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Higher Chemistry > Structure & Reactions > Flashcards

Structure & Reactions Flashcards

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1
Q

Examples of slow everyday reactions

A
  • Milk going sour
  • Iron rusting
  • Bread going mouldy
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2
Q

Examples of fast everyday reactions

A
  • Fireworks
  • Popcorn
  • Starting a car
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3
Q

How do you increase the rate of reaction ~ collision theory

A

DECREASE particle size

INCREASE concentration
INCREASE temperature

ADD a suitable catalyst

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4
Q

What is needed for a chemical reaction to occur?

A
  1. The reacting particles must collide together
  2. Collisions must have sufficient energy to overcome the energy barrier
  3. The reacting particles must have the correct collision geometry
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5
Q

Why does decreasing particle size speed up a reaction?

A

The smaller the particle size, the larger the surface area.

The more surfaces there are, the more collisions can take place, therefore increasing the reaction rate

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6
Q

Why does increasing concentration speed up the rate of reaction?

A

The higher the concentration, the higher the no.of molecules there are present and available for collisions.

Therefore collisions occur more frequently and the reaction goes faster

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7
Q

Why does increasing temperature speed up the rate of reaction?

A

At higher temperatures more particles have beget equal to the activation energy.

The activation energy is the amount of energy required for a successful collision to take place.

Therefore more successful collisions take place and the reaction goes faster

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8
Q

Why do catalysts speed up reaction rate?

A

They lower the activation energy.

That is they reduce the amount of energy particles need in order to achieve successful collisions.

Therefore successful collisions occur more easily, more frequently, and the reaction goes faster.

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9
Q

Formula for average rate of reaction

A

Average rate = change in quantity/ change in time

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10
Q

Formula for relative rate of reaction

A

Rate = 1/t

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11
Q

What is enthalpy (H)

A

A measure of the energy stored in a chemical

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12
Q

What is an exothermic reaction?

A

A reaction in which heat is given off/ lost

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13
Q

What is an endothermic reaction?

A

A reaction in which chemicals absorb energy/ gain energy

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14
Q

What is a reaction profile?

A

Diagrams showing the difference in the enthalpy the reactants have at the start compared with the enthalpy the products have at the end

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15
Q

Formula for enthalpy change

A

Hproducts- Hreactants

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16
Q

What are the units for enthalpy change?

A

kJmol-1

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17
Q

What is activation energy (Ea)?

A

The minimum kinetic energy that colliding molecules must have in order for a reaction to occur

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18
Q

What is an activated complex?

A

An unstable intermediate in which old bonds are partially broken and new bonds are partially formed

Eg. The reaction between hydrogen and bromine

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19
Q

Why do reacting particles need the correct collision geometry to collide successfully?

A

If particles do not collide with the correct collision geometry the activated complex cannot form

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20
Q

What is temperature?

A

Temperature is a measure of the average kinetic energy of the particles in a substance

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21
Q

What is an energy distribution diagram?

A

Diagrams which show the effect that temperature has on the number of successful collisions and hence reaction rate

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22
Q

Where on an energy distribution diagram can the particles with sufficient energy for successful collision be found?

A

On the right side of the activation energy line

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23
Q

What difference does a small increase in temperature make to the number of particles which have energy equal to or greater than the required Ea?

A

It increases the no.of particles which have energy equal to or greater than the Ea

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24
Q

What effect does a small increase in temperature have on the no.of successful collisions?

A

It increases the no.of successful collisions as more particles have the required activation energy

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25
Q

What are catalysts?

A

Catalysts speed up the rate of a chemical reaction and can be recovered after the reaction

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26
Q

What are homogeneous catalysts?

A

Catalysts which are in the same state as the reactants

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27
Q

What are heterogenous catalysts?

A

Catalysts which are in a different state to the reactants

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28
Q

How do catalysts work?

A

Catalysts speed up chemical reactions by providing an alternative reaction pathway which has a lower activation energy

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29
Q

What are the three stages involved with catalysts?

A

Adsorption

Reaction

Desorption

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30
Q

What is adsorption?

A

Reactants form a temporary bond with the catalyst.

This weakens the bonds within the molecules.

Thus, the activation energy is lowered

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31
Q

What is the reaction stage of a catalyst?

A

The molecules react on the catalyst surface.

The collision geometry is more favourable since one of the molecules is fixed

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32
Q

What is desorption?

A

The product molecules leave the catalyst and the vacant site can be occupied by another reaction molecule.

In other words, catalysts can be reused

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33
Q

What type of diagrams can catalysts be shown on?

A

Potential energy diagrams

Energy distribution diagrams

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34
Q

Who constructed the periodic table?

A

Dimitri Mendeleev

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35
Q

What group number are the Noble Gases?

A

0/8

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36
Q

What group number are the Halogens?

A

7

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37
Q

What group number are the Alkaline Earth Metals?

A

2

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38
Q

What group number are the Alkali Metals?

A

1

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39
Q

Why do elements in the same group have similar chemical properties?

A

They have the same number of outer electrons

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40
Q

What is bonding?

A

A term that describes how atoms join together

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41
Q

What is structure?

A

Describes how the atoms in the element are arranged

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42
Q

What are properties of an element?

A

The characteristics of the substance resulting from its bonding and structure, whether physical or chemical

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43
Q

What type of structures do metallic substances form?

A

Lattice structures

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44
Q

What allows the outer electrons of metallic atoms to be delocalised?

A

They’re held loosely

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45
Q

What does metallic bonding consist of?

A

The atoms losing their outer electrons to a common ‘pool’ of delocalised electrons.

As each atom has lost one or more electrons, the atoms become positively charged ions.

The charged metal ions are now attracted to the pool of electrons

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46
Q

Why do Metals conduct electricity?

A

As the electrons are free to move

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47
Q

Why are Metals shiny/ have a lustre?

A

When light is shone onto a metal, the delocalised electrons absorb the energy from the light and then re-emit the light which gives the metal its characteristic shiny appearance

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48
Q

Why are most metallic substances solid at room temperature?

A

Metallic bonding is strong

Due to the energy required to break these bonds and change them into a liquid

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49
Q

What may covalent substances take the form of?

A
  • Discrete covalent molecules

* Covalent molecules

50
Q

What is the trend between the properties of covalent molecular substances?

A

Low melting and boiling points

51
Q

Why do discrete covalent molecular substances have low melting and boiling points?

A

Due to the fact that molecules of these substances are held together by weak intermolecular bonds called London Dispersion Forces.

As these intermolecular forces are weak it doesn’t require a lot of energy to overcome them

52
Q

What structure do sulphur atoms form?

A

Closed, eight membered, puckered rings

53
Q

What structure do phosphorus atoms form?

A

Tetrahedral molecules

54
Q

What are Fullerenes?

A

Discrete, covalent bonded molecules of carbon

Spherical

Buckminsterfullerene

55
Q

What are covalent networks?

A

Huge complex structures of thousands of atoms covalently bonded together

56
Q

What are the four covalent network structures?

A
  • Carbon Diamond
  • Carbon graphite
  • Silicon
  • Boron
57
Q

Why is carbon diamond exceptionally hard and ridged?

A

There are no discrete molecules; all carbon atoms are joined together in the structure

58
Q

Describe the structure of a carbon diamond network

A

Each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement

59
Q

Why does carbon diamond not conduct electricity?

A

There are no free electrons

60
Q

Why is carbon diamond transparent?

A

Due to the way the structure is arranged, ‘light tunnels’ form between the atoms

61
Q

Describe the structure of carbon graphite

A

Graphite has a structure based on three covalent bonds from each atom, forming layers of hexagonal rings

Each carbon atom has a fourth unpaired electron

This causes a pool of electrons

The result is strong covalent bonding within the layers but only weak interactions between the layers

62
Q

Why does graphite conduct?

A

Fourth unpaired electron

Pool of electrons

63
Q

What is silicon oxide?

A

Sand

64
Q

How common is silicon?

A

Second most common element found on earth

65
Q

Example of a silicon compound

A

Silicon oxide - sand

66
Q

What is silicon used for?

A

In electronics in the form of silicon chips

67
Q

Example of a use of boron

A

Pyrex glassware - boron oxide

Strength and ability to withstand high temperatures

68
Q

What makes boron almost as hard as diamond?

A

It’s covalent network structure

69
Q

What’s the trend between the properties of covalent networks?

A

Exceptionally high melting and boiling points

70
Q

Why do covalent networks have high melting and boiling points?

A

The strong covalent bonds within the structure require high quantities of energy to overcome

71
Q

What are monatomic elements?

A

Single atoms

The Noble gases

72
Q

What is the covalent radius?

A

Size of an atom

The covalent radius of an atom is taken to be half the distance between the nucleus of 2 of its bonded atoms

73
Q

Atomic radius ______ decreases as you go across a period

A

Decreases

74
Q

Atomic radius ______ as you go down a group

A

Increases

75
Q

Explain why atomic radii increases on going down a group

A
  • although the no.of protons and therefore the nuclear charge increases, its effects are overpowered by the no. Of electron layers increasing
  • every additional electron layer ‘shields’ the outer electrons from the positive nucleus - they stop them getting as close to the nucleus so the protons pull is weaker
  • This results in atomic size increasing
  • this ‘sheilding’ effect is called screening
76
Q

Explain why atomic radii decreases on going across a period

A
  • As you go across a period, the atomic number increases so the nuclear charge increases
  • That then causes a greater attraction for the electrons and they are pulled in closer
  • This results in the atomic size decreasing
77
Q

What is the covalent bond a result of?

A

Two positive nuclei being held together by their common attraction for the shared pair of electrons

78
Q

What is electronegativity?

A

A measure of the attraction an atom involved in a bond has for the electrons of the bond

79
Q

As you go down a group, electronegativity ______

A

Decreases

80
Q

As you go across a period, electronegativity _____

A

Increases

81
Q

Explain why electronegativity across a period increases

A

Because nuclear charge increases each time we move across a period the attraction between the electrons and the nucleus gets stronger.

In addition, each time the nuclear charge increases the distance between the outer electrons and the nucleus decreases.

Therefore, the attraction between the outer electrons and the nucleus becomes stronger.

Increased nuclear charge and decreased distance from the nucleus means that electronegativity across a period increases.

82
Q

Explain why electronegativity decreases down a group.

A

Because an additional electron cloud is added to an atom each time we move down a group the distance between the outer electrons and the nucleus increases.

Therefore, the attraction between the electrons and the nucleus becomes weaker.

In addition, each time an electron cloud is added it acts like a screen between the nucleus and outer electrons and weakens the attraction between the electrons and the nucleus.

Increased distance from the nucleus and screening means that electronegativity decreases down a group.

83
Q

What is the first ionisation energy of an element?

A

The energy required to remove one mole of electrons from one mole of atoms in the gaseous state

84
Q

As you go down a group, ionisation energy ____

A

Decreases

85
Q

As you go across a period the ionisation energy _____

A

Increases

86
Q

Explain why ionisation energy decreases as you go down a group

A

Because an additional electron cloud is added to an atom each time we move down a group the distance between the outer electrons and the nucleus increases.

Therefore, the attraction between the electrons and the nucleus becomes weaker.

In addition, each time an electron cloud is added it acts like a screen between the nucleus and outer electrons and weakens the attraction between the electrons and then nucleus.

Increased distance from the nucleus and screening means that ionisation energy decreases down a group.

87
Q

Explain why ionisation energy increases as you go across a period.

A

Because nuclear charge increases each time we move across a period the attraction between the electrons and the nucleus gets stronger.

In addition, each time the nuclear charge increases the distance between the outer electrons and the nucleus decreases.

Therefore, the attraction between the outer electrons and the nucleus becomes stronger.

Increased nuclear charge and decreased distance from the nucleus means that ionisation energy across a period increases.

88
Q

Elements with the greatest electronegativity is more likely to ____ electrons to form ____ ions

A

Gain

Negative

89
Q

Elements with the

Smallest electronegativity is more likely to ____ electrons to form ____ ions

A

Lose

Positive

90
Q

Conductivity of ionic substances?

A

Conducts when molten/ in solution due to lattice breaking allowing electrons to move freely

91
Q

What is a pure covalent bond?

A

When atoms of the same element form a covalent bond they share the pair of electrons as they have the same electronegativity

92
Q

What type of covalent bond would hydrogen and hydrogen form?

A

Pure covalent bond

93
Q

Example of pure covalent bond

A

Hydrogen-hydrogen

94
Q

What’s a polar covalent bond?

A

When atoms with different electronegativities form bonds and the electrons are pulled closer to the the atom with the highest electronegativity

95
Q

Example of polar covalent bond

A

Hydrogen fluoride

96
Q

What are Van der Waal’s forces?

A

The weak forces of attraction between atoms

97
Q

What are the 3 types of Van der Waal’s forces?

A
  • London Dispersion Forces
  • Permanent dipole - permanent dipole attractions
  • Hydrogen bonding
98
Q

What’s a temporary dipole and why is it temporary?

A

The entire electron cloud tends to wobble around which means at any one moment the majority of electrons are in one side of the atom, creating a slightly negative side, and the nucleus is on the other side, creating a slightly positive side.

Temporary bc the electron cloud could wobble in another direction a moment later

99
Q

What is the London Dispersion Force?

A

A temporary, weak electrostatic attraction between the DELTA + end of a temporary dipole of one atom and the DELTA - end of a temporary dipole of another atom.

LDF are the weakest type of Van der Waal’s attraction

100
Q

The _____ the size of an atom, the greater the electron cloud wobble

The greater the electron could wobble the greater the ____ in the DELTA+ and DELTA- ends of the dipole, which results in a greater attraction between the two atoms and ___ LDF

A

Greater

Difference

More

101
Q

What are permanent dipole- permanent dipole attractions?

A

Weak electrostatic attractions that exist between polar molecules

102
Q

Why are permanent dipole- permanent dipole attractions stronger than London Dispersion Forces?

A

Because they occur between permanent dipoles rather than momentary (temporary) dipoles

103
Q

What’s hydrogen bonding?

A
  • strongest Van der Waal’s force
  • Weak electrostatic attraction that exists between polar molecules when the molecules contain a hydrogen atom bonded to nitrogen, oxygen, or fluorine
104
Q

Why do N, O, and F form strong hydrogen bonds?

A

They’re all extremely electronegative compared with hydrogen so the bonds they form are highly polar

105
Q

Example of molecule with hydrogen bonding

A

Water

106
Q

What’s ionic character?

A

The bond’s tendency to be ionic

107
Q

Ionic character _____ as the difference in electronegativity increases

A

Increases

108
Q

The larger the difference in electronegativities between bonded atoms, the more ____ the bond will be

A

Polar

109
Q

Where do LDF exist?

A

Between all molecules + atoms, but the only force between non-polar molecules and monatomic elements.

Weakest attractive force.

110
Q

Example of molecules which contain LDF

A

Between halogen molecules and between bible gas atoms

111
Q

Example of molecule which contains permanent dipole- permanent dipole attractions

A

Between ICl molecules

112
Q

Describe viscosity

A

The stronger the intermolecular force of attraction between the molecules the more viscous the liquid will be

113
Q

Why do stronger intermolecular forces result in a more viscous liquid?

A

Bc the stronger the attraction there is between molecules the more tightly the molecules will be held together and the thicker the liquid will be

114
Q

What’s the solubility of a substance?

A

It’s ability to dissolve in a liquid

115
Q

What’s the miscibility of a substance?

A

It’s ability to mix with another substance

116
Q

When are substances said to be immiscible?

A

If they don’t mix and separate into separate layers

117
Q

What’s the general rule for solubility/miscibility?

A

Like dissolves like

118
Q

Polar and ionic substances tend to dissolve in ____ solvents

A

Polar

119
Q

Non polar substances tend to dissolve in _______ substances

A

Non polar

120
Q

Polarity of water?

A

Extremely polar

121
Q

Why is ice less dense than water?

A
  • when water freezes the H bonding between molecules forces them into a hexagonal shape
  • This makes molecules spread out more than they do as liquid

Higher Chemistry (2 decks)

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