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1.3 Bonding Flashcards

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1
Q

What is a compound?

A

Compounds are atoms of different elements chemically bonded together.

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2
Q

What is ionic bonding?

A

Ionic bonding is when ions are held together by electrostatic attraction.

(see page 24 in the revision guide for some more information)

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3
Q

What ions are formed in group 1?

A

1+ ions

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4
Q

What ions are formed in group 2?

A

2+ ions

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5
Q

What ions are formed in group 6?

A

2- ions

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6
Q

What ions are formed in group 7?

A

1- ions

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7
Q

What charge is a sulphate ion?

A

SO4 2-

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8
Q

What charge is a hydroxide ion?

A

OH-

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9
Q

What charge is a nitrate ion?

A

NO 3-

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10
Q

What charge is a carbonate ion?

A

CO3 2-

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11
Q

What charge is ammonium?

A

NH 4+

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12
Q

Describe the structure of sodium chloride.

A

(see page 25 in the revision guide)

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13
Q

How does ionic structure explain the behaviour of ionic compounds?

A
  • Ionic compounds conduct electricity when they’re molten or dissolved.
  • Ionic compounds have high melting points.
  • Ionic compounds tend to dissolve in water.

(see page 25 in the revision guide for a detail of each)

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14
Q

What is a molecule?

A

A molecule is two or more atoms bonded together.

Molecules are help together by strong covalent bonds.

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15
Q

How do electrons relate to a covalent bond?

A

A single covalent bond contains one shared pair of electrons.

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16
Q

How does the structure of graphite explain its properties?

A

(see page 26 in the revision guide)

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17
Q

How does the structure of diamond explain its properties?

A

(see page 27 in the revision guide)

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18
Q

What is dative bonding?

A

Dative bonding is where both electrons come from one atom in a covalent bond.

(see page 27 in the revision guide for an example)

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19
Q

Draw the shape of a linear molecule, what is its bond angle?

A

(see page 29 in the revision guide)

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20
Q

Draw the shape of a trigonal planar model, what is its bond angle?

A

(see page 29 in the revision guide)

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21
Q

Draw the shape of a tetrahedral molecule, what is its bond angle?

A

(see page 29 in the revision guide)

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22
Q

Draw the shape of a trigonal pyramidal, what is its bond angle?

A

(see page 29 in the revision guide)

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23
Q

Draw the shape of a bent molecule, what is its bond angle?

A

(see page 29 in the revision guide)

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24
Q

Draw the shape of a trigonal bipyramidal molecule, what is its bond angle?

A

(see page 29 in the revision guide)

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25
Q

Draw the shape of a seesaw molecule, what is its bond angle?

A

(see page 29 in the revision guide)

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26
Q

Draw the shape of an octahedral molecule, what is its bond angle?

A

(see page 29 in the revision guide)

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27
Q

Draw the shape of a square planar molecule, what is its bond angle?

A

(see page 29 in the revision guide)

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28
Q

Predict the shape of the molecule H2S

A

(see page 29 in the revision guide)

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29
Q

What is electronegativity?

A

Electronegativity is the atoms ability to attract the electron pair in a covalent bond.

(see page 30 in the revision guide to see the electronegativity of some elements)

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30
Q

How are covalent bonds polarised?

A

(see page 30 in the revision guide)

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31
Q

Describe Van der Waals forces.

A

(see page 31 in the revision guide)

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32
Q

Describe hydrogen bonding.

A

(see page 32 in the revision guide)

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33
Q

What are the three types of intermolecular forces?

A
  • Induced dipole-dipole or Van der Waals forces.
  • Permanent dipole-dipole forces.
  • Hydrogen bonding.
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34
Q

Draw and describe the structure of metals.

A

Metal elements exist as giant metallic lattice structures.

see page 33 in the revision guide for a diagram and description

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35
Q

How does metallic bonding explain the properties of metals?

A

(see page 33 in the revision guide)

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36
Q

Describe the physical properties of a solid, liquid and a gas.

A

(see page 33 in the revision guide)

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37
Q

What type of bonding is NaCl and MgCl2?

A

Ionic.

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38
Q

What type of bonding is CO2, I2 and H2O?

A

Simple covalent (molecular)

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39
Q

What type of bonding is diamond, graphite and SiO2?

A

Giant covalent (macromolecular)

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40
Q

What type of bonding is Fe, Mg and Al?

A

Metallic.

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41
Q

Do NaCl and MgCl2 have high or low melting points?

A

High.

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42
Q

Do CO2, I2 and H2O have high or low melting points?

A

Low, (involves breaking intermolecular forces but not covalent bonds)

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43
Q

Do diamond, graphite and SiO2 have high or low melting points?

A

High.

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44
Q

Do Fe, Mg and Al have high or low melting points?

A

High.

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45
Q

What is the state of NaCl and MgCl2 at room temperature?

A

Solid.

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46
Q

What is the state of CO2, I2 and H2O at room temperature?

A

May be solid (like I2) but usually liquid or gas.

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47
Q

What is the state of diamond, graphite and SiO2 at room temperature?

A

Solid.

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48
Q

What is the state of Fe, Mg and Al at room temperature?

A

Solid.

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49
Q

Do NaCl and MgCl2 conduct electricity?

A

No because the ions are held in place.

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50
Q

Do CO2, I2 and H2O conduct electricity?

A

No

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51
Q

Do diamond, graphite and SiO2 conduct electricity?

A

No (except graphite)

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52
Q

Do Fe, Mg and Al conduct electricity?

A

Yes (delocalised electrons)

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53
Q

Substance X has a melting point of 1045K. When solid, it is an insulator, but once melted it conducts electricity. Identify the type of structure present in substance X.

A

Ionic.

see page 35 in the revision guide

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54
Q

Draw the shape of BeCl2

A

(see page 29 in the revision guide)

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55
Q

Draw the shape of BF3

A

(see page 29 in the revision guide)

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56
Q

Draw the shape of NH4+

A

(see page 29 in the revision guide)

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57
Q

Draw the shape of PF3

A

(see page 29 in the revision guide)

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58
Q

Draw the shape of H2O

A

(see page 29 in the revision guide)

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59
Q

Draw the shape of PCl5

A

(see page 29 in the revision guide)

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60
Q

Draw the shape of SF4

A

(see page 29 in the revision guide)

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61
Q

Draw the shape of ClF3

A

(see page 29 in the revision guide)

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62
Q

Draw the shape of SF6

A

(see page 29 in the revision guide)

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63
Q

Draw the shape of XeF4

A

(see page 29 in the revision guide)

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64
Q

Give the definition of ionic bonding.

A

Ionic bonding is the electrostatic force of attrition between oppositely charged ions formed by electron transfer.

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65
Q

What ions do metal atoms form?

A

Metal atoms lose electrons to form +ve ions.

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66
Q

What ions do non-metal atoms form?

A

Non-metal atoms gain electrons to form -ve ions.

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67
Q

Is the melting point higher in smaller, higher charged ions or in larger, lower charged ions? Explain your answer.

A

Ionic bonding is stronger and the melting points higher when the ions are smaller and/or have higher charges.

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68
Q

Which has the higher melting point, NaCl or MgO?

A

MgO has a higher melting point than NaCl as the ions involved (Mg2+ and O2-) are smaller and have higher charges than those in NaCl (Na+ and Cl-)

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69
Q

Give the definition of a covalent bond.

A

A covalent bond is a shared pair of electrons.

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70
Q

Give the definition of a dative covalent bond / co-ordinate bonding.

A

A dative covalent bond forms when the shared pair of electrons in the covalent bond come from only one of the bonding atoms.

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71
Q

Draw the 3D structure of H3O+.

A

(see page 1 in the chemrevise revision guide)

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72
Q

Draw the 3D structure of NH4+

A

(see page 1 in the chemrevise revision guide)

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73
Q

Draw the 3D structure of NH3BCl3

A

(see page 1 in the chemrevise revision guide)

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74
Q

What shape is an NH4+ molecule?

A

Tetrahedral

75
Q

Give the definition for metallic bonding.

A

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

76
Q

What 3 main factors affect the strength of metallic bonding?

A
  1. No. protons / strength of nuclear attraction.
    - The more protons, the stronger the bond.
  2. No. delocalised electrons per atom (the outer shell electrons are delocalised)
    - The more delocalised electrons, the stronger the bond.
  3. Size of ion
    - The smaller the ion, the stronger the bond.
77
Q

Which metal has the stronger ionic bonding and higher melting point, Mg or Na?

A

Mg has stronger metallic bonding than Na and hence a higher melting point. The metallic bonding gets stronger because in Mg there are more electrons in the outer shell that are released to the sea of electrons. The Mg ion is also smaller and has one more proton. There is therefore a stronger electrostatic attraction between the positive metal ions and the delocalised electrons and higher energy is needed to break the bonds.

78
Q

What bonding is in a giant ionic lattice?

A

Ionic bonding.

79
Q

Give 2 examples and draw the structure of a compound with a giant ionic lattice structure.

A

Sodium chloride.
Magnesium chloride.

(see page 2 in the chemreivse revision guide)

80
Q

What bonding is in a simple molecular structure?

A

Covalent bonding.

  • With intermolecular forces (van der waals, permanent dipoles, hydrogen bonds) between molecules.
81
Q

Give 5 examples of a molecule with a simple molecular structure.

A
  • Iodine
  • Ice
  • Carbon dioxide
  • Water
  • Methane
82
Q

What bonding is in a macromolecular structure?

A

Covalent bonding.

83
Q

Give 4 examples of a molecule with a macromolecular structure.

A
  • Diamond
  • Graphite
  • Silicon dioxide
  • Silicon
84
Q

What bonding is in a giant metallic lattice?

A

Metallic

- electrostatic force of attraction between the metal positive ions and the delocalised electrons.

85
Q

Give 2 examples of an element with a metallic giant lattice. Draw the structure.

A
  • Magnesium
  • Sodium
    (all metals)

(see page 2 in the chemrevise revision guide)

86
Q

Do giant ionic lattices have a or low melting and boiling point? Why?

A

High- because of the giant lattice of ions with strong electrostatic forces between the oppositely charged ions.

87
Q

Do simple molecular structures have high or low melting and boiling points? Why?

A

Low - because of the weak intermolecular forces between molecules (specify type, e.g. van der waals, hydrogen bonds)

88
Q

Do macromolecular structures have high or low melting and boiling points? Why?

A

High - because of many strong covalent bonds in the macromolecular structure. It takes a lot of energy to break the many strong bonds.

89
Q

Do giant metallic lattices have high or low melting and boiling points? Why?

A

High - because of the strong electrostatic forces between the positive ions and the sea of delocalised electrons.

90
Q

How good is giant ionic lattices solubility in water?

A

Generally good.

91
Q

How good is simple molecular structures solubility in water?

A

Generally poor.

92
Q

How good is macromolecular structures solubility in water?

A

Insoluble.

93
Q

How good is giant metallic lattices solubility in water?

A

Insoluble.

94
Q

How good is giant ionic lattices conductivity when a solid?

A

Poor: the ions can’t move/ fixed in a lattice.

95
Q

How good is simple molecular structures conductivity when a solid?

A

Poor: there are no ions to conduct and electrons are localised (fixed in place)

96
Q

How good is macromolecular structures conductivity when a solid?

A

Diamond and sand are poor because the electrons cannot move, they are delocalised.

Graphite is good because it has free delocalised electrons between layers.

97
Q

How good is giant metallic lattices conductivity when a solid?

A

Good: The delocalised electrons can move through the structure.

98
Q

How good is a giant ionic lattices conductivity when molten?

A

Good: the ions can move

99
Q

How good is a simple molecular structures conductivity when molten?

A

Poor: There are no ions

100
Q

How good is a macromolecular structures conductivity when molten?

A

Poor

101
Q

How good is a giant metallic lattices conductivity when molten?

A

Good

102
Q

What physical state are giant ionic lattices usually in?

A

Solids

Crystalline

103
Q

What physical state are simple molecular structures usually in?

A

Mostly gases and liquids.

104
Q

What physical state are macromolecular structures usually in?

A

Solids.

105
Q

What physical state are giant metallic structures usually in?

A

Shiny metal.

106
Q

Why are giant metallic lattices malleable?

A

Malleable as the positive ions in the lattice are all identical. So the planes of ions can slide easily over one another.

107
Q

How many bonding pairs are in a linear molecule?

A

2

108
Q

How many bonding pairs are in a trigonal planer molecule?

A

3

109
Q

How many bonding pairs are in a tetrahedral molecule?

A

4

110
Q

How many bonding pairs are in a trigonal pyramidal molecule?

A

3

111
Q

How many bonding pairs are in a bent molecule?

A

2

112
Q

How many bonding pairs are in a trigonal bipyramidal molecule?

A

5

113
Q

How many bonding pairs are in a octahedral molecule?

A

6

114
Q

How many lone pairs are in a linear molecule?

A

0

115
Q

How many lone pairs are in a trigonal planar molecule?

A

0

116
Q

How many lone pairs are in a tetrahedral molecule?

A

0

117
Q

How many lone pairs are in a trigonal pyramidal molecule?

A

1

118
Q

How many lone pairs are in a bent molecule?

A

2

119
Q

How many lone pairs are in a trigonal bipyramidal molecule?

A

0

120
Q

How many lone pairs are in a octahedral molecule?

A

0

121
Q

What is the bond angle for a linear molecule?

A

180

122
Q

What is the bond angle for a trigonal planar molecule?

A

120

123
Q

What is the bond angle for a tetrahedral molecule?

A

109.5

124
Q

What is the bond angle for a trigonal pyramidal molecule?

A

107

125
Q

What is the bond angle for a bent molecule?

A

104.5

126
Q

What is the bond angle for a trigonal bipyramidal molecule?

A

120 and 90

127
Q

What is the bond angle for a octahedral molecule?

A

90

128
Q

Give 4 examples of molecules with a linear shape. Draw them.

A

CO2, CS2, HCN, BeF2

129
Q

Give 5 examples of molecules with a trigonal planar shape. Draw them

A

BF3, AlCl3, SO3, N(O3)-, C(O3)2-

130
Q

Give 4 examples of molecules with a tetrahedral shape. Draw them.

A

SiCl4, S(O4)2-, Cl(O4)-, N(H4)+

131
Q

Give 4 examples of molecules with a trigonal pyramidal shape. Draw them.

A

NCl3, PF3, ClO3, H3O-

132
Q

Give 4 examples of molecules with a bent shape. Draw them.

A

OCl2, H2S, OF2, SCl2

133
Q

Give 1 example of molecules with a Trigonal bipryamidal shape. Draw them.

A

PCl5

134
Q

Give 1 example of molecules with a octahedral shape. Draw them.

A

SF6

135
Q

Which electrons repel each other more, lone pairs or bonding pairs?

A

Lone pairs repel more than bonding pairs.

  • If there are no lone pairs in the molecule, the bonding pairs repel each other equally.
136
Q

Explain the shape of water.

A

(see bottom of page 3 in the chemrevise revision guide)

137
Q

How much do lone pairs reduce the bond angle by?

A

Lone pairs repel more than bonding pairs so they reduce the bond angles by about 2.5 degrees per lone pair.

138
Q

How many bonding pairs are in a square planar molecule?

A

4

139
Q

How many lone pairs are in a square planer molecule?

A

2

140
Q

Draw the shape of XeF4. What is the bond angle?

A

Square planar, (see page 4 in the chemreive revision guide)

141
Q

Draw the shape of BrF4. What is the bond angle?

A

(see page 4 in the chemrevise revision guide)

142
Q

Draw the shape of (I3)-. What is the bond angle?

A

(see page 4 in the chemrevise revision guide)

143
Q

Draw the shape of CIF3. What is the bond angle?

A

(see page 4 in the chemrevise revision guide)

144
Q

Draw the shape of SF4. What is the bond angle?

A

(see page 4 in the chemrevise revision guide)

145
Q

Draw the shape of I(F4)+. What is the bond angle?

A

(see page 4 in the chemrevise revision guide)

146
Q

Give the definition of electronegativity.

A

Electronegativity is the relative tendency of an atom in covalent in a molecule to attract electrons in a covalent bond to itself.

147
Q

What are the 4 most electronegative atoms?

A

F, O, N and Cl

148
Q

What is the most electronegative element?

A

Fluorine

149
Q

Does electronegativity increase or decrease across a period? Explain your answer.

A

Electronegativity increases across a period as the number of protons increases and the atomic radius decreases because the electrons in the same shell are pulled in more.

150
Q

Does electronegativity increase or decrease down a group?

A

Electronegativity decrease down a group because the distance between the nucleus and the outer electrons increases and the shielding of inner shells increases.

151
Q

How can you tell when compounds are purely covalent?

A

A compound containing elements of similar electronegativity and hence a small electronegativity difference will be purely covalent.

152
Q

How can you tell when compounds are purely ionic?

A

A compound containing elements of very different electronegativity and hence a very larger electronegativity difference (>1.7) will be ionic.

153
Q

When does a polar covalent bond form?

A

When the elements in the bond have different electronegativities.

154
Q

What happens to the electrons in a polar covalent bond and what does this form?

A

When a bond is a polar covalent bond it has an unequal distribution of electrons in the bond and produces a charge separation, (dipole) partially + and partially - ends.

155
Q

Will the element with the larger electronegativity in a polar compound be the partially + or partially - end?

A

The element with the larger electronegativity in a polar compound will be the partially - end. (e.g. HCl, Cl is the partially - end)

156
Q

Do symmetrical molecules form polar bonds? Explain your answer.

A

A symmetric molecule will not be polar even if the individual bonds within the molecule are polar.

The individual dipoles on the bonds ‘cancel out’ due to the symmetrical shape of the molecule.

There is no net dipole moment. The molecule is non polar.

157
Q

What is a symmetric molecule?

A

A molecule where all bonds are identical and have no lone pairs.

158
Q

Is CO2 a polar molecule?

A

CO2 is a symmetrical molecule and is a non-polar molecule.

159
Q

Is CCl4 a polar molecule?

A

CCl4 is a non-polar molecule.

160
Q

Is CH3Cl a polar molecule?

A

CH3Cl is a polar molecule.

161
Q

Where do Van der Waals’ forces occur and not occur?

A

Van der Waals’ forces occur between all molecular substances and noble gases.
They occur between all simple covalent molecules and the separate atoms in noble gases.

They do not occur in ionic substances.

162
Q

What is another name for Van der Waals’ forces?

A

transient, induced dipole-dipole interactions.

163
Q

How do Van der Waals’ forces form?

A

In any molecule the electrons are moving constantly and randomly. As this happens the electron density can fluctuate and parts of the molecule become more or less negative, (e.g. small temporary or transient dipoles form)

These instantaneous dipoles can cause dipoles to form in neighbouring molecules. These are called induced dipoles. The induced dipole is always the opposite sign to the original.

164
Q

What is the main factor that affects the size of the Van der Waals?

A

The more electrons there are in the molecule the higher the chance that temporary dipoles will form. This makes the Van der Waals stronger between the molecules and so boiling points will be greater.

165
Q

Why do the boiling points increase down group 7?

A

There are increasing numbers of electrons in the bigger molecules causing an increase in the size of the Van der Waals between molecules. This is why I2 is a solid whereas Cl2 is a gas.

166
Q

Why do the boiling points increase across the alkane homologous series?

A

There are increasing numbers of electrons in the bigger molecules causing an increase in the size of the Van der Waals between molecules.

167
Q

How can the shape of a molecule have an effect on the size of the Van der Waals forces?

A

Long chain alkanes have a larger surface areas of contact between molecules for van der Waals forces to form than compared to spherical shaped branched alkanes and so have stronger Van der Waals.

168
Q

Where do permanent dipole-dipole forces occur?

A

Permanent dipole-dipole forces occurs between polar molecules.

169
Q

Why do molecules with permanent dipole-dipole forces have higher boiling points than molecules with Van der Waals forces?

A

It is a stronger bond than Van der Waals.

170
Q

Do asymmetrical molecules form polar molecules?

A

Yes, polar molecules are asymmetric. They have a bond where there is a significant difference in electronegativity between the atoms.

171
Q

Draw two molecules of HCl interacting with each other.

A

(see page 6 in the chemrevise revision guide)

172
Q

Draw two molecules of propanone interacting with each other.

A

(see page 6 in the chemrevise revision guide)

173
Q

Where does hydrogen bonding occur?

A

Hydrogen bonding occurs in compounds that have a hydrogen atom attached to one of the three most electronegative atoms of nitrogen, oxygen and fluorine, which must have an available lone pair of electrons.

e.g. -O-H, -N-H, F-H bond.

There is a large electronegativity difference between the H and the O,N,F

174
Q

Fact: Hydrogen bonding occurs in addition to Van der Waals forces.

and

Permanent dipole-dipole forces occurs in addition to Van der Waals forces.

A

know this?

175
Q

Draw two molecules of HF interacting with each other.

A

(see page 6 in the chemrevise revision guide)

176
Q

Draw two molecules of NH3 interacting with each other.

A

(see page 6 in the chemrevise revision guide)

177
Q

Draw two molecules of H2O interacting with each other.

A

(see page 6 in the chemrevise revision guide)

178
Q

Which is the strongest intermolecular bond?

A

Hydrogen bonding.

179
Q

Which 4 functional groups are able to form hydrogen bonding?

A
  • Alcohols
  • Carboxylic acids
  • Proteins
  • Amides
180
Q

Draw the molecular structure of Iodine. And describe the structure in words.

A

(see page 7 in the chemrevise revision guide)

181
Q

Draw the molecular structure of ice. And describe the structure in words.

A

(see page 7 in the chemrevise revision guide)

182
Q

Draw the macromolecular structure of diamond. And describe the structure in words.

A

(see page 7 in the chemrevise revision guide)

183
Q

Draw the macromolecular structure of graphite. And describe the structure in words.

A

(see page 7 in the chemrevise revision guide)

184
Q

Why do diamond and graphite have a very high melting point?

A

Because of the strong covalent forces in the giant structure. It takes a lot of energy to break the many strong covalent bonds.

Chemistry (36 decks)

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