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Higher Chemistry - 2021/2022 Version 2 > Unit 1 - Chemical Changes and Structure > Flashcards

Unit 1 - Chemical Changes and Structure Flashcards

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1
Q

How do metallic bonds form?

A

Between metals

Positive nuclei attracted to negative electrons

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2
Q

Why do metals conduct electricity?

A

Delocalised electrons

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3
Q

What increases the strength of a metallic bond?

A

Increased number of delocalised electrons, increased charge on metal ions, increased strength of metallic bonds

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4
Q

What is the structure of metallic bonds?

A

Metal lattice, sea of delocalised electrons

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5
Q

Give example of metals that form metallic bonds.

A

Period 1 and 2 (metals)

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6
Q

What are monatomic gases?

A

Gases that exist as separate or discrete atoms

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7
Q

Give examples of monatomic gases.

A

Helium, Hydrogen, Neon, Argon, Krypton, etc

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8
Q

What is a covalent molecular gas?

A

Diatomic elements

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9
Q

How do diatomic elements bond?

A

Covalently

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10
Q

What is a covalent bond?

A

A shared pair of electrons

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11
Q

Give examples of elements that are covalent molecular gases.

A

Diatomic gases

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12
Q

What increases the strength of intermolecular bonds?

A

As the size of the atom increases, the number of electrons also increases, increasing the strength of the intermolecular forces between atoms
This increases the melting and boiling points

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13
Q

Give examples of covalent molecular solids.

A

Phosphorous (P4)

Sulfur (S8)

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14
Q

What is a covalent molecular solid?

A

An element of a set number (P4) bonding with itself covalently

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15
Q

What type of intermolecular force do covalent molecular solids have?

A

Weak London Dispersion Forces

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16
Q

What increases the strength of an intermolecular force?

A

Increased size = increased number of electrons = increased strength

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17
Q

What is a fullerene?

A

Discrete molecules containing 60 or more atoms covalently bonded

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18
Q

What does discrete mean?

A

An compound that has a set amount of atoms

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19
Q

What is a covalent network solid?

A

Atoms of an element that bond to other atoms of the same element

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20
Q

Give examples of covalent network solids?

A

Boron, Carbon (graphite and diamond) and Silicon

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21
Q

Why do covalent network elements have high melting and boiling points?

A

Have extremely strong intramolecular covalent bonds that require a lot of energy to break

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22
Q

Why do monatomic elements exist as a gas at room temperature?

A

Held together by weak intermolecular forces

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23
Q

Why do noble gases have no intramolecular bonds?

A

Have full electron shells so there are no bonding electrons available to form a covalent bond

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24
Q

Why do diamonds not conduct electricity?

A

Every carbon atom is bonded to four other carbons and therefore has no delocalised electrons for electricity to pass through

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25
Q

What structure do diamonds have?

A

Tetrahedral

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26
Q

What structure does graphite have?

A

Hexagonal plate arrangement

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27
Q

What are the three trends observed in the periodic table?

A

Covalent radius
Electronegativity
Ionisation Energy

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28
Q

What is covalent radius?

A

Half of the distance between two nuclei of covalently bonded atoms of the same element

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29
Q

What is the trend for covalent radius across a period?

A

Across a period, the number of protons increase, increasing the nuclear charge, attracting the outer electrons closer to the nucleus, decreasing the covalent radius

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30
Q

What is the trend for covalent radius down a group?

A

Down a group, the number of filled electron shells increases, increasing the shielding effect, preventing the outer electrons from being strongly attracted to the nucleus, increasing the covalent radius

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31
Q

Why is covalent radius a periodic property?

A

Follows a clear trend across a period and down a group

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32
Q

Why is there no value for covalent radius for the noble gases?

A

Monatomic so don’t bond so the covalent radius can’t be calculated

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33
Q

What is electronegativity?

A

A measure of an atom’s nuclear attraction for the electrons in a covalent bond

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34
Q

What is the trend for electronegativity across a period?

A

Across a period, the number of protons increase, increasing the nuclear charge, drawing the outer electrons closer to the nucleus, increasing the electronegativity

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35
Q

What is the trend for electronegativity down a group?

A

Down a group, the number of filled electron shells increases, increasing the shielding effect, preventing the outer electrons from being strongly attracted to the nucleus, decreasing the electronegativity

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36
Q

Why is electronegativity classed as a periodic property?

A

Follows a clear trend across a period and down a group

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37
Q

Why aren’t there electronegativity values for the noble gases?

A

Monatomic, have full electron shells and so there are no bonding electrons

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38
Q

What is ionisation energy?

A

The energy required to remove an electron from an atom in a gaseous state

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39
Q

What is the trend for ionisation energy across a period?

A

Across a period, the number of protons increases, increasing the nuclear charge, attracting the outer electrons closer to the nucleus, thus increasing the ionisation energy

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40
Q

What is the trend for ionisation energy down a group?

A

Down a group, the number of filled electron shells increases, increasing the shielding effect and preventing the outer electrons from being strongly attracted to the nucleus, decreasing the ionisation energy

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41
Q

Is ionisation an endothermic process or an exothermic process?

A

Endothermic as it requires energy

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42
Q

Why is ionisation energy a periodic property?

A

Follows a clear trend across a period and down a group

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43
Q

How do you write ionisation energy equations?

A

1st IE = Mg(g) –> Mg+(g) + e- Enthalpy change = +738kJmol
2nd IE = Mg+(g) –> Mg2+(g) + e- = +1451kJmol
= +2189kJmol

44
Q

Why is there a large difference in ionisation energies between the first and second ionisations of Lithium?

A

Second electrons requires the breaking of a full electron shell that is closer to the nucleus

45
Q

Why isn’t there a fourth ionisation energy for Lithium?

A

Lithium only has three electrons

46
Q

What are ionic compounds?

A

Compounds that form when a metal and a non metal have a large electronegativity difference.

47
Q

What state are ionic compounds in at room temperature?

A

Solid, high melting point

48
Q

What is the order of states of matter?

A

Solid -> liquid -> gas

49
Q

In an ionic compound, what loses an electron and what gains the electron?

A

`Metal loses electron, non metal gains electron

50
Q

When do ionic compounds conduct electricity?

A

When molten or in solution

51
Q

What state are discrete covalent molecular compounds in?

A

Liquid or gas

52
Q

In summary, what are the properties of covalent network solids?

A

High melting and boiling points

Don’t conduct electricity (except graphite)

53
Q

In summary, what are the properties of covalent molecular solids?

A

Low melting and boiling points
Insulators/don’t conduct electricity
P and S

54
Q

In summary, what are the properties of covalent molecular liquids and gases?

A

Low melting and boiling points

Insulators/ don’t conduct

55
Q

What is a pure covalent bond?

A

When there is an equal attraction in each atom (same electronegativity values)

56
Q

Where do electrons sit in a pure covalent bond?

A

In the centre

57
Q

What kind of bond do diatomic elements have?

A

Pure covalent

Temporary dipoles

58
Q

What is a dipole?

A

When one end of a mole becomes slightly positive and the other slightly negative as a result of electrons ‘wobbling’

59
Q

How do you compare electronegativity?

A

The nucleus of __1__ has a greater attraction for bonding electrons than __2__. Electrons will always be nearer __1__. This results in a polar covalent bond because each end has a different charge. The molecule has a permanent dipole

60
Q

What does it mean by ‘least polar’?

A

The compound with the least difference in electronegativity values

61
Q

What is a permanent dipole?

A

An unequal sharing of electrons between molecules

62
Q

What is the intramolecular bonding continuum?

A 
The differences in electronegativity
0 = pure polar covalent
0 - 0.4 = non polar covalent
0.4 - 1.7 = polar covalent
1.7 -> = ionic
63
Q

What determines the overall polarity of a molecule?

A

The symmetry of the charge distribution

64
Q

Are hydrocarbons polar or non-polar molecules?

A

Non-polar

65
Q

How do you decide if a molecule is overall polar or non-polar?

A

Draw out individual compounds
Find electronegativity differences
Decide whether polar or non-polar
See if they can cancel out

66
Q

How do you cancel out polar charges when dealing with symmetry?

A

Take note of polar and non polar bonds
If polar bonds are divisible by 2, they cancel out, making it non polar
(e.g. 3x non-polar bonds, 1x polar bond - don’t cancel out - Polar molecule)
(e.g. 2x non-polar bonds, 2x polar bonds - cancel out - Non polar molecule)
(e.g. 4x polar bonds - cancel out - Non polar)

67
Q

Is water polar or non-polar?

A

Polar

68
Q

Give an experiment used to test whether a liquid is polar or non-polar.

A

Fill burette with liquid on top of a beaker
Place a charged rod near the stream of water
If the water deflects towards the rod, the liquid is polar.
If the water has a normal path, the liquid is non-polar

69
Q

What are the three types of intermolecular forces, from weakest to strongest?

A

London Dispersion Forces
Permanent dipole - permanent dipole interactions
Hydrogen bonding

70
Q

What is the alternative name for intermolecular forces?

A

Van der Waals

71
Q

What is the alternative name for London dispersion forces?

A

Temporary dipole - temporary dipole interactions

72
Q

What determines the type of intermolecular force present in a compound?

A

The intramolecular bonds (the polarity)

73
Q

What compounds have LDF’s?

A

Molecules with pure covalent bonds
Monatomic gases
Diatomic elements
Covalent molecular elements

74
Q

What increases the strength of an LDF? Why?

A

An increase in the number of electrons increases the strength of an LDF as it contributes to the partial charge when dipoles appear

75
Q

What can confirm the strength of an intermolecular force?

A

The melting and boiling points of noble gases

low, weak bonds

76
Q

Where do intermolecular forces occur?

A

Between molecules, between the slightly positive atom of one molecule and the slightly negative atom of another molecule

77
Q

Moving down the halogen group, why does the state change from gases to liquids to solids?

A

Down the group, the number of electrons increase, creating a stronger partial charge when dipoles occur. This increases the strength of the LDF’s and so the molecules are brought closer together, increasing the melting and boiling points

78
Q

Which element has a higher melting and boiling point : Sulphur or Phosphorous? Why?

A

Sulphur
Sulphur has 128 electrons (S8) while phosphorous have 60 (P4). Sulphur therefor has a stronger partial charge and stronger LDF’s
This brings the molecules closer together, increasing the melting and boiling points

79
Q

When do permanent dipole-permanent dipole interaction forces form?

A

When atoms have difference electronegativity values

80
Q

What increases the strength of a permanent dipole-permanent dipole interaction?

A

An increase in the difference of electronegativity values

81
Q

When does hydrogen bonding occur?

A

When a hydrogen atom is bonded to nitrogen, oxygen, or fluorine

82
Q

Hydrogen bond is not the bond between hydrogen and NOF, but the intermolecular forces between each of the molecules

A
83
Q

What is stronger : Intramolecular bonds or Intermolecular bonds?

A

Intramolecular

84
Q

Why does water have an unusually high melting and boiling point?

A

Hydrogen bonds present

85
Q

What increases the strength of an intermolecular bond?

A

The partial charge on the atoms

86
Q

In terms of bonding, what increases the melting and boiling points?

A

Stronger inter and intramolecular bonds

87
Q

How would you construct an experiment to investigate the factors affecting the viscosity of alcohols?

A

Select alcohols of similar mass with differing number of hydrogen bonds
(e.g. propanol, propan-1,2-diol, propan-1,2,3-triol)
Fill up to set volume in a measuring cylinder with a ball bearing at the bottom of each flask.
Turn the flask upside down and measure the time taken for the ball bearing travel the distance of the flask.
The longer it takes = the more viscous the alcohol is = the more hydrogen bonds are present = stronger intermolecular forces

88
Q

What is miscability?

A

How well something mixes

89
Q

What is solubility?

A

How well something dissolves

90
Q

What substances are the best solvent for non-polar substances?

A

Non polar substances

91
Q

What substance is the best solvent for polar substances and ionic compounds?

A

Polar substances

92
Q

Give examples of polar solvents.

A

Water

Short chain alcohols (ethanol)

93
Q

What is oxidisation?

A

A loss of electrons

Metal

94
Q

What is reduction?

A

A gain of electrons

Non metals

95
Q

How do you write a REDOX equation?

A

Look at pg12 of data book to find equations
(May need to flip it around)
Scale up to get same number of electrons in both equations
Combine equations, leaving out the electrons

96
Q

What is an oxidising agent?

A

A substance that promotes the oxidisation of another substance by accepting electrons from it. Therefore, as it is gaining electrons, the oxidising agent is being reduced.

97
Q

What is the reducing agent?

A

A substance which promotes another substance gain its donated electrons. Therefore, the reducing agent is being oxidised as it has lost electrons

98
Q

Where can you find strong reducing agents?

A

Top right of pg12 of databook

99
Q

Where can you find strong oxidising agents?

A

Bottom left of pg12 in databook

100
Q

Give examples of oxidising agents.

A

Hydrogen peroxide
Potassium Permanganate
Sulphur Dioxide

101
Q

What is an everyday use of hydrogen peroxide?

A

Bleaching of hair and clothes
Teeth whitening
Antiseptic

102
Q

What is an everyday use of potassium permanganate?

A

Antiseptic

103
Q

What is an everyday use of sulphur dioxide?

A

Bleaching

104
Q

What is a redox reaction?

A

A transfer of electrons between atoms

105
Q

How would you write a REDOX equation from scratch?

A

Should be given a formula of a reactant and a product
Balance the main element
Balance oxygen by adding water
Balance hydrogen by adding hydrogen ions
Balance the overall charge by adding electrons

106
Q

How would you do a redox titration?

A

Scale for electrons
Combine equations
Mole ratio
Solve

107
Q

Why would a titration not need an indicator?

A

Self indicating

Higher Chemistry - 2021/2022 Version 2 (3 decks)

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