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Higher Chemistry > Unit 1: Chemical Change And Structure > Flashcards

Unit 1: Chemical Change And Structure Flashcards

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1
Q

What is a catalyst?

A

A substance that increases the rate of a chemical reaction

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2
Q

What happens to a catalyst after a reaction occurs?

A

It remains the same

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3
Q

Name the two types of catalyst

A

Biological catalyst and chemical catalyst

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4
Q

Define a chemical and a biological catalyst

A

A chemical catalyst is elements or substances added to induce a reaction

A biological catalyst is naturally occuring

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5
Q

Define a heterogenous catalyst

A

It’s a different state from the reactants

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6
Q

Define a homogenous catalyst

A

It’s in the same state as the reactants

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7
Q

What do catalysts do?

A

They improve geometry and more successful collisions with lower activation energy

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8
Q

Name the 4 types of bonding amongst the first 20 elements

A

Metallic
Covalent network
Covalent molecular
Monatomic

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9
Q

Which group have monatomic bonding and what is it?

A

Group 8 have monatomic bonding and it means they consist of single, unbounded atoms

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10
Q

Name the 7 elements amongst the first 20 have metallic bonding

A

Lithium
Beryllium
Sodium
Magnesium
Aluminium
Potassium
Calcium

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11
Q

What are covalent networks and what’s their state? (What do covalent networks consist of?

A

They consist of thousands of atoms joined together with covalent bonds and they are solids

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12
Q

What’s the structure of diamond? (Arrangement)

A

Tetrahedral arrangement, all four electrons used to make bonds

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13
Q

Whats the structure of graphite? (Arrangement)

A

3 covalent bonds arranged in hexagonal rings, 4th electron is delocalised in layers

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14
Q

Name 4 of the 20 elements that are diatomic

A

Nitrogen
Oxygen
Fluorine
Chlorine

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15
Q

What are the diatomic molecules at room temperature? And what’s their size?

A

Gases and small

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16
Q

What are the diatomic molecules in the first 20 elements examples of?

A

Covalent molecular molecules

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17
Q

What are the 3 elements in the first 20 that are covalent molecular solids?

A

Carbon
Phosphorus
Sulphur

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18
Q

What forces do the 3 elements that are covalent molecular solids have? What’s their state at room temperature?

A

Some weak attraction forces, and can be solid at room temperature

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19
Q

How was the periodic tables invented? (Order)

A

Invented in order of increasing atomic mass

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20
Q

What does melting point and boiling point depend on?

A

The strength of forces between the particles

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21
Q

Why are elements on the left of the periodic tables generally stronger?

A

They have stronger intermolecular forces

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22
Q

What happens to the melting point and boiling point as you go down group 1?

A

It decreases

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23
Q

What’s the density of a substance? Calculation?

A

Mass per unit volume, in g/cm3

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24
Q

Define covalent radius

A

Half the distance between the centres, nuclei, of 2 bonded atoms

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25
Q

Other name for covalent radius?

A

Atomic size

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26
Q

Here’s covalent radius measured? Unit?

A

Measured in picometers

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27
Q

What happens as you go across a period?

A

As we go across a period the nuclear charge and number of outer electrons increases

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28
Q

What happens as you go down a group?

A

As we go down a group, the number of electron shells or energy levels increases, but the number of outer electrons stays the same

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29
Q

In terms of the trends in atomic size: what happens as you go across a period?

A

The atomic size decreases as nuclear charge increases and attracts the outer electrons closer to the nucleus

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30
Q

In terms of trends in atomic size: what happens as you go down a group?

A

As you go down a group the atomic size increases as an extra electron shell is added

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31
Q

Define the term “First ionisation energy”

A

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

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32
Q

What does the first ionisation of magnesium look like?

A

Mg (g) ———> Mg+(g) + e-

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33
Q

Which elements are involved in metallic bonding?

A

Metals

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34
Q

What happens to first ionisation energy as you go down a group?

A

It decreases

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35
Q

What happens to the first ionisation energy as you go across a period?

A

It increases

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36
Q

Define “second ionisation energy”

A

The amount of energy required to remove one mole of electrons from one mole of gaseous 1+ ions

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37
Q

Show the second ionisation energy formula for magnesium

A

Mg (g)+ ——> Mg(g) 2+ + e-

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38
Q

Why’s the 2nd ionisation energy of an element always greater than the 1st?

A
  • In the 2nd ionisation energy of negative electrons are being removed from + ions rather than = ones
  • in the positive ion there’s a greater attraction for the electron so more energy is needed to remove the 2nd mole of electrons
  • successive ionisation energies increases as the atom becomes more positive
  • there’s a large jump in ionisation energy when the electron to be removed comes from a new shell, closer to the nucleus
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39
Q

Define electronegativity

A

It’s a measure of an atoms attraction for the shared pair of electrons in a bond

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40
Q

Trends in terms of electronegativity: what happens across a period?

A

Electronegativity increases, the charge in the nucleus increases

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41
Q

Trends in terms of electronegativity: what happens down a group?

A

As we go down a group, electronegativity decreases, atoms have a bigger radius

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42
Q

What is covalent bonding a result of?
What do atoms share?

A

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

Atoms share electrons

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43
Q

What’s ionic bonding and what’s its force of attraction?

A

It’s the electrostatic force of attraction between the oppositely charged ions

Strong forces of attraction

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44
Q

What type of ions is ionic bonding between? What do ionic compounds form?

A

It’s between positive and negative ions

Ionic compounds form lattice structures of oppositely charged ions

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45
Q

How do polar covalent bonds form?
What’s it sometimes called?

A

When distribution of electrons is unequal a polar covalent bond forms (this is sometimes called a dipole)

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46
Q

What is there in polar covalent bonds? (Electronegativity)

A

There is a difference in electronegativity

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47
Q

What charge does the less electronegative atom develop? Polar covalent bonds

A

A slightly positive charge and vice versa for the more electronegative atom

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48
Q

Where are electrons pulled in polar covalent bonds?

A

Closest to the atom with the greatest electronegativity

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49
Q

How does pure covalent bonding occur?

A

When the elements are the same as one another or the electronegativities are equal or almost equal

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50
Q

What’s the distribution of electrons like in pure covalent bonds?

A

Bonding electrons are equally shared

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51
Q

What are pure covalent and ionic bonding considered to be on the opposite of?
What lays between?

A

The bonding continuum, with polar covalent bonding laying between the two extremes

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52
Q

What happens as the difference in electronegativity between two atoms occurs?

A

The more polar the bond will be, upto 1.9

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53
Q

If the different in electronegativity is 2, what does this mean for the movement of bonding electrons?

A

The movement of bonding electrons from the element with the lower electronegativity is complete resulting in the formation of ions

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54
Q

What type of bond is it if the difference in electronegativity is 0->0.4?

A

Pure covalent bonds

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55
Q

What type of bond is it if the difference in electronegativity is 0.5->1.9?

A

Polar covalent bonds

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56
Q

What type of bonding would it be if the different in electronegativity is 2-»»

A

Ionic bonding

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57
Q

As a general rule, what’s the electronegativity difference in non polar covalent bonds?

A

<0.5

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58
Q

As a general rule, what’s the difference of electronegativity in polar covalent bonds?

A

> 0.5

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59
Q

As a general rule, what’s the difference in electronegativity of ionic bonds?

A

> 1.7

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60
Q

What’s the melting point of ionic lattices and covalent networks?

A

High

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61
Q

What’s the melting point of covalent molecular?

A

Low

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62
Q

What’s the melting point of a metallic lattice?

A

Can vary, usually high

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63
Q

How do ionic lattices conduct electricity? In which states?

A

In liquid and dissolved states. Not solid

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64
Q

How do covalent networks conduct electricity? In which states?

A

They don’t. Only as graphite

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65
Q

How does covalent molecular conduct electricity? In which state?

A

It doesn’t

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66
Q

How do metallic lattices conduct electricity? In which state?

A

They conduct in all states

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67
Q

What type of bond is this?

H - Br
2.2 - 2.8

A

Polar covalent

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68
Q

What type of bond is this?
N —- Cl
3.0 - 3.0

A

Pure covalent (non polar)

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69
Q

What must you take into consideration when there’s a molecule containing more than two atoms? (Polarity in molecules)

What does the determine?

A
  • all polar bonds within the molecule
  • the shape of the molecules

It’s to determine whether it has a permanent dipole ion (an overall polar structure)

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70
Q

Name the two types of covalent molecule (electronegativity)

A

Polar molecules and non polar molecules

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71
Q

Define polar molecules

A

Theses are molecules with overall polarity, one end is clearly more negative than the other

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72
Q

Define non polar molecules

A

These are molecules with no overall polarity

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73
Q

Define a permanent dipole

A

There is a positive and negative end of the molecule

74
Q

If a molecule contains polar bonds, what determines whether a molecule has an overall polarity?

A

Spatial arrangement

75
Q

If a molecule contains polar bonds, what determines whether a molecule has an overall polarity?

A

Spatial arrangement

76
Q

What’s the structure of CO2? Is it polar? Why?

A

O=C=O
It’s non polar because it’s linear so both ends have a slightly negative charge

77
Q

Generally are tetrahedral molecules polar or non polar?

A

Non polar

78
Q

What are intermolecular forces also known as?

A

van der waals forces

79
Q

Where are intermolecular forces present?

A

Between all molecules

80
Q

What are the three types of van der waals forces?

A

1) London dispersion forces (LDFs)
2) Permanent dipole to permanent dipole interaction
3) hydrogen bonding

81
Q

What break when you melt or boil a substance?

A

The van der waals forces, NOT the covalent bonds within the molecule

82
Q

Is it possible for a molecule to have more than one type of van der waals force?

A

Yes

83
Q

LDFs: what do electrons do and what does this mean?

A

Electrons move around atoms randomly which means that more often than not there’s more electrons on one side of the atom than the other

84
Q

Where do LDFs exist?

A

Between all atoms and molecules

85
Q

Why do LDFs arise?

A

Because electrons move around the atoms

86
Q

Which molecules are LDFs the main force?

A

Non- polar molecules

87
Q

What does the uneven distribution of LDFs cause?

A

A slightly negative and a slightly positive charge on either side of the atom, known as a temporary dipole

88
Q

What is an induced dipole and how does it happen? (LDFs)

A

When another atoms is nearby, the electrons of that atom are repelled by the electron density of the temporary dipole, pushing them to the other side of that atom, creating a second dipole. This is called an induced dipole.

89
Q

Define an LDF

A

The forces of attraction between temporary dipoles and induced dipoles caused by movement of electrons in atoms and molecules

90
Q

What does the strength of LDFs depend on?

A

The size of the molecule/atom

91
Q

What does the strength of LDFs depend on?

A

The size of the molecule/atom

92
Q

How are larger dipoles established?

A

Larger atoms and molecules have more electrons. This leads to larger dipoles being established.

93
Q

What is a permanent dipole to permanent dipole interaction and how does it occur?

A

When two polar molecules are near each other there’s an attraction between the negative end of one of the molecules and positive end of the other. This is called a permanent dipole to permanent dipole interaction.

94
Q

What’s stronger? LDFs or permanent dipole to permanent dipole interaction?

A

Permanent dipole to permanent dipole interaction

95
Q

Rank all the van der waals forces by strength
1 = strongest

A

1 - hydrogen bonds
2- permanent dipole to permanent dipole interaction
3 - London dispersion forces (LDFs)

96
Q

Where do hydrogen bonds occur? Van der waals forces

A

Between molecules where there is a very high electronegativity

97
Q

For a molecule to have hydrogen bonding, it must have one or more of which bonds?

A

Nitrogen (N-H)
Oxygen (O-H) All highly polar
Fluorine (F-H)

NOF bonding**

98
Q

Where do permanent dipole to permanent dipole interactions occur?

A

Between polar molecules

99
Q

Do polar molecules react to an electric field?

A

Yes

100
Q

What would happen if you help a charged rod near a stream of a polar liquid?

A

The liquid will be attracted or repelled

101
Q

What is viscosity?

A

The measure of how thick a liquid is

102
Q

Do viscous liquids have strong intermolecular forces?

A

Yes

103
Q

What is an oxidising agent?

A

A substance that accepts electrons, whilst being reduced

104
Q

What does an oxidising agent assist?

A

Oxidation

105
Q

Which elements tend to be oxidising agents?

A

Those with high electronegativities tend to form ions by gaining electrons, so act as oxidising agents

106
Q

What’s a reducing agent?

A

A substance that donates electrons, whilst it is oxidised

107
Q

What does a reducing agent assist?

A

Reduction

108
Q

Which elements tend to be reducing agents?

A

Those which have low electronegativities tend to form ions by losing electrons, so act as reducing agents

109
Q

Which elements tend to be reducing agents?

A

Those which have low electronegativities tend to form ions by losing electrons, so act as reducing agents

110
Q

Which group have the strongest reducing agents?

A

Group 1

111
Q

Which group have the strongest oxidising agents?

A

Group 7

112
Q

Example of a strong (strongest) reducing agent

A

Carbon monoxide

113
Q

Example of oxidising agent

A

Hydrogen peroxide H2O2

114
Q

Uses for oxidising agents

A

They can inactivate viruses and their ideal for bleach due to the oxidation process being an effective means of breaking down coloured compounds

115
Q

Where are the strongest oxidising agents placed in the electrochemical series?

A

Bottom left hand column of the electrochemical series

116
Q

Where are the strongest reducing agents placed in the electrochemical series?

A

The top of the right hand column

117
Q

What’s the balanced redox equation for this?

Al(s) ——-> Al3+ (aq) + 3e-
2H+ (aq) + e- ——-> H2 (g)

A

6H+ + 2AL ——> 2AL3+ + 3H2

118
Q

What’s always cancelled out in redox equations?

A

Electrons

119
Q

What’s miscibility?

A

The ability for a liquid to “dissolve” another liquid

120
Q

What aren’t immiscible liquids miscible?

A

They have a boundary between them
Example: water and hexane
Water is polar, hexane is not

121
Q

Define solvent

A

The liquid in which the substance dissolves

122
Q

What group and one element make covalent network structures?

A

Group 4 and boron

123
Q

What is thermochemistry?

A

The study in changes in heat energy which occur during a chemical reaction

124
Q

What’s an endothermic reaction?

A

The reaction takes in heat

125
Q

What’s an exothermic reaction?

A

A reaction that releases heat into the atmosphere

126
Q

Which has higher products energy after a reaction? Endothermic or exothermic

A

Endothermic because it takes IN heat

127
Q

What is an activated complex?

A

An unstable arrangement of atoms formats as react at bonds are broken and product bonds are formed.

128
Q

When does an activated complex occur?

A

At the maximum point of the potential energy diagram

129
Q

What does it mean if an atom has a symmetrical shape? Polarity…

A

It means there’s no polarity, therefore it’s not permanent

130
Q

How does sulphur exist?

A

S8

131
Q

How does phosphorus exist?

A

P5

132
Q

Define the term “nuclear charge”

A

The force of attraction between the nucleus and outer electrons

133
Q

Name all the diatomic elements

A

Chlorine
Hydrogen
Iodine
Nitrogen
Fluorine
Oxygen
Bromine

134
Q

Alkali metals: what happens as you go down a group?

A
  • Reactivity increases
  • melting point and boiling point decreases
135
Q

Alkali metals: what happens when they react with water? Final product?

A

React with water to release hydrogen and form hydroxides which dissolve in water making alkaline solutions

136
Q

Alkali metals: density?

A

Low density, lithium, sodium and potassium float on water

137
Q

Transition metals: what do they form and are they catalysts?

A

They’re catalysts and they form coloured compounds

138
Q

Transition metals: compared with group 1, melting point, strength, density, reactivity?

A

Melting point - higher minus per hey
Strength - stronger
Density - harder
Reactivity - less reactive and don’t react as vigorously with water or oxygen

139
Q

Transition metals: properties (similar/different) etc

A

Similar properties and some special ones because a lower energy level is being filled in the atoms of elements

140
Q

Noble gases: 3 features (bonding, reactivity, colour)

A

Bonding - monatomic
Reactivity - unreactive
Colour - colourless

141
Q

Halogens: trends down a group? (Mp/bp, reactivity)

A

Reactivity decreases down the group
Melting point and boiling point increase

142
Q

Halogens: what two things can they form?

A

1) ionic salts with metals
2) simple covalent molecules with other non-metals

143
Q

What happens when you put a more reactive halogen and a less reactive halogen in an aqueous solution?

A

The more reactive halogen can displace the less reactive one

144
Q

What’s a fullerene? Bonding

A

Discrete covalent bond

145
Q

Covalent networks, strength?
What three elements can form these?

A

Very strong forces between bonds
3 elements: carbon, boron, silicon

146
Q

Define a covalent bond
(What creates it?)

A

When 2 positive nuclei are held together by their common attraction to a shared pair of electrons

147
Q

Name the 4 factors of reaction rate

A

Temperature
Concentration
Particle size
Use of a catalyst

148
Q

Name 5 signs that a chemical reaction has occurred

A

1 - colour change
2 - smell
3 - change in pH
4 - temperature change
5 - fizzing

149
Q

What must happen for a reaction to occur?

A

Particles must collide with a minimum energy to react, called the activation energy

150
Q

How does pressure affect reaction rate?

A

Increasing the pressure in a reaction involving gases, reduces the volume for particles to move around in. This results in more collisions, therefore increasing the reaction rate.

151
Q

Define activation energy

A

The activation energy, Ea, is the minimum kinetic energy required by colliding particles before a reaction may occur

152
Q

What affect does a catalyst have on activation energy?

A

They lower the required activation energy

153
Q

What is the stored energy that every substance contains called?

A

Enthalpy

154
Q

What must happen for a successful collision to take place?

A

The collision geometry must be correct (the reactant molecules have to be facing the right way! So that the activated complex can be formed

155
Q

Define metallic bonding

A

The electrostatic force of attraction between positively charged ions and delocalised electrons

156
Q

Metal + oxygen —-> ?

A

Metal oxide

157
Q

Metal + water ——> ?

A

Metal hydroxide + hydrogen

158
Q

Metal + dilute acid ——-> ?

A

Salt + hydrogen

159
Q

Name the four states of matter

A

Liquid. Solid, gas and aqueous

160
Q

What is an atomic number?

A

Number of protons

161
Q

What’s a mass number?

A

Protons + neutrons

162
Q

What are atoms made of?

A

Three subatomic particles
Neutrons, protons and electrons

163
Q

What’s an isotope?

A

Atoms with the same atomic number but a different mass number

164
Q

Name the four structures of molecules

A

Linear
Angular
Trigon pyramidal
Tetrahedral

165
Q

An element has the atomic number 11. Using the periodic table you can identify it as?

A

Sodium

166
Q

Label this nuclide notation

35
17 Cl

A

35 —-> mass number
17 —-> atomic number
Cl ——€ symbol

167
Q

What’s an ionic lattice?

A

A regular repeating arrangement of metal and non metal ions which creates compounds with very high melting points

168
Q

The difference between the first ionisation energies of sodium and chlorine is mainly due to the different in?

A

Number of protons

169
Q

Explain why electronegativity values decrease going down group 7

A

Covalent radius increases so attraction of the nucleus for the outer electrons decreases

170
Q

Explain fully why the boiling points of the halogens increase going down group 7 (3 marks)

A

It’s because the intermolecular forces increase going down a group

LDFs are forces between the molecules

The more electrons the stronger the LDFs

171
Q

Explain the decrease in atom size going across the third period from sodium to argon

A

Increase in the number of protons in the nucleus

172
Q

Silicon nitride has a melting point of 1900°C and does not conduct electricity when molten

Explain fully, in terms of structure and bonding, why silicon nitride has a high melting point (2 marks)

A

Silicon nitride is a covalent network.
Strong covalent bonds are broken

173
Q

Explain fully why, of these three chlorides, silicon tetrachloride is the most soluble in hexane.
Silicon tetrachloride, phosphorus and sulfur (2 marks)

A

Silicon tetrachloride and hexane are non-polar.

Silicon tetrachloride is non-polar due to its shape cancelling out

174
Q

The different between the first ionisation energies of sodium and chlorine is mainly due to the difference in the…

A

Number of protons

175
Q

Which of the following contains pure covalent bonds?
CO2, H2S, PH3, CF4

A

PH3

176
Q

A student is carrying out a titration. Which of the following would help the student to accurately observe the end-point?
1) repeating the titration
2) using a white tile under the flask
3) rinsing the flask between titrations
4) disregarding the rough titre

A

Using the white tile under the flask

177
Q

Write the equation for the first ionisation of phosphorus

A

P(g) ——> P+(g) + e-

178
Q

Explain why the melting point of phosphorus P5, is much higher than that of nitrogen N2. Refer to the intermolecular forces involved. 3 marks

A

1) stronger intermolecular forces between phosphorus compared to nitrogen
2) LDFs are the intermolecular forces present
3) There are more electrons in P5 compared to N2

179
Q

Why does silicon dioxide have a high melting point? (2 marks)

A

Silicon dioxide is a covalent network
Strong covalent bonds are broken

180
Q

Explain full why chloromethane has a lower boiling point than water. In your answer you should refer to the intermolecular forces involved

A

Chloromethane has permanent dipole to permanent dipole interactions whereas water has hydrogen bonding

Hydrogen bonding is stronger than permanent dipole to permanent dipole interactions

181
Q

Titanium chloride is a liquid at room temperature. Suggest the type of bonding and structure present in titanium chloride

A

Covalent molecular

182
Q

Write the ion-electron equation for the oxidation of magnesium atoms

A

Mg(s) —-> Mg2+ (aq) + 2e-

Higher Chemistry (2 decks)

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