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Chem A-level > 3.1.3 Bonding Knowledge > Flashcards

3.1.3 Bonding Knowledge Flashcards

1
Q

what type of bonding is in a substance with only metal elements?

A

metallic bonding in a giant metallic lattice; electrostatic attraction between positive metal ions and delocalised electrons

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2
Q

what type of bonding is in a substance with metal and non-metal elements?

A

ionic bonding in a giant ionic lattice; electrostatic attraction between oppositely charged ions

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3
Q

what type of bonding is in a substance with only non-metal elements?

A

covalent bonding

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4
Q

what type of bonding is in C, Si, or SiO2?

A

covalent in a macromolecular structure; covalent bonds between atoms

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5
Q

what type of bonding is in the substance if it has H-F, H-O, or H-N bonds?

A

covalent bonding between atoms in a simple molecular structure with hydrogen bonding between molecules

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6
Q

what type of bonding is the substance if it has covalent bonds and is polar?

A

covalent bonding between atoms in a simple molecular structure with permanent dipole-dipole attractions between molecules

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7
Q

what type of bonding is the substance if it has covalent bonds and is not polar?

A

covalent bonding between atoms in a simple molecular structure with induced dipole-dipole attractions between molecules

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8
Q

what are the 3 types of bonding?

A

metallic, ionic, covalent

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9
Q

what are the 2 types of covalently bonded structures?

A

macromolecular and simple molecular

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10
Q

what is the structure of metallic bonding?

A

giant metallic lattice

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11
Q

what is the structure of ionic bonding?

A

giant ionic lattice

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12
Q

what is the type of bonding in a simple molecular structure?

A

Covalent bonds between atoms. IMF between molecules.

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13
Q

what are the particles within a metallic structure?

A

positive metal ions, delocalised electrons

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14
Q

what are the particles within an ionic structure?

A

negative and positive ions

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15
Q

what are the particles within a macromolecular structure?

A

atoms

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16
Q

what are the particles within a simple molecular structure?

A

molecules

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17
Q

define metallic bonding

A

Metallic bonding is the strong electrostatic attraction of positive metal ions surrounded by a sea of delocalised electrons

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18
Q

is the bonding stronger in Mg or Na, and why?

A
  • the bonding in magnesium is stronger than in sodium because:
  • Mg²⁺ ions have a higher charge than Na⁺ ions (2+ vs +)
  • Mg²⁺ has double the number of delocalised electrons than Na⁺
  • Mg²⁺ ion is smaller than Na⁺ ion
  • therefore, Mg²⁺ has a greater electrostatic attraction to the sea of delocalised electrons, than Na⁺
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19
Q

what are the properties of metals?

A
  • All metals are good electrical and thermal conductors
  • the majority of metals are very strong
  • metals are malleable and ductile
  • generally high MP and BP
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20
Q

why are metals good electrical and thermal conductors?

A

The delocalised electrons help transfer energy through the metal very efficiently and can flow and carry charge

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21
Q

why are the majority of metals very strong?

A

there is a strong electrostatic attraction between the positive metal ions and delocalised electrons

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22
Q

why are metals malleable and ductile?

A

the layers of metal ions can slide past one another

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23
Q

what makes a metal have a higher MP or BP than another metal?

A

the strength of the metallic bonds

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24
Q

define ionic bonding

A

strong electrostatic attraction between oppositely charged ions

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25
Q

what are the physical properties of an ionic substance?

A
  • high MP/BP
  • electrical conductivity when aqueous or molten
  • brittle, shatter easily
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26
Q

what is the charge on an ammonium ion?

A

1+

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27
Q

what is the formula of an ammonium ion?

A

NH₄⁺

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28
Q

What is the formula of a hydroxide ion?

A

OH⁻

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29
Q

what is the charge on a hydroxide ion?

A

1-

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30
Q

what is the formula of a nitrate ion?

A

NO₃⁻

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31
Q

what is the charge on a nitrate ion?

A

1-

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32
Q

what is the formula of a nitrite ion?

A

NO₂⁻

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33
Q

what is the charge on a nitrite ion?

A

1-

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34
Q

what is the formula of a hydrogencarbonate ion?

A

HCO₃⁻

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35
Q

what is the charge on a hydrogencarbonate ion?

A

1-

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36
Q

what is the formula of a chlorate (I) ion?

A

ClO⁻

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37
Q

what is the charge on a chlorate (I) ion?

A

1-

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38
Q

what is the formula of a chlorate (V) ion?

A

ClO₃⁻

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39
Q

what is the charge on a chlorate (V) ion?

A

1-

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40
Q

what is the formula of a carbonate ion?

A

CO₃²⁻

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41
Q

what is the charge on a carbonate ion?

A

2-

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42
Q

what is the formula of a sulfate ion?

A

SO₄²⁻

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43
Q

what is the charge on a sulfate ion?

A

2-

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44
Q

what is the formula of a sulfite ion?

A

SO₃²⁻

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45
Q

what is the charge on a sulfite ion?

A

2-

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46
Q

what is the formula of a dichromate ion?

A

Cr₂O₇²⁻

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47
Q

what is the charge on a dichromate ion?

A

2-

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48
Q

what is the formula of a phosphate ion?

A

PO₄³⁻

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49
Q

define a covalent bond

A

A covalent bond is a shared pair of electrons between two atoms

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49
Q

what is the charge on a phosphate ion?

A

3-

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50
Q

how many covalent bonds does each carbon have in diamond?

A

4

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51
Q

what is the structure of diamond?

A

macromolecular - tetrahedral

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52
Q

what are the properties of diamond?

A
  • has a very high melting point
  • is very hard
  • is a non-conductor
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53
Q

why is diamond a non-conductor?

A

there are no delocalised electrons or free-moving charged particles to carry current/energy

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54
Q

how many covalent bonds does each carbon have in graphite?

A

3

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55
Q

what is the structure of graphite?

A

macromolecular - arranged in layers which can slide over each other

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56
Q

why can layers slide over each other in graphite?

A

layers held together by weak IMF

57
Q

what are the properties of graphite?

A
  • malleable
  • conductor
  • high melting point
58
Q

why is graphite malleable?

A

arranged in layers which can slide over each other; layers held together by weak IMF

59
Q

why is graphite a conductor?

A

each carbon atom has a delocalised electron which can carry charge/current through the structure

60
Q

why do simple molecular structures have a low BP/MP?

A
  • when a simple molecular substance melts/boils it is the IMF that are broken
  • IMF are much weaker than covalent bonds so simple molecular compounds have low MP
61
Q

write the ionic formula for lithium oxide

A

Li₂O

62
Q

write the ionic formula for calcium nitrate

A

Ca(NO₃)₂

63
Q

how many bonds does beryllium usually form?

A

2

64
Q

how many bonds does boron usually form?

A

3

65
Q

how many bonds does carbon usually form?

A

4

66
Q

how many bonds does nitrogen usually form?

A

3

67
Q

how many bonds does oxygen usually form?

A

2

68
Q

how many bonds does fluorine usually form?

A

1

69
Q

draw the dot-cross diagram for oxygen

A
  • each atom has 6 valence electrons
  • Oxygen likes to form 2 bonds, so must form a double bond
70
Q

draw the dot-cross diagram for N₂H₄

A
  • hydrogen likes to form 1 bond - nitrogen likes to form 3
  • because nitrogen likes to form more bonds, put the bond between nitrogen and nitrogen in first
  • then add the hydrogens
71
Q

Draw the dot-cross diagram for C₂H₂

A
  • carbon likes to form 4 bonds - hydrogen can only form 1 bond
  • therefore the bonds Hydrogen can make are with carbon, and the carbons must triple bond to each other to gain 8
72
Q

what is a lone pair?

A

a pair of valence electrons that are not shared with another atom

73
Q

what is a bonding pair?

A

The electron pair being shared by the atoms

74
Q

what is a coordinate bond?

A

A coordinate bond is a shared electron pair which have both come from the same atom

75
Q

how is a coordinate bond represented?

A

by an arrow - the direction of the arrow points from the atom which donates to pair of electrons, to the atom which accepts the pair of electrons

76
Q

what are the features of a coordinate bond?

A

once a coordinate bond has formed, it behaves in exactly the same way as a covalent bond - same length and strength

77
Q

draw a diagram to represent the coordinate bond in an ammonium ion

A
78
Q

what is Valence Shell Electron Pair Repulsion Theory?

A

Bonding pairs and lone pairs of electrons are arranged in charged clouds that repel each other; to minimise repulsion, electron pairs try to get as far apart as possible

79
Q

summarise the strength of repulsion of different bonds according to VSEPR theory

A
  • lone pairs repel more than bonding pairs
  • lone pair⇒lone pair is strongest repulsion
  • then lone pair⇒bonding pair
  • bonding pair⇒bonding pair is weakest repulsion
80
Q

draw the VSEPR diagram for a water molecule; give its shape name and bond angle

A
  • name: bent/non-linear
  • bond angle: 104.5
81
Q

what is the bond angle in a bent/v-shaped molecule with 2LP and 2BP?

A

104.5

82
Q

what is the shape of a molecule with 2 LP and 2 BP?

A

bent/non-linear

83
Q

draw the VSEPR diagram for boron trifluoride; give its shape name and bond angle

A
  • shape: trigonal planar
  • bond angle: 120
84
Q

what is the bond angle in a trigonal planar molecule?

A

120

85
Q

what is the shape of a molecule with 3 BP?

A

trigonal planar

86
Q

draw the VSEPR diagram for methane; give its shape name and bond angle

A
  • shape: tetrahedral
  • bond angle: 109.5
87
Q

what is the bond angle in a tetrahedral molecule?

A

109.5

88
Q

what is the shape of a molecule with 4 BP?

A

tetrahedral

89
Q

draw the VSEPR diagram for Phosphorus pentachloride; give its shape name and bond angle

A
  • shape: trigonal bipyramidal
  • bond angle: 120 and 90
90
Q

what is the bond angle in a trigonal bipyramidal molecule?

A

120 and 90

91
Q

What is the shape of a molecule with 5 BP?

A

trigonal bipyramidal

92
Q

draw the VSEPR diagram for sulphur hexafluoride; give its shape name and bond angle

A
  • shape: octahedral
  • bond angle: 90
93
Q

what is the bond angle in an octahedral molecule?

A

90

94
Q

what is the shape of a molecule with 6 BP?

A

octahedral

95
Q

draw the VSEPR diagram for Dichlorocarbene; give its shape name and bond angle

A
  • shape: bent/v-shaped
  • bond angle: 117.5
96
Q

what is the bond angle in a bent/v-shaped molecule with 2 BP and 1 LP?

A

117.5

97
Q

what is the shape of a molecule with 2BP and 1 LP?

A

bent/v-shaped

98
Q

draw the VSEPR diagram for ammonia; give its shape name and bond angle

A
  • shape: pyramidal
  • bond angle; 107
99
Q

what is the bond angle in a pyramidal molecule?

A

107

100
Q

what is the shape of a molecule with 3 BP and 1 LP?

A

pyramidal

101
Q

draw the VSEPR diagram for sulphur tetrafluoride; give its shape name and bond angle

A
  • shape name: seesaw
  • bond angle: 119 and 89
102
Q

what is the bond angle in a seesaw molecule?

A

119 and 89

103
Q

what is the shape of a molecule with 4BP and 1 LP?

A

seesaw

104
Q

draw the VSEPR diagram for iodine pentafluoride; give its shape name and bond angle

A
  • shape name: square pyramidal
  • bond angle: 89
105
Q

what is the bond angle in a square pyramidal molecule?

A

89

106
Q

what is the shape of a molecule with 5 BP and 1 LP?

A

square pyramidal

107
Q

what is the shape of a molecule with 1 BP and 2 LP

A

linear

108
Q

what is the shape of a molecule with 2BP?

A

linear

109
Q

draw the VSEPR diagram for iodine trifluoride; give its shape name and bond angle

A
  • shape name: T-shaped
  • bond angle: 90
110
Q

what is the shape of a molecule with 3BP and 2LP?

A

T-shaped

111
Q

what is the bond angle in a T-shaped molecule?

A

90

112
Q

draw the VSEPR diagram for xenon tetrafluoride; give its shape name and bond angle

A
  • shape name: square planar
  • bond angle: 90
113
Q

what is the shape of a molecule with 4BP and 2LP?

A

square planar

114
Q

what is the bond angle in a square planar molecule?

A

90

115
Q

define electronegativity

A

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

116
Q

what is the Pauling Scale

A

used as a measure of electronegativity. Each element is assigned a number between 0 and 4; the higher the number the more electronegative the element is

117
Q

what 3 factors determine electronegativity?

A
  • Nuclear charge
  • Atomic radius
  • Shielding
118
Q

describe the trend in electronegativity across period 2

A
  • nuclear charge increases
  • atomic radius decreases
  • shielding remains the same
  • the ability to attract electrons in a covalent bond increases
119
Q

When and how does polarity occur?

A
  • When there is a difference in electronegativity within a molecule:
  • electron distribution is asymmetrical
  • this produces a polar covalent bond
  • the atom with a greater electron density is given a delta - charge, and atoms with a lower electron density are given a delta + charge
120
Q

How do you work out if a molecule is polar?

A
  • If a molecule has a lone pair, it is polar (excluding square planar)
  • If the central atom of the molecule has different atoms attached, it is polar
  • Charge of an ion doesn’t affect its polarity
121
Q

is BF3 polar or non-polar, and why?

A

dipoles act in exactly opposite directions as the molecule is symmetrical, so the dipoles cancel, therefore non-polar

122
Q

is NH3 polar or non-polar, and why?

A

dipoles do not cancel as the molecule is not symmetrical, therefore polar

123
Q

is CHCl3 polar or non-polar, and why?

A

dipoles do not cancel as the molecule is not symmetrical, therefore polar

124
Q

what are the 3 types of IMF?

A
  • induced dipole-dipole (Van der Waals) forces
  • permanent dipole-dipole forces
  • hydrogen bonding
125
Q

how do you work out high high the BP/MP of a substance will be?

A
  • once you know the strongest type of IMF a molecule will have, you can make statements about which molecule will have a higher MP/BP
  • The stronger the force between the molecules, the more energy required to break the IMF, and the higher the MP/BP
126
Q

when does hydrogen bonding occur?

A

occurs between H (bonded to N, O, or F) and lone pair on a N, O, or F atom on another atom

127
Q

what is the strongest IMF?

A

hydrogen bonding

128
Q

describe how a hydrogen bond arises

A
  • very large difference in electronegativity between oxygen and hydrogen [state which atoms!]
  • creates a dipole on the O-H bond [state which bond!]
  • lone pair on oxygen atom [state which atom!] in one molecule strongly attracts a partially positive hydrogen atom on a different molecule
129
Q

draw a hydrogen bond

A
  • M1: lone pairs on each O atom and delta + and delta - on each O-H bond
  • M2: dotted line shown between lone pair on one molecule and correct H on another
  • M3: O - - - - - H-O in straight line
130
Q

how strong are p.d.d?

A

generally weaker than hydrogen bonding (although a bigger molecule with p.d.d will have higher MP than hydrogen bonding molecule which is smaller)

131
Q

when does a p.d.d arise?

A

occurs between polar molecules

132
Q

how does p.d.d. arise?

A
  • difference in electronegativity leads to bond polarity
  • dipoles don’t cancel, therefore molecule has overall permanent dipole
  • attraction between delta + on one molecule and delta - on another
133
Q

how is p.d.d drawn?

A
134
Q

how strong is i.d.d.?

A

generally weakest force but can be stronger than both hydrogen bonds and p.d.d if the molecule is large, due to the number of electrons increasing within the molecule

135
Q

when does i.d.d. occur?

A

occurs between all molecules (and atoms of noble gases) but is the important force for non-polar molecules (molecules that do not also have other IMFs

136
Q

how does i.d.d. arise?

A
  • Random movement of electrons in one molecule (atom) leads to an uneven distribution of electrons, creating a temporary dipole in one molecule (atom)
  • this induces a dipole in a neighbouring molecule (atom)
  • dipoles attract
137
Q

how is i.d.d. drawn?

A
138
Q

in what 3 things is hydrogen bonding important and why?

A
  • ice is less dense than water because the hydrogen bonds in ice hold the molecules further apart
  • proteins are held in their complex 3D shapes by hydrogen bonds
  • the 2 strands of DNA are held together by hydrogen-bonding; these hydrogen bonds are strong enough to hold the strands together, but weak enough so they can separate during cell division
139
Q

describe the trends in MP/BP across period 3

A
  • general trend is that MP/BP increase and then decrease across the period
  • From Na ⇒ Mg ⇒ Al the MP increases due to stronger metallic bonds - Na forms Na+ ions, Mg forms Mg2+, Al forms Al3+ and atomic radius decreases
  • silicon has highest MP due to its giant macromolecular structure, very similar to diamond, held together by strong covalent bonds; these require a lot of energy to break, hence Si has a very high MP
  • P, S, Cl, Ar all exist as simple covalent molecules. P forms P4 tetrahedrons, S forms S8 rings, Cl forms Cl2 diatomic molecules. As S8 rings are the biggest of the 3 molecules, they have the strongest i.d.d forces, and so have the higher MP/BP. P4 tetrahedrons are next in size order, and so have a slightly lower MP/BP. Cl2 molecules are the smallest of the 3, have weakest i.d.d. forces, and so lowest MP/BP
  • Silicon has the highest MP but Aluminium has the highest BP. In order to melt silicon a lot of energy is needed to break strong covalent bonds and hence a high temperature is needed however once molten, relatively little more energy is then needed to vapourise it, and so the BP isn’t that much higher than MP. compared to Al, once molten a lot of energy is still needed to overcome strong electrostatic metallic bonds, and so Al has a very high BP
140
Q

why is the BP of Br different from Mg, and why is Mg liquid over a much greater temperature range?

A
  • bromine is simple molecular structure with weak i.d.d forces between molecules
  • magnesium has metallic bonding in a giant metallic lattice, and more energy is needed to overcome stronger metallic bonds
  • magnesium has a much greater liquid range because forces of attraction in liquid are stronger
141
Q

Why does I2 have a higher BP than HCl?

A
  • I2 has i.d.d forces as it is non-polar
  • HCl has p.d.d. forces as it is polar
  • the i.d.d. IMF between I2 molecules is stronger than the p.d.d. IMF between HCl molecules because I2 has a lot more electrons than HCl

Chem A-level (13 decks)

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