonding

A
  1. Number of protons/strength of nuclear attraction
  2. Number of delocalised electrons per atom
  3. Size of ion
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50
Q

Metallic Bonding

More protons = ….

A

stronger bond

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51
Q

Metallic Bonding

More delocalised electrons per atom = ….

A

stronger bonding

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52
Q

Metallic Bonding

Smaller the ion = …

A

stronger the bond

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53
Q

Explain why Mg has stronger metallic bonding than Na and a higher melting point

A
  1. In Mg: more electrons in outer shell that are released to sea of electrons
  2. Mg ion is smaller and has more than one proton
  3. ∴ stronger electrostatic attraction between positive metal ions and delocalised electrons = higher energy is needed to break bonds
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54
Q

Illustrate a giant ionic lattice of sodium chloride

A
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55
Q

Illustrate metallic bonding in magnesium

A
56
Q

What does the shape of a molecule depend on?

A

The number of pairs of electrons in outer shell of central atom

57
Q

Bonding pairs and lone pairs electrons exist as ___ ____

A

charge clouds

58
Q

What are charge clouds?

A

Area where you have really big chance of finding an electron pair

59
Q

Why does a pair of electrons in an outer shell of an atom sit as far apart from each other as possible?

A

Electrons = negatively charged ∴ repel each other

60
Q

Shape of charge cloud effects…

A

how much it repels other charge clouds

e.g. lone-pair charge clouds repel more than bonding-pair charge clouds

61
Q

Why are the bond angles between bonding pairs reduced when lone pairs of electrons are added?

A

∵ they’re pushed together by lone-pair repulsion

62
Q

_______ angles are the biggest

A

Lone-pair/lone-pair

63
Q

_______ angles are the second biggest

A

Lone-pair/bonding-pair

64
Q

_______ angles are the smallest

A

Bonding-pair/bonding-pair

65
Q

Name the shape of a molecule with 2 electron pairs (& no lone pairs)

A

Linear

66
Q

Draw BeCl2

State the bond angles of the molecule

2 electron pairs (& no lone pairs)

A
67
Q

Name the shape of a molecule with 3 electron pairs (& no lone pairs)

A

Trigonal planar

68
Q

Draw BF3

State the bond angles of the molecule

3 electron pairs (& no lone pairs)

A
69
Q

Name the shape of a molecule with 4 electron pairs (& no lone pairs)

A

Tetrahedral

70
Q

Draw NH4+

State the bond angles of the molecule

4 electron pairs (& no lone pairs)

A
71
Q

Name the shape of a molecule with 3 electron pairs & 1 lone pair

A

Trigonal Pyramidal

72
Q

Draw PF3

State the bond angles of the molecule

3 electron pairs & 1 lone pair

A
73
Q

Name the shape of a molecule with 2 electron pairs & 2 lone pairs

A

Bent

74
Q

Draw H2O

State the bond angles of the molecule

2 electron pairs & 2 lone pairs

A
75
Q

Name the shape of a molecule with 5 electron pairs (& no lone pairs)

A

Trigonal Bipyramidal

76
Q

Draw PCl5

State the bond angles of the molecule

5 electron pairs (& no lone pairs)

A
77
Q

Name the shape of a molecule with 4 electron pairs & 1 lone pair

A

Seesaw

78
Q

Draw SF4

State the bond angles of the molecule

4 electron pairs & 1 lone pair

A
79
Q

Name the shape of a molecule with 3 electron pairs & 2 lone pairs

A

T-shaped

80
Q

Draw ClF3

State the bond angles of the molecule

3 electron pairs & 2 lone pairs

A
81
Q

Name the shape of a molecule with 6 electron pairs (& no lone pairs)

A

Octahedral

82
Q

SF6

State the bond angles of the molecule

6 electron pairs (& no lone pairs)

A
83
Q

Name the shape of a molecule with 4 electron pairs & 2 lone pairs

A

Square planar

84
Q

Draw XeF4

State the bond angles of the molecule

4 electron pairs & 2 lone pairs

A
85
Q

Describe how you can find out how many bonding and
lone electron pairs there are on central atom of molecule

A
  1. Find central atom (one that all other atoms are bonded to)
  2. Work out no. of electrons in outer shell of central atom (use periodic table)
  3. Add one to this number for every atom that central atom is bonded to
  4. Divide by 2 to find no. of electron pairs on central atom
  5. Compare no. of electron pairs to no. of bonds to find no. of lone pairs and no. of bonding pairs on central atom
86
Q

Predict the shape of the molecule H2S (show all your steps)

A
87
Q

Define Electronegativity

A

The power of an atom to attract a pair of electrons in a covalent bond

88
Q

____ is most electronegative element

A

Fluorine

89
Q

How are polar covalent bonds created?

A

In a covalent bond between 2 atoms of different electronegativities, bonding electrons will be pulled towards the more electronegative atom = makes bond polar

90
Q

Covalent bond between 2 atoms of same element is ____

A

non-polar

91
Q

Why is a covalent bond between 2 atoms of same element non-polar?

A

∵ atoms have equal electronegativities = electrons equally attracted to both nuclei

92
Q

Some elements (e.g. C & H) have very similar electronegativities = bond essentially ____

A

non-polar

93
Q

In a polar bond, difference in electronegativity between 2 atoms causes a ____ ___ to form

A

permanent dipole

94
Q

What is a dipole?

A

Difference in charge between 2 atoms caused by shift in electron density in bond

95
Q

Greater difference in electronegativity between atoms = …

A

more polar the bond

96
Q

What makes a polar molecule?

A

When a molecule contains polar bonds that give an uneven distribution of charge across the whole molecule

97
Q

When are molecules with polar bonds not polar and why is this?

A

When polar bonds are arranged symmetrically in molecule = charges cancel out & there’s no permanent dipole

98
Q

Name 3 Intermolecular Forces

A
  • Permanent dipole-dipole forces
  • Van der Waals forces or induced dipole-dipole forces
  • Hydrogen bonding
99
Q

What type of molecules have permanent dipole-dipole forces?

A

Polar molecules

100
Q

Describe how permanent dipole-dipole forces form

A

In a substance made from molecules with permanent dipoles = they’ll be weak electrostatic forces of attraction between δ+ and δ- charges on neighbouring molecules

101
Q

Explain why if you put a charged rod next to a jet of polar liquid (e.g water), the liquid will move towards the rod

A
  1. ∵ polar liquids contain molecules with permanent dipoles
  2. (Doesn’t matter if rod is postively or negatively charged)
  3. Polar molecules in liquid can turn around so the opposite charged end is attracted towards the rod
102
Q

Where are Van der Waals forces found?

A

Found between all atoms and molecules

103
Q

Describe how Van der Waals forces form

A
  1. Electrons in charge clouds = always moving quickly
    1. At any moment, electrons in atom are likely to be more to one side than the other
    2. At this moment = atom has temporary dipole
  2. Dipole causes another temporary dipole in opposite direction on neighbouring atom
    1. 2 dipoles are attached to each other
  3. 2nd dipole causes another dipole in 3rd atom
  4. Dipoles are being created and destroyed constantly ∵ electrons are constantly moving
    1. Overall effect = atoms are attracted to each other
104
Q

Van der Waals forces hold molecules together in ___

A

lattice

105
Q

Describe and explain the structure of iodine at room temp

A
  1. Iodine atoms are held together in pairs by strong covalent bonds to from I2 molecules
  2. Molecules held together in molecular lattice arrangement by weak van der Waals attractions (this causes iodine to be solid at room temp.)
106
Q

Name 3 factors that affect the strength of the Van der Waals forces

A
  1. Size of molecules
  2. Shape of molecules
  3. Number of electrons
107
Q

Explain how the size of molecules affects the strength of van der Waal forces

A

Larger molecules = larger electron clouds = stronger van der Waals forces

108
Q

Explain how the shape of molecules affects the strength of van der Waal forces

A

Long, straight molecules lie closer than branched ones = closer together 2 molecules get = stronger the forces between them are

109
Q

When does hydrogen bonding occur?

A

When hydrogen is covalently bonded to fluorine, nitrogen or oxygen

110
Q

Hydrogen Bonding is the _____ intermolecular force

A

Strongest

111
Q

Describe hydrogen bonding

A
  1. F, N & O = very electronegative ∴ they draw bonding electrons away from hydrogen atom
  2. Bond is polarised + hydrogen has high charge density = hydrogen atoms form weak bonds with lone pair of electrons on F, N or O atoms or other molecules
112
Q

Molecules with H-bonding usually contain ____ or ____ groups

A

-OH or -NH groups

113
Q

Draw hydrogen bonding occuring in water

A
114
Q

Draw hydrogen bonding occuring in ammonia

A
115
Q

Substances with h-bond have ____ boiling/melting points than similar molecules

A

Substances with h-bond have higher boiling/melting points than similar molecules

116
Q

Why do substances with h-bond have higher boiling/melting points than similar molecules?

A

∵ of extra energy needed to break h-bonds

117
Q

Anomalously high boiling points of H2O, NH3 & HF are
caused by ___ ____ between molecules

A

Anomalously high boiling points of H2O, NH3 & HF are
caused by hydrogen bonding between molecules

118
Q

Explain why ice is less dense than liquid water

A
  1. As liquid water cools to form ice, molecules make more h-bonds & arranged themselves into regular lattice structure
  2. In this structure, H2O molecules are further apart on average than molecules in liquid water
119
Q

Explain why simple covalent compounds have lower melting/boiling points than macromolecules. (4)

A
  1. To melt/boil simple covalent compound = just need to overcome the van der Waals forces between molecules
  2. These forces are weak
  3. To melt/boil macromolecules = many covalent bonds have to be broken
  4. Covalent bonds = strong
120
Q

Explain how the solubility of a substance in water depends on type of particles it contains

A
  • Water = polar solvent
  • ∴ substances that are polar or charged will dissolve in it
  • Whereas non-polar or uncharged substances won’t
121
Q

Fill in the blanks

A
122
Q

Fill in the blanks

(3 examples)

A
123
Q

Fill in the blanks

(3 examples)

A
124
Q

Fill in the blanks

(3 examples)

A
125
Q
A
126
Q

Explain why CF4 has a bond angle of 109.5° (2)

A
  • Around carbon = 4 bonding pairs of electrons
  • ∴ these repel equally & spread as far apart as possible
127
Q

State what is meant by macromolecular (1)

A

Means a giant molecule with covalent bonding

128
Q

Predict the shape of AlCl4-. Draw a diagram of the specie to show its 3D shape. Name the shape and suggest a value for the bond angles. Explain your reasoning. (4)

A
129
Q

Application Question

Perfume is a mixture of fragrant compounds dissolved in a volatile solvent. When applied to the skin the solvent evaporates, causing the skin to cool for a short time. After a while, the fragrance may be detected some distance away. Explain the observations. (4)

A
  1. Solvent has low boiling point or weak intermolecular forces
  2. Solvent needs energy, taken from the skin, to overcome intermolecular forces and evaporate
  3. Perfume molecule slowly spreads through the room
  4. By random diffusion of the perfume
130
Q

Draw a diagram to show how 1 molecule of ammonia is attracted to 1 molecule of water. (3)

(hint h-bonding)

A
131
Q

Draw NH3BCl3

A
132
Q

What type of bonding does H3O+ have?

A

Dative Covalent Bonding

133
Q

Fill in the blanks

A
134
Q

Showing outer electrons only, draw a dot-and-cross diagram to indicate the bonding in calcium oxide (2)

A

Ionic Bonding

135
Q

Explain why the boiling temperature of PH3 is greater than that of CH4 (3)

A
  • PH3 has dipole-dipole
  • between molecules
  • stronger than in CH4
136
Q

Explain why the H-F bond in HF is polar (2)

A
  • Difference in electronegativity / F more electronegative than H
  • Bonding pair of electrons attracted (drawn) towards F (nucleus) in the covalent bond

A-level Chemistry (34 decks)

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