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Higher Chemistry > Unit 1 > Flashcards

Unit 1 Flashcards

1
Q

Steps to Reacting Masses Calculations

A

Identify two chemicals and write mole statement, convert moles to GFM and use proportion to calculate answer

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2
Q

Reaction rates can be ______ by chemists

A

Controlled

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3
Q

Rates of reaction can be affected by changes in…

A

Pressure, Particle Size, Temperature, Catalysts and Concentration

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4
Q

If reaction rates are too low a manufacturing process will not be ______

A

Economically Viable

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5
Q

If reaction rates are too high there’s a high risk of _____

A

Thermal Explosion

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6
Q

Where’s a reactions rate the quickest and why?

A

At the start of the reaction and as there’s the highest concentration of reactants

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7
Q

Calculating Average Rate

A

Change in concentration or mass or volume/ Time OR 1/Time

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8
Q

Collision Theory

A

States that reactants must collide with correct geometry and posses a minimum energy (Activation Energy) before a successful reaction occurs

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9
Q

In a dilute solution, particles are __A__ apart, meaning ___B___ likely to collide and react

A 
A= Further
B= Less
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10
Q

In a concentrated solution, ___A___ particles are present in the same volume. They are __B__ likely to meet and react

A 
A = More
B = More
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11
Q

(Using the Collision Theory) Explain why increasing the concentration increases the rate

A

It increases the number of particles, meaning more successful collisions

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12
Q

(Using the Collision Theory) Explain why increasing the pressure of a gas increases the rate

A

This increases the concentration of the gas, meaning more successful collisions

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13
Q

(Using the Collision Theory) Explain why decreasing the particle size increases the rate

A

This increases the surface area of the reactant, meaning more successful collisions

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14
Q

Temperature

A

A measure of the average kinetic energy of the particles of a substance

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15
Q

Why does a 10 Degree rise in temperate double a reaction?

A

The average kinetic energy of the colliding particles increases. Meaning more particles collide with >/= Activation energy. Resulting in more successful collisions

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16
Q

Kinetic Energy

A

Movement

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17
Q

Catalysts

A

Increase the rate of a reaction by providing an alternative pathway with a lower activation energy

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18
Q

Temperate is a __A__ of the average __B__ of the particles of a substance

A 
A= Measure
B= Kinetic Energy
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19
Q

The activation energy is the __A__ energy required by __B__ before the reaction will occur

A 
A= Minimum Energy 
B= Colliding Particles
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20
Q

What can be used to explain the effect of changing temperature on the kinetic energy of particles?

A

An energy distribution diagram

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21
Q

Potential energy diagrams show the __A__ change that occurs in a reaction and the __B__

A 
A= Enthalpy
B= Activation energy
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22
Q

Catalysts __A__ up a reaction by __B__ the activation energy

A 
A= Speed
B= Lowering
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23
Q

Activated Complex

A

High energy with a unstable arrangement of atoms

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24
Q

Where’s the activated complex located?

A

Top of the activation energy barrier

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25
Q

Two types of catalysts

A

Heterogeneous and Homogeneous

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26
Q

Heterogeneous Catalysts

A

Catalysts that exist in a different state to the reactants

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27
Q

Homogeneous Catalysts

A

Catalysts that exist in the same state to the reactants

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28
Q

Exothermic reactions have a __A__ change as the products have __B__ energy as the reactants

A 
A= Negative Enthalpy
B= Less
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29
Q

Endothermic reactions have a __A__ change as the products have __B__ energy as the reactants

A 
A= Positive Enthalpy 
B= More
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30
Q

To find Enthalpy change

A

Energy of products - Energy of reactants

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31
Q

To find Activation Energy

A

Activation Energy - Energy of Reactants

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32
Q

What happens when a catalyst is added to a reaction (According to potential energy diagrams)

A

Enthalpy stays the same, Activation energy is lowered

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33
Q

Elements are arranged in order of …

A

Increasing atomic mass and number

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34
Q

Periodicity

A

The regular recurrence of similar properties when elements are arranged in order of increasing atomic number

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35
Q

Explain what happens to the pattern in atomic size going across a period

A

The covalent atomic radius decreases because the nuclear charge increases, attracting the outer electrons more strongly, pulling them closer without any additional shielding.

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36
Q

Shielding

A

The electrons shells between the outer shell and the nucleus “Shield” (The outer electron from the nuclear charge)

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37
Q

Explain the pattern in atomic size going down a group

A

The covalent radius increases (atoms get larger) because of the addition of extra electron shells

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38
Q

First ionisation energy

A

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

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39
Q

Second ionisation energy

A

The energy required to remove one mole of electrons from one mole of 1+ ions in the gaseous state

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40
Q

The ____ group has the higher 1st ionisation energy

A

Noble gases

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41
Q

The ___ group has the lowest 1st ionisation energy

A

Alkali Metals

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42
Q

As you go across each period the ionisation energy ____

A

Increases

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43
Q

Why does the 1st ionisation energy increases going across a period?

A

Nuclear charge increases, attracting the outer electrons more strongly, meaning more energy is needed to remove the outer electron.

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44
Q

As you go down a group the first ionisation energy ____

A

Decreases

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45
Q

Why does the 1st ionisation energy decreases going down a group?

A

The energy level is further away so the electron is attached to it more strongly, meaning less energy to remove it and the outer electrons are more shielded from the nucleus

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46
Q

Why does the ionisation energy increases as more electrons are removed from an element?

A

Because of the larger attraction of the nucleus for the remaining electrons

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47
Q

Electronegativity

A

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

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48
Q

Across a period the electronegativity _____

A

Increases

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49
Q

Looking down a group the electronegativity _____

A

Decreases

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50
Q

Why is there no electronegativity value for the Noble gases?

A

They don’t form bonds

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51
Q

Intramolecular Bonds

A

Bonds formed between atoms to get a full outer shell

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52
Q

Main types of intramolecular bonds

A

Metallic, Ionic, Covalent (Network and Discrete)

53
Q

Only ____ group can’t form intramolecular bonds

A

The noble gases

54
Q

Structure of Metallic Bonds

A

Giant metallic lattice of positive ions in a sea of delocalised outer electrons

55
Q

Delocalised

A

Free to Move

56
Q

Definition of Metallic Bonds

A

Electrostatic attraction between positive ions and delocalised outer electrons

57
Q

Metallic bonds are found in ….

A

Liquid and Solid Metals

58
Q

Strength of Metallic Bonds

A

Varies but mostly strong

59
Q

Metallic bonds get __A__ as you go down a group because the electrons are __B__ from the nuclear charge and are __C__ strongly attracted

A 
A= Weaker
B= Further
C= Less
60
Q

Properties of Metallic Bonds

A
  • Conducts electricity because delocalised electrons are free to move
  • Most have high mpts and bpts due to strong metallic bonds
61
Q

Structure of ionic bonds

A

Lattice of positive and negative ions

62
Q

Definition of ionic Bonds

A

The electrostatic attraction between positive and negative ions

63
Q

Ionic bonds are found in….

A

Compounds containing metal/ammonium and a non-metal

64
Q

Strength of Ionic bonds

A

Very strong bonds leading to high mpts and bpts

65
Q

Properties of ionic bonds

A
  • Conducts when molten in ‘aq’ solution as the ions are free to move
  • Soluble in water
66
Q

The formula of ionic bonds

A

Ratio of -

1 grain of salt : 2+ and negative ions

67
Q

Structure of covalent network bonds

A

A giant lattice of covalently bonded atoms

68
Q

Diamond is a covalent network bond. Properties:

A

Very Hard

69
Q

Graphite is also a covalent network bond. Properties:

A

Weak forces between the layers make it soft and flaky

70
Q

Properties of covalent network bonds

A
  • Used as abrasives as they’re very hard
  • Most are non-conductors (Except graphite)
  • Very high mpts and bpts
71
Q

Only three elements with covalent network structures

A
  • Boron
  • Carbon(Graphite and Diamond)
  • Silicon(Silicon Carbide and Silicon Dioxide)
72
Q

Structure of Covalent Discrete Bonds

A

Discrete molecules held together by weak intermolecular forces

73
Q

Most covalent substances are made up of covalent __A__ molecules. These can be compounds involving __B__

A 
A= Discrete
B= Non- Metals
74
Q

Discrete

A

There is a definite number of atoms bonded together in a molecule

75
Q

Properties of a covalent discrete bond

A
  • Low mpt and bpt

- Non-Conductor

76
Q

Ionic Character

A

The greater the difference int he electronegativity, the more ionic character it has

77
Q

Pure covalent bond

A

The electronegativities of both atoms are identical and the bonding electrons are shared equally between both atoms.

78
Q

Polar Covalent Bonds

A

When there is small difference between the electronegativites of both non-metal atoms in the covalent bond

79
Q

Order of ionic character between pure covalent bonds, ionic bonds and polar covalent bonds

A

Pure covalent bonds, polar covalent bonds and ionic character.

80
Q

The three Van Der Waals Forces

A
  • Ldf Forces
  • Permanent dipole-Permanent dipole interactions
  • Hydrogen Bonding
81
Q

London Dispersion forces

A

The result of electrostatic attraction between temporary dipoles and induced dipoles and are caused by the continual movement of electrons in atoms and molecules

82
Q

How ldf arise

A
  • The electron cloud around the molecule or atom won’t be symmetrical. Meaning a more negative charge on one side of a molecule giving it a temporary dipole
  • This dipole then induces dipoles in neighbouring molecules or atoms. The small positive end attracts to the small negative end of the induced dipole
83
Q

Ldf are the only bond found between __A__. Polar molecules can have __B__ intermolecular forces between them.

A 
A= Non- Polar Molecules
B= Additional
84
Q

Ldf are _____ than covalent bonds

A

Weaker

85
Q

Why is it that the larger the molecule of atom, the stronger the ldf?

A

The electron cloud is larger and has a higher number of temporary dipoles available to induce dipoles in neighbouring molecules or atoms

86
Q

The higher the ldf, the ___ the boiling point

A

Higher

87
Q

As you go down group 7, the molecule __A__ and the strength of the ldf __B__

A

A & B= Increases

88
Q

Permanent dipole - permanent dipole interactions

A

Additional electrostatic forces of attraction between polar molecules

89
Q

To find out if a molecule has permanent dipole-permanent dipole interactions

A
  • Use electronegativities to see whether they are pure or polar covalent
  • Work out the GFM
90
Q

Hydrogen Bonds

A

Intermolecular forces which are formed between molecules in which hydrogen is bonded to a strongly electronegative atom like fluorine, oxygen or nitrogen

91
Q

When hydrogen bonds are present, the substance will have a __A__ mpt and bpt than another substance of __B__ mass

A 
A= Higher
B= Similar
92
Q

Hydrogen bonds are the ___ intermolecular bond

A

Strongest

93
Q

Properties of water that are affected by hydrogen bonding

A
  • High mpt and bpt
  • Surface tension
  • It’s more viscous than ethanol
  • Water is miscible with methanol but immiscible with tetrachloromethane
  • The density of ice is less than water at low temperatures
  • Water is less volatile
94
Q

Why is there high mpts and bpts in water

A

There’s a lot of strong hydrogen bonds between the molecules and must be broken before ice can melt and water can boil

95
Q

Why has water got surface tension

A

Because of hydrogen bonds

96
Q

Why is water more viscous than ethanol

A

As it has more hydrogen bonds between the molecules

97
Q

Surface tension

A

Acts as if it’s surface is encloses in an elastic skin

98
Q

Viscous

A

It’s molecules can’t flow over one another as easily, making it thicker

99
Q

As the number of O-H groups increases, the amount of hydrogen bonds __A__ and the viscosity __B__

A

A & B = Increases

100
Q

Miscible

A

Mixes

101
Q

Immiscible

A

Doesn’t mix

102
Q

Why is water more miscible with methanol than with tetrachloromethane?

A

It has the same type of intermolecular forces

103
Q

Water is _____ with liquids with weak intermolecular forces

A

Immiscible

104
Q

What type of liquid would be miscible with tetracholoromethane?

A

Another non-polar covalent molecule

105
Q

Why does ice float on water?

A

Ice is less dense than water because in ice the water molecules form an open lattice structure held together by lots of hydrogen bonds

106
Q

Why is water less volatile than ethanol?

A

Water forms more hydrogen bonds, meaning more energy is needed for the molecules to separate

107
Q

Volatile

A

Evaporates less easily

108
Q

What molecules or compounds dissolve best in polar solvents like water and alcohols?

A
  • Ionic compounds

- Polar Molecules

109
Q

What dissolves best in non-polar solvents?

A

Non-polar compounds

110
Q

What needs to happen for a solute to dissolve in a solvent?

A

Enough energy must be releases when new bonds form between the solvent and solute to break the bonds between the solute particles

111
Q

How are all acids formed

A

By dissolving polar covalent molecules in water

112
Q

Hydrogen Chloride (HCl) dissolves to form…

A

Hydrochloric Acid

113
Q

Hydrogen Nitrate (HNO3) dissolves to form…

A

Nitric Acid

114
Q

Hydrogen Sulphate (H2SO4) dissolves to form…

A

Sulphuric Acid

115
Q

Why can an ionic compound dissolve in water

A

The attraction between the ions and the polar bonds in water molecules release enough energy to break up the ionic lattice

116
Q

Why can a polar covalent compound dissolve in water?

A

The attraction between the polar molecule and water molecules release enough energy to break up the covalent bonds in the molecule

117
Q

Why are ionic compounds insoluble in non-polar solvents?

A

The non-polar molecules mean they can’t attract ions strongly to break up the ionic lattice

118
Q

Why are non- polar substances insoluble in water?

A

There’s no strong attraction between non-polar molecules and water, so there’s not enough energy released to break the covalent bonds within it

119
Q

Why are non- polar substances soluble in non-polar solvents?

A

When new ldf form between the two molecules, enough energy is released to break the ldf between the non-polar substances

120
Q

Elements with high mpt and bpt have….

A

Strong metallic or covalent (Network) bonds

121
Q

Elements with low mpt and bpt have…

A

Ldf

122
Q

In general the mpt and bpt of elements __A__ across period 2 and 3 reaching a __B__ at group 4

A 
A= Increases 
B= Peak
123
Q

Why does the metallic bonds increases going across a period?

A
  • The charge of the ion increases

- The number of outer electrons which become delocalised increases

124
Q

What happens to the mpt and bpt in Group 4?

A

It has the highest mpt and bpt because lots of covalent bonds have to be broken when it melts or boils

125
Q

Why is H and He in period 1 and elements in group 5 to 0 have low mpt and bpt?

A
  • In group 0, weak ldf must be broken between the atoms

- H2, Group 5,6, and 7 elements must have broke ldf between the molecules

126
Q

Explain the mpt and bpt going down group 1

A

It decreases because the metallic bonds get weaker

127
Q

Explain the mpt and bpt as you go down group 7

A

The mpt and bpt increase because the molecules get larger, increasing the surface area and strength of ldf

128
Q

Sublime

A

They change directly from a solid into a gas

129
Q

Which elements are in sublime

A

Carbon and Arsenic

Higher Chemistry (3 decks)

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