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Chemistry > Bonding And Shapes Of Molecules UNIT 1 > Flashcards

Bonding And Shapes Of Molecules UNIT 1 Flashcards

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0
Q

What is ionic bonding

A

The electrostatic force of attraction between ions of opposite charge formed by the transfer of electrons

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1
Q

What is metallic bonding

A

The electrostatic force of attraction between 2 positive metal atoms and the surrounding sea of delocalised electrons

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2
Q

What is covalent bonding

A

The electrostatic force of attraction between 2 negative non metal atoms, consists of sharing a pair of electrons

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3
Q

Examples of giant covalent compounds

A

Graphite

Diamond

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4
Q

What are lone pairs

A

Pairs of electrons in the outer shell which aren’t involved in the covalent bond

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5
Q

What is a similarity between coordinate/ dative covalent bonding and normal covalent bonding

A

They have the same length and strength bonds between the same atoms

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6
Q

Describe coordinate/ dative covalent bonding

A

Normally electrons shared between pair of atoms. In this case only 1 atom provides the pair.
The atom that accepts the pair is an atom that doesn’t have a full outer shell (electron deficient)
The atom that’s donating the pair has a pair of electrons which are not used in bond (lone pair)

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7
Q

What is electron deficiency

A

An atom that doesn’t have a full outer shell

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8
Q

Example of coordinate/dative covalent bonding

A

Ammonia into ammonium

The nitrogen atom in an ammonia molecule donates a pair of electrons to a proton from the hydrogen atom to form ammonium

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9
Q

What is bond polarity

A

Electron pair not always shared equally. If joined atoms are different the nucleus from one atom may attract the electron stronger than other

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10
Q

What is electronegativity

A

The ability to attract the bonding electrons in an atom in a covalent bond

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11
Q

Which is the most electronegative element

A

Flourine

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12
Q

Example of a double covalent bond

A

Oxygen

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13
Q

What is a diatomic molecule

A

They are non polar as atoms have same electronegativities attracting electrons to nucleus

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14
Q

What affects electronegativity

A

Nuclear charge
Distance between nucleus and outer electrons
Shielding of nuclear charge by electrons in inner shells

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15
Q

What results in greater electronegativity

A

The smaller the atom the closer the nucleus is to the shared outer main level electrons

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16
Q

What are dipoles

A

The difference in charge between two atoms which is caused by a difference in electron density

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17
Q

Do all molecules with polar atoms form polar molecules

A

NO depends on shape of molecules and polarity of bonds

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18
Q

What are the three intermolecular forces

A

Van der waal (tempory dipole)
Permanent Dipole-dipole
Hydrogen

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19
Q

Order the intermolecular forces, From strongest to weakest

A

Hydrogen
Permanent dipole dipole
Vdw

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20
Q

Permanent dipole dipole bonding happens between……..

A

Polar molecules

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21
Q

What happens in permanent dipole dipole forces

A

Happens between polar molecules
2 molecules with dipoles will orientate themselves so that atoms with opposite charges are next to each other, so the 2 molecules attract

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22
Q

What type of molecules are needed for permanent dipole dipole forces

A

Polar molecules

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23
Q

Which of the intermolecular forces is present in everything

A

Vdw

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24
Q

Vdw forces are very….

A

Weak

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25
Q

How strong are vdw forces

A

Very weak

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26
Q

The size of the vdw increases with…

A

Number of electrons

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27
Q

What increases the vdw

A

Number of electrons

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28
Q

How do vdw forces work

A

In an atom the electrons are not fixed in a particular place. They occupy orbitals. Electrons can be found anywhere in orbital. Means in any instant the electron density may be more in one direction= temporary dipole. This dipole can cause another temporary dipole in the opposite direction on a neighbouring atom. The 2 dipoles are then attracted to each other. The second dipole can cause another dipole in a third atom. Because the electrons are constantly moving, the dipoles are being created and destroyed all the time. Even though the dipoles keep changing, the overall effect is for the atoms to be attracted to each other.

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29
Q

Why does H2S have a much lower boiling point compared with H2O

A

It only contains vdw forces, compared with H2O which contains hydrogen bonds

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30
Q

What compound is the exception to the rule of hydrogen bonding

A

Trichloromethane

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31
Q

Why can trichloromethane form hydrogen bonds

A

The combined effect of the 3 chlorines ( which are quite electronegative) is enough to leave the hydrogen virtually electron less

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32
Q

Which element can form hydrogen bonds

A

Nitrogen
Oxygen
Fluorine

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33
Q

Why can only nitrogen , oxygen and fluorine form hydrogen bonds

A

The are the only atoms that are electronegative enough to leave hydrogen virtually electron less

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34
Q

When are hydrogen bonds formed

A

When hydrogen is bonded to a very electronegative element

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35
Q

What happens during hydrogen bonding

A

The very electronegative element pulls the bonding electrons away from the hydrogen atom, leaving it with virtually no electron density. This leaves an exposed proton. The bond is so polarised, and the hydrogen
has such a high charge density because it’s so small, that the hydrogen atoms form weak bonds with lone pairs of electrons. Then begins to form a dative covalent bond with atom that has a lone pair

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36
Q

Example of hydrogen bonded compound

A

Water

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37
Q

How does oxygen form hydrogen bonds

A

Oxygen is very electronegative so pulls shared electrons strongly leaving the hydrogen virtually electron less/no electron density and an exposed proton. It can start to from a dative covalent bond with hydrogen

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38
Q

Which is the only atom that can take place in hydrogen bonding

A

Hydrogen, due to its small size

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39
Q

Why can only hydrogen take place in hydrogen bonding

A

Due to its small size

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40
Q

What are compounds

A

Atoms of different elements bonded together

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41
Q

What are the 2 main types of bonding in compounds

A

Ionic and covalent

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42
Q

What happens when atoms of different elements bond together

A

A compound is formed

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43
Q

In ionic bonding, what are the simplest ions

A

The single atoms, which have either lost or gained 1,2 or 3 electrons so they have a full outer shell

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44
Q

An easy way to work out how many electrons are lost from elements

A

Look at periodic table. Elements in same group all have same number of outer electrons. means they lose or gain same number of electrons to achieve full outer shell. Means they all form atoms with same charges

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45
Q

What does electrostatic attraction hold together

A

Positive and negative ions

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46
Q

What force holds together positive and negative ions

A

Electrostatic force of attraction

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47
Q

Sodium chloride and magnesium oxide are…

A

Ionic compounds

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48
Q

Examples of ionic compounds

A

Sodium chloride

Magnesium oxide

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49
Q

Example of giant ionic lattice structure

A

Sodium chloride

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50
Q

What is sodium chloride an example of

A

Giant ionic lattice structure

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51
Q

Ionic crystals are what shape

A

Giant lattices of ions

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52
Q

What type of crystal is a giant ionic lattice

A

Ionic crystal

53
Q

How is the giant lattice structure ‘giant’

A

As it is made up of the same basic unit repeated over and over

54
Q

Different ionic compounds have ………. …….. Structures

A

Differently shaped

55
Q

What does the structure of the ionic compound affect

A

Their physical properties

56
Q

What affects physical properties in ionic compounds

A

The structure of the ionic compound

57
Q

3 physical properties of ionic compounds

A

Conduct electricity
High melting points
Tend to dissolve in water

58
Q

Info on ionic compounds conducting electricity

A

They conduct when molten or dissolved- not when solid
The ions in a liquid are free to move (and they carry a charge). In a solid they are fixed in position by strong ionic bonds

59
Q

Why can’t solid ionic compounds conduct electricity

A

The ions are fixed in a position by strong ionic bonds

60
Q

What allows molten or dissolved ionic compounds to conduct electricity

A

The ions are free to move and also carry a charge

61
Q

Info on ionic compounds and high melting points

A

The giant ionic lattices are held together by strong electrostatic forces. It looks a large amount of energy to overcome these forces, meaning a high melting point (801c for sodium chloride)

62
Q

What causes high melting points in ionic compounds

A

Strong electrostatic forces within the giant lattice

63
Q

Info on ionic compounds and dissolving in water

A

Water molecules are polar, part of the molecule has a small negative charge, and the other part has small positive charge. The water pulls the ions away from the lattice and causes it to dissolve

64
Q

Why do ionic compounds tend to dissolve

A

The water molecules pulls the ions away from the lattice

65
Q

What is a molecule

A

A Group of atoms bonded together

66
Q

What happens when groups of atoms are bonded together

A

A molecule is formed

67
Q

Molecules are the ……….. Parts of a compound that can take part in chemical reactions

A

Smallest

68
Q

What is the smallest part of a compound that can take part in a chemical reaction

A

A molecule

69
Q

Does it matter if the atoms are the same or different in a molecule

A

No, they will still bond together to form a molecule

70
Q

What type of bond holds together molecules

A

Strong covalent bonds

71
Q

Strong covalent bonds hold together…

A

Molecules

72
Q

What are giant covalent structures

A

They are structures which have a huge network of covalently bonded atoms.

73
Q

What is the name of the structure with a huge network of covalently bonded atoms

A

Giant covalent structure

74
Q

Another name for giant covalent structure

A

Macromolecular structures

75
Q

Macromolecular structures are the same as….

They are made up of what….

A

Giant covalent structures

A huge network of covalently bonded atoms

76
Q

Why can carbon atoms form giant covalent structures

A

They can each form 4 strong covalent bonds

77
Q

Because carbon atoms can each form 4 strong covalent bonds, this means they can form….

A

Giant covalent structures/ Macromolecular

78
Q

Structure of graphite

A

The carbon atoms are arranged in sheets of flat hexagons covalently bonded with 3 bonds each. The fourth outer electron of each carbon atom is delocalised.
The sheets of hexagons are bonded together by weak vdw forces

79
Q

Properties of graphite

A

The weak bonds between the layers are easily broken, so layers can slide over each other. Graphite feels slippery, used as dry lubricant and pencils
The delocalised electrons aren’t attached to any particular carbon atoms and are free to move along the sheets so can carry an electric current
The layers are quite far apart compared to the length of the covalent bonds, so graphite has a low density. Used to make strong, lightweight sports equipment
Because of the strong covalent bonds in the hexagon sheets, graphite has a very high melting point ( it sublimes at over 3900K)
Graphite is insoluble in any solvent. The covalent bonds in the sheets are too difficult to break

80
Q

What does sublimes mean

A

The molecule changes straight from a solid to a gas, skipping the liquid stage

81
Q

Why is graphite insoluble in any solvent

A

The covalent bonds within the sheets are too difficult to break

82
Q

Why can graphite conduct electricity

A

The delocalised are not attached to a particular carbon atom so are free to move along the sheets, so can carry an electric current

83
Q

Why does graphite have such a high melting point

A

Due to the strong covalent bonds within the hexagon sheets

84
Q

Why does graphite have a low density

A

The layers are quite far apart compared with the length of the covalent bonds

85
Q

Why does graphite feel slippery

A

There are weak bonds between the layers which are easily broken, so the layers can slide over each other

86
Q

Structure of diamond

A

Each carbon atom is covalently bonded to 4 other carbon atoms. The atoms arrange themselves in a tetrahedral shape (crystal lattice structure)

87
Q

Why do diamond gemstones sparkle

A

It’s structure makes it reflect light

88
Q

Properties of diamond

A

Very high melting point (sublimes at over 3800K)

Extremely hard (used in diamond-tipped drills and saws)

Vibrations travel easily through the stiff lattice = good thermal conductor

It can’t conduct electricity- all the outer electrons are held in localised bonds

Insoluble

89
Q

What does molecular shape depend on

A

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

90
Q

Electron pairs exist as

A

Charge clouds

91
Q

What is a charge cloud

A

An area where you have a really large chance of finding an electron pair

92
Q

What do the electrons do inside the charge cloud

A

They don’t stay still, they are constantly moving around

93
Q

Bonding pairs and lone pairs of electrons exist as….

A

Charge cloud

94
Q

Do Electron charge clouds repel or attract each other? Why?

A

Repel, as they are all negatively charged

95
Q

What affects how much a charge cloud repels other charge clouds

A

The shape of the charge cloud

96
Q

Which repels the most? Lone pair charge clouds or bonding pair charge clouds?

A

Lone pair charge clouds. So the greatest angles are between lone pairs of electrons, and bond angles between bonding pairs are often reduced because they are pushed together by lone pair repulsion

97
Q

What does the Valence-Shell Electron-Pair Repulsion Theory mean

A

That the shapes and bond angles of many molecules can be predicted

98
Q

What is the name of the rule that means that shapes and bond angles of molecules can be predicted

A

The Valence Shell Electron Pair Repulsion Theory

99
Q

Write down the order of the strength of repulsion between different kinds of electron pair

A

Lone pair/ lone pair bond angles are the biggest

Lone pair/ bonding pair bond angles are the second biggest

Bonding pair/ bonding pair bond angles are the smallest

100
Q

Molecules containing hydrogen bonding are usually…..

A

organic, containing OH or NH groups

101
Q

Properties of molecules containing hydrogen bonding

A

High melting and boiling point because of the extra energy needed to break the hydrogen bonds

102
Q

Why is ice less dense then liquid water

A

Ice has more hydrogen bonds, hydrogen bonds are quite long. So the H2O molecules in ice are further apart, making it less dense

103
Q

The greater the difference in electronegativity the ……….. …….. The bond

A

More polar

104
Q

When the elements have pretty similar electronegativities, will the bonds be polar or non polar

A

Non polar, as they are attracted to the nuclei by almost the same amount

105
Q

What happens if you put an electrostatically charged rod next to a jet of a polar liquid

A

The liquid will move towards the rod, as it is polar and the molecule contains permanent dipoles. It doesn’t matter if the rod is positively or negatively charged, the polar molecules in the liquid can turn around so the oppositely charged end is attracted towards the rod

106
Q

Van der Waals forces causes ……. atoms and molecules to be ……… to each other

A

All

Attracted

107
Q

Van der Waals forces can hold molecules in a…..

A

Lattice

108
Q

Example of van der waal lattice

A

Iodine….iodine atoms are held together in pairs by strong covalent bonds to form molecules
But the molecules are then held together in a molecular lattice arrangement by weak van der waals attractions

109
Q

Stronger van der waals forces mean …….. boiling points

A

Higher

Not all van der waals forces are the same strength-larger molecules have larger electron clouds meaning stronger vdw forces. Molecules with greater surface areas also have stronger vdw forces because they have a more exposed electron cloud. Stronger vdw= higher boiling point

110
Q

Metal atoms exist as…..

A

Giant metallic lattice structures

111
Q

Structure of metals

A

Giant metallic lattice structure

The outermost shell of electrons of a metal atom is delocalised-the electrons are free to move about the metal. This leaves a positive metal ion. The positive metal ions are attracted to the delocalised negative electrons. They form a lattice of closely packed positive ions in a sea of delocalised electrons

112
Q

Properties of metals

A

The number of delocalised electrons per atom affect the melting point. The more there are the stronger the bonding will be and the higher the melting point. Mg2+ has two delocalised electrons per atom, so it’s got a higher melting point that Na+, which only has one. The size of the metal ion and the lattice structure also affect the melting point.
As there are no bonds holding specific ions together, the metal ions can slide over each other when the structure is pulled, so metals are malleable and ductile (can be drawn into a wire).
The delocalised electrons can pass kinetic energy to each other= good thermal conductor
Metals also good electrical conductors because the delocalised electrons can carry a current.
Metals are insoluble, except in liquid metals, because of strength of metallic bonds.

113
Q

Physical properties of solids

A

Particles very close together= high density, incompressible

The particles vibrate about a fixed point and can’t move about freely

114
Q

Physical properties of liquids

A

Similar density to solids, virtually incompressible

The particles move about freely and randomly within the liquid, allowing it to flow

115
Q

Physical properties of gases

A

The particles have loads of energy, much further apart compared sign solids and liquids. The density low, very compressible. The particles move about freely, with not a lot of attraction between them, so they’ll quickly diffuse to fill a container

116
Q

Melting and boiling points depend on ………… Between ………

A

Attraction

Particles

117
Q

The closer the particles are the……

A

Greater the density

118
Q

What do delocalised electrons enable structures to do

A

Conduct electricity

119
Q

As water is a polar substance, it tends to only dissolve….

A

Other polar substances

120
Q

Polar substances will be dissolved by water because….

A

Water is also a polar molecule

121
Q

If a solid has a regular structure it’s called a…..

A

Crystal. The structure is a giant lattice

122
Q

What is a crystal

A

A solid with a regular pattern. It has a crystal lattice

123
Q

What is the exception to the rule of covalent bonds not breaking during boiling and melting

A

Giant molecular structures, like diamond

124
Q

Why don’t covalent bonds break during melting and boiling

A

To melt or boil a simple covalent compound you only have to overcome the vdw forces or hydrogen bonds that hold the molecules together

You don’t need to break the much stronger covalent bonds that hold the atoms together in the molecules

That’s why simple covalent compounds have low melting and boiling points

125
Q

Example of a simple covalent molecule in terms of melting and boiling point

A

Chlorine has stronger covalent bonds than bromine. But under normal conditions, chlorine is a gas and bromine is a liquid. Bromine has the higher boiling point because it’s molecules have more electrons, giving stronger vdw forces

126
Q

Bonding models explain…

A

How substances behave

127
Q

Example of a bonding model

A

‘The physical properties of ionic compounds provide evidence that supports the theory of ionic bonding’

They have high melting points-tells you that the atoms are held together by a strong attraction. Positive and negative ions are strongly attracted, so the model fits the evidence

They are often soluble in water but not in non-polar solvents-tells you that the particles are charged. The ions are pulled apart by polar molecules like water, but not by non-polar molecules. Again, the model of ionic structures fits the evidence

128
Q

Limitations of bonding models

A

Dot and cross models of ionic and covalent bonding are great for explaining what’s happening. But it can be oversimplified

Most bonds aren’t purely ionic or purely covalent but somewhere in between. This is due to bond polarisation. Most compounds end up with a mixture of ionic and covalent properties

129
Q

Why is the electron distribution in a covalent bond not symmetrical

A

It can be symmetrical if two identical nonmetal atoms are attracted to the same pair of electrons but if two different elements are bonded, one is always higher in electronegativity and pulls the electron pair closer to itself causing a slight asymmetry

Chemistry (17 decks)

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