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PAST PAPER CONTENT Flashcards

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1
Q

Fc6Would Ions with a greater charge be deflected more in a Mass Spectrometer

A

Yes !

Ions with a greater charge would be more deflected in a Mass Spectrometer

Because there is a greater force

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2
Q

Dative Covalent Bonding Definition?

A

A Two center, Two Electron covalent bond in which these two electrons derive from the same atom

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3
Q

How to Draw Carbon Monoxide ?

A

Diagram

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4
Q

Polar vs Non-Polar

A

Hydrocarbons are non-polar

Symmetry = Polar

Polar>0.5 EN

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5
Q

Examples of Polar Molecules?

A

H- N,O,F

H20, NH3, HF

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6
Q

What line is it on the EM if an electron emits energy to N=1, N=2, N=3 ?

A

UV-Radiation n=1

Visible Light Transition n=2

Infrared Radiation Transition n=3

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7
Q

Is Combustion exothermic or endothermic

A

exothermic

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8
Q

Is Decomposition endo or exo?

A

endo

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9
Q

Is neutralisation exo or endo?

A

exo

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10
Q

What three processes go on in Mass Spectrometry?

A

Ionization, Acceleration and Deflection

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11
Q

Outline the three processes in Mass Spectrometry.

A

ionization: sample bombarded with high-energy or high-speed
electrons/OWTTE;

acceleration: electric field/oppositely charged plates;
deflection:

(electro)magnet/magnetic field;

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12
Q

Electrolytic Cell Diagram?

A

See Answer

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13
Q

Define the term activation energy

A

the minimum quantity of energy which the reacting species must possess in order to undergo a specified reaction.

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14
Q

Distinguish between a strong and a weak acid

A

Weak Acid- Partially Ionizes in Solution

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15
Q

Enthalpy Change

A

The Enthalpy change is independent of the route taken

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16
Q

The definition of the standard enthalpy of formation, ΔHf, is:

A

The enthalpy change when 1 mole of a compound is formed from its elements; [1]

In their standard states; [1]

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17
Q

The equation for the formation of the following compounds are?

Methanol

Carbon Dioxide

Ethane

A

C (s) + ½O2 (g) + 2H2 (g) → CH3OH (l)

C (s) + O2 (g) → CO2 (g)

2C (s) + 3H2 (g) → C2H6 (g)

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18
Q

Which properties increase as you go down group 1?

A

Atomic Radius

Electronegativity

First Ionization Energy

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19
Q

Which properties decrease as you go down group 1?

A

Melting Point

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20
Q

What does the strength of Metallic Bonding depend on?

A

Ionic radius and charge

Smaller Ionic Radius –> Stronger Bonding

Greater Charge –> Stronger Bonding

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21
Q

Are enthalpy changes accurate and how do they compare to enthalpy change of formation?

A

Bond Enthalpies are not specific to compounds - averages are taken from the data booklet

Enthalpy Changes of Formation are specific

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22
Q

Equation for Equilibrium Constant?

A

[C}X{D}/[A}X[B]

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23
Q

Does a weak acid have poor conductivity?

A

Yes

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24
Q

Are neutralisation reactions redox?

A

Neutralisation reactions are not redox

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25
Q

Describe the reaction between halogen and ethane?

A

Free radical substitution and homolytic

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26
Q

In the equation Pv = nRT, what is the relationship between n and T?

A

Inverse Relationship

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27
Q

S,P,D,F on the periodic table?

A

SEE

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28
Q

What is true about Strong Acids

A

They are good proton donors and they have weak conjugate bases

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29
Q

What Compund is the exception to the rule that O2 has a -2 Oxidation number?

A

H2O2 - Hydrogen Peroxide where it has an oxidation number of -1

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30
Q

Suggest two reasons why atoms are no longer regarded as the indivisible units of matter.

A

subatomic particles «discovered»
OR

particles smaller/with masses less than atoms «discovered»

OR
«existence of» isotopes «same number of protons, different number of neutrons»✔
charged particles obtained from «neutral» atoms

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31
Q

Lewis Structure for Ammonia

A

See?

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32
Q

Suggest why the Bond enthalpy might not be right?

A

«N-H» bond enthalpy is an average «and may not be the precise value in NH3»

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33
Q

Suggest an experiment to show that one metal is more reactive than the other

MG and ZN

A

Alternative 3
construct a cell with Mg and Zn electrodes ✔
bulb lights up
OR
shows (+) voltage
OR
size/mass of Mg(s) decreases «over time»
OR
size/mass of Zn increases «over time»

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34
Q

Brønsted–Lowry acid.

A

fully ionizes/dissociates ✔
proton/H+ «donor »✔

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35
Q

Explain the trend of ionisation energy across a period

A

Left to right

Ionisation energy increases

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36
Q

Ionisation Energy

A

Minimum energy required to remove the most loosely bound electron of an isolated gaseous atom, positive ion, or molecule

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37
Q

do Metallic Bonds have low ionisation energies?

A

Yes - low Ionisation energies

They lose electrons so removing them is a tough one for them!

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38
Q

State the trend between ionisation energy and group number?

A

Going down a group the ionisation energy decreases

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39
Q

Do Metallic Bonds have a large atomic radii?

A

Yes - Metallic Bonds have laerg atomic radii

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40
Q

Do Metallic Bonds have high EM values?

A

No - on the lower end

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41
Q

Can you arrange the electron configuration in a periodic table?

A

SEE

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42
Q

State the trend between Atomic Radius and a group and explain the reason why?

A

As you go down the group the Atomic Radius increases

This is because there are more Shells

Although there may be greater nuclear charge –> effect counteracted due to a greater number of protons

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43
Q

State the trend between Atomic Radius and the period number? and explain why?

A

Atomic radius decreases across a period

BECAUSE

Nuclear charge increases with no significant change in shielding

Shielding remains approximately constant –. Atoms in the same period have the same number of inner shells

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44
Q

Compare the size of ionic and atomic radii?

A

Ionic radii of positive ions are smaller than their atomic radii

Ionic radii of negative ions are greater than their atomic radii

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45
Q

Why is Na larger than Na+?

A

Na has one extra shell

They both have the same nuclear charge

Greater amount of electron repulsion –> Greater amount of e-

Larger e- cloud in Na because there are more electrons repelling for the same nuclear charge

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46
Q

How does charge change with ionic radius?

A

Positive ions –> Decrease in radius as the charge increases

Negative –> Increase in radius as the charge increases

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47
Q

How Volatile are Ionic Compounds?

A

Low volatility

Strong electrostatic forces of attraction

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48
Q

How soluble are Ionic Compounds in Water?

A

Often soluble in water.

Water is a polar solvent and energy is released when the ions are hydrated by being surrounded by water molecules.

Energy pays back energy required to break apart the ionic lattice

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49
Q

How soluble are Ionic Compounds in non-Polar solvents?

A

Great deal of energy required to break apart the ionic lattice

Energy is not paid back by the energy released when the non-polar solvent forms interactions with ions

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50
Q

Metal Oxide + Water

A

Metal Hydroxide

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51
Q

Metal + Base

A

Salt + Hydrogen

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52
Q

Acid + Metal + Carbonate

A

Salt + Water + Hydrogen

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53
Q

Are Metal Oxides basic or acidic?
Are Non Metal Oxides basic or acidic?

A

Basic
Acidic - non Metal

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54
Q

Outline the electron configuration of a Hydrogen Emission spectrum?

A

E- occupy discrete energy levels

Energy Levels converge

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55
Q

BOND ANGLE DIAGRAM

A

SEE

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56
Q

4 E- Pair and 1 lone pair

A

Trigonal Pyramidal

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57
Q

rate of reaction def

A

the increase in the concentration of product per unit time.

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58
Q

metalloid?

A

Elements with a mixture of metallic and non-metallic physical properties

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59
Q

Closed system

A

A system in which neither matter nor energy can be lost or gained. An equilibrium can never occur in a system that is not closed

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60
Q

Homogeneous reaction

A

A reaction in which all the species involved are in the same phase.

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61
Q

Le Chatelier’s principle

A

If an equilibrium is subjected to a stress, the equilibrium will shift to minimize the effect of the stress.

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62
Q

Amphoteric and Amphriprotic

A

Amphiprotic species specifically refer to species that can donate or accept a proton (H+ ion). So, a molecule or ion that is amphiprotic must have at least one hydrogen atom that can be ionized, allowing it to act as an acid or a base.

On the other hand, amphoteric species refers to substances that can react with both acids and bases, not necessarily by donating or accepting a proton. This means that an amphoteric species can act as an acid in the presence of a base, and as a base in the presence of an acid.

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63
Q

Orbital:

A

A region of space around the nucleus of an atom that is occupied by a maximum of two electrons at any given time.

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63
Q

Orbital:

A

A region of space around the nucleus of an atom that is occupied by a maximum of two electrons at any given time.

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64
Q

Electronegativity

A

Relative measure of the ability an atom has to attract a shared pair of electrons.

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65
Q

Allotropes

A

Occur when an element can exist in different crystalline forms, such as in carbon, which can exist as graphite, fullerene and diamond.

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66
Q

Allotropes

A

Occur when an element can exist in different crystalline forms, such as in carbon, which can exist as graphite, fullerene and diamond.

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67
Q

Electron affinity

A

Enthalpy change when an electron is added to an isolated atom in the gaseous state.

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67
Q

Electron affinity

A

Enthalpy change when an electron is added to an isolated atom in the gaseous state.

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68
Q

Buffer

A

A solution that resists changes in pH when small amounts of acid or alkali are added to it.

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69
Q

Enthalpy of Combustion

A

The energy released when one mole of a compound is burned in excess oxygen.

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70
Q

Standard enthalpy of formation

A

The energy change when one mole of a compound is formed under standard conditions from its constituent elements in their standard states.

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71
Q

Standard enthalpy of solution

A

: The energy change when one mole of a substance is dissolved in an infinite amount of water under standard conditions.

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72
Q

Aufbau principle

A

The principle that states that lowest energy levels are filled first.

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73
Q

Aufbau principle

A

The principle that states that lowest energy levels are filled first.

74
Q

Ionic Radius trends down the group?

A
  • As the number of shells filled
  • Valence electrongs less strongly attracted
  • Due to Greater Shielding
  • Distance from the nucleus
  • Increase in Atomic Radius
75
Q

Atomic Radius Trends across a period?

A
  • As the number of proton increases
  • increasing nuclear charge attracts the valence electrons more strongly
  • Decrease in Atomic radius
76
Q

Ionization Energy definition?

A
  • Minimum energy required
  • to eject an electron
  • out of a neutral atom
  • in its ground states
77
Q

Factors that affect ionization energy?

A
  • Force of attraction between the valence electron and the nucleus
  • The stronger the force, the more energy is required to pull it away
78
Q

Ionization Energy trend across Period?

A
  • As the number of protons increases
  • The increasing nuclear charge attracts the valence electrons more strongly
79
Q

Ionization Energy trends going down the group?

A
  • As the number of filled shells increases
  • The valence elctrons are less strongly attracted
  • Due to greater shielding and distance from the nucleus
80
Q

Electronegativity trends across period?

A
  • As the number of protons increases
  • Increasing Nuclear charge attracts the electrons more strongly
81
Q

How does Electronegativity vary going down the group?

A

Greater Atomic Radius and shielding weakens the attractive force on the electron

82
Q

Amphoteric?

A

A chemical species that behaves as both an acid and a base

83
Q

Reactions of the Oxides period 3 with Water?

A

Na2O + H2O –> 2NaOH(aq)

MgO + H2O –> Mg(OH)2 (s)

Al2O3 + H2O –> no reaction

SiO2 + H2O –> no reaction

P4O6 + 6H2O –> 4H3PO3

SO2 + H2O –> H2SO3

Cl2O + H2O –> 2HClO

84
Q

Allotropes?

A

Allotropes are different molecular or crystalline structures of an element

85
Q

Graphite characteristics?

A

**Conductivity - High **

Only 3 electrons used in bonding - the fourth electron is found in a delocalised sea of electrons

Geometry - 2-dimensional

Hardness - Very Soft

2 dimension sheets slide over each other because of van der Waals’ forces between the sheets

86
Q

Fullerence characteristics?

A

Conductivity - Intermediate

Surface of the “ball” is conductive

Geometry

2 dimensonal folded into a sphere

Soft

87
Q

How do you determine if something is an London Forces?

A

Non-polar molecules

Any molecular element like Cl2, Br2, O2, P4, S8

Hydrocarbons

88
Q

How do you do you know something is a Dipole-Dipole IMF?

A
  • A covalent molecule that has different non-metals, not symmetrically oriented
89
Q

Physical properties of Solids

Giant Ionic

A
  • High Boiling Point
  • Low Volatility
  • High Delta H vap
90
Q

Giant Metallic Physical Properties?

A

High Boiling Point

Low Volatility

High Delta H vap

91
Q

Giant Covalent Physical Properties?

A

High Boiling Point

Low Volatility

High Delta H vap

92
Q

Molecular Covalent Physical Properties of Solids?

A

low boiling point

high volatility

low delta H vap

93
Q

Assumptions in Calorimetry?

A
  • Specific Heat Capacity of the solution is the same as pure water
  • Zinc in the reaction mixture did not absorb some of the energy
  • Density of the solution was 1.0g/cm3
  • The Calorimeter itself did not absorb any heat
  • Heat Loss
94
Q

Equilibrium definition?

A

Equilibrium is achieved when the forward rate of reaction and the reverse rate of reaction are equal

95
Q

How does the Removal of a Species changes its equilibrium constant?

A

Shifts towards that species to maintain its concentration

Usualy accomplished by adding something else that reacts with only that species

96
Q

Continuos Spectrum?

A

contains all wavelengths of electromagentic radiation

97
Q

Line spectrum?

A

contains only certain energies of electromagnetic radiation

98
Q

Line spectrum?

A

contains only certain energies of electromagnetic radiation

99
Q

Shapes of orbitals?

A

S - spherical

p - dumbell

d - 4-leaf clover

f - pointing to corners of a cube

100
Q

Electronic structure of the Periodic Table?

A

**Group 1 **

1s1 –> 5s1

**Group 2 **

2s2 –> 5s2

**Transition metals **

1st row: 3d1 –> 3d10
2nd row: 4d1 –> 4d10

Group 3:

2p1 –> 5p1

Group 4:

2p2 –> 5p2

Group 5:

2p3 –> 5p3

Group 6:

2p4 –> 5p4

Group 7:

2p5 –> 5p5

Group 8:

1s2 –> 5p6

101
Q

Where does oxidation and reduction take place in Voltaic cell electrolysis?

A

Anode - where oxidation occurs (negative)
Cathode - where Reduction occurs (positive)

102
Q

Voltaic Cell Drawing Check List?

A
  • electrodes labelled with metal type
  • external circuit
  • salt bridge
  • half reactions
  • anode/cathode labelled
  • electrode polarity labelled
  • electron flow in external circuit shown
103
Q

Purpose of salt bridge?

A
  • Salt bridge completes the circuit
  • allows for the conduction of ions
  • allows for the balancing of ionic charge
104
Q

Common mistakes made with the salt bridge?

A
  • Electrons flow through the wire in the external circuit
  • Ions flow through the salt bridge
105
Q

Dynamic Equilibrium?

A

The rate of the forward reaction is equal to the rate of the backwards reaction

106
Q

What characteristics do equilibrium reactions have?

A
  • Forward and Reverse reactions occur at the same rate (see graph)
  • The concentration of reactants and products are equal
  • Equilibrium requires a closed system
  • There is no change in macroscropic properties
107
Q

What does the value of Kc determine?

A

To which side the reaction has gone/

108
Q

Values of Kc and what they mean?

A

Forward reaction hardly proceeds - Kc < 10^-10

Reactants dominant at equilibrium - 0.01

Equal - 1

Products predominant at the reaction - 100

Reaction - almost goes to completion - Kc > 10^10

109
Q

If a reaction is reversed, what will the reciprocal of the reaction be equal to?

A

If the forward reaction value of Kc = 500, the reverse reaction would be equal to 1/500

110
Q

If the reaction is doubled then what happens to the Kc?

A

The Kc is also doubled

111
Q

Le Chatleiers Principle

A
  • When a system at equilibrium is subject to change, the system will react to minimise the effect of that change
112
Q

Which factors affect Kc?

A

ONLY temp

113
Q

Why do catalysts not effect the rate of reaction?

A

A catalyst increases the rate of the forward and backward reactions equally

tHEY INCREASE THE RATE OF REACTION WITHOUT BEING CHEMICALLY CHANGED as they offer a lower activation energy

114
Q

Functional Groups

A

SEE

115
Q

Factors affecting Volatility

A
  • Low volatility = High Boiling Point
  • The more polar the molecule
116
Q

Equilibrium Constant equation?

A

[D]^x/[A]^x [B]^x

117
Q

What side will the equilibrium shift if the temp is increased?

A

When temp increases –> Moves to endothermic side –> Moves Forward/Backwards (depending on the reaction)

118
Q

First Ionisation Eneegry

A

the energy required to remove one mole of the most loosely held electrons from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+

119
Q

What is the difference between Electron Domain and Molecular Geometry?

A

Electron domain geometry describes the arrangement of electron domains (either bonding pairs or lone pairs) around a central atom in a molecule, regardless of the positions of the atoms. In other words, electron domain geometry is based solely on the number and arrangement of electron pairs around the central atom, and it determines the overall shape of the molecule. The electron domain geometry is usually determined by counting the number of electron pairs around the central atom and using that to identify the molecular shape.

On the other hand, molecular geometry describes the actual arrangement of atoms in a molecule, taking into account the positions of the atoms and not just the electron pairs. Molecular geometry is also known as the “shape” of the molecule, and it is often used to explain the properties of the molecule, such as polarity, reactivity, and intermolecular forces. Molecular geometry is determined by the electron domain geometry and the positions of the atoms around the central atom.

120
Q

Collision Theory

A
  • Particles myst collide in their correct orientation
  • Particles require sufficient activation energy
  • Sufficient Energ y> Actyivation Energy
121
Q

Reactions of Electroolysis of Copper Bromide at the Anode and the Cathode?

A

2Br - + 2e- –> Br2
Cu –> Cu 2+ + 2e-

122
Q

Orbital Diagram of Scandium?

A

SeeS

123
Q

SPDF ON THE PERIODIC TABLE

A

SEE

124
Q

Electronic configuration of Co

and Co2+

A

1S2 2S2 2P6 3S2 3P6 3D7 4S2: ATOM

1S2 2S2 2P6 3S2 3P6 3D7 4S2: ION

125
Q

FUNCTIONAL GROUP OF ALKYNE

A
  • C (TRIPLE BOND) C -

(Alkynes)

CnH2n - 2

126
Q

General formula of Halogenalkanes?

A
  • Cn H2n+1X
127
Q

Formula group of Aldehydes?

A

R-CHO

(-al)

Cn H2n + 1 CHO

128
Q

Ketones Functional group?

A

R-CO-R

carbonyl (-one)

Cn H2n + 1 COC x H2x+1

129
Q

Ether Functional group?

A

R-O-R

(alkoxy)

Cn H2n+1O CmH2m+1

130
Q

Ester

A

R–COO–R’

ester

CnH2n+1CO2CmH2m+1

131
Q

Amines Functional group?

A

R–NH2

amine or amino

CnH2n+1NH2

132
Q

Amides

A

R–CONH2

carboxamide or amido

CnH2n+1CONH2

133
Q

Nitriles

A

R–CN

nitrile

CnH2n+1CN

–C≡N

134
Q

What does Volatility mean?

A

The measure of how easily a substance evaporates is referred to as its volatility

135
Q

What is the relationship between melting point and volatility?

A

The higher the volatility the lower the boiling point

136
Q

Factors affecting Volatility?

A
  • Boiling Point
  • Is it branched chained?
  • Shape of the molecule
137
Q

How does the nature of the functional group affect the volatility?

A
  • Polar functional groups –> Strong dipole-dipole interactions between the molecules and higher boliling points
138
Q

Functional groups in order of boiling point?

A

Alkanes, Halogenalkane, Aldehyde, Ketone, Amine, Alcohol, Carboxylic Acid

139
Q

Factors affecting the solubility of organic compounds?

A

Depends on their ability to form hydrogen bonds with water molecules

Compunds that are able to form hydrogen bonds in the water wiull be soluble

The first three alcohols - methanol, ethanol and propanol - are considered to be completely soluble in water.

As the length of the hydrocarbon chain increases, however, the solubility decreases. Alcohols with 10 or more carbon atoms, such as decanol (CH3(CH2)9OH), are considered to be insoluble in water.

This is due to the increasing size of the hydrocarbon chain, which increases the hydrophobic character of the molecule, decreasing its solubility in polar solvents such as water.

140
Q

How to classify primary, secondary and tertiary?

A
  • Alcohols and Halogenalkanes –> Refers to the nmber of carbon atoms that are bonded to the carbon atom with the Functional group
  • Amines –> Number of alkyl groups that are bonded to the Amine
141
Q

Why should you not innhale benzene?

A

Vapours are quite hazardous –> Benzen is a carcinogen

142
Q

Write the equations for the initian, propogation and termination of Free Radical substitution between Clhorine and Methane?

(- represents the dot)

A

Cl2 –> 2Cl’ (Initiation)

Propogation

CH4 + Cl’ –> ‘CH3 + HCl
‘CH3 + Cl2 –> CH3Cl + Cl’

Termination:
Cl’ + Cl’ –> Cl2

‘CH3 + ‘Cl –> CH3Cl

143
Q

Can you explain what happens in the initiatoon step of free-radical substituion?

A
  • Bond between Cl2 atoms breaks in the presence of UV light
  • Homolytic Bond fission –> forms free radicals
144
Q

Propogation

A
  • Chlorine free radical + Molecule of methane
  • Makes a chlorine free radical and Hydrogen Chloride
145
Q

What is the theoretical enthalpy change in combustion equation?

A

Sum of reactants - Sum of products

1- You take big (number of moles ) into account
2- Tink about each indivudal bond

146
Q

Why does Mg have a higher ionisation energy than Na?

A
  • Mg has a greater number of protons
  • greater attraction by Mg for valence electrons –> Lower Atomic Radius
147
Q

Why does Calcium have a greater melting point than K

A

Ca2+ > K+

  • greater attraction between delocalised e
148
Q

Why is the boiling point of Na2O > SO3?

A

Sodium oxide is ionic whereas SO3 is covalent

Na2O has SEFoA whereas SO3 has weak IMF’s

Forces more easily broken than in ionic bonds

149
Q

Definition of enthalpy bond change?

A

Enenrgy needed to break one mole of a bond in a gaseous molcule; averaged over similar compounds

150
Q

How soluble are Alchohol?

A
  • Alcohols are soluble in water as they can form H bonds in Water
151
Q

How soluble are Alchohol?

A
  • Alcohols are soluble in water as they can form H bonds in Water
152
Q

Partial oxidation of Primary Alcohols?

A

Catalyst - Potassium Dichromate

Aldehyde

If the aldehyde is the desired product, it can be removed as it forms by distillation

. This is possible because the boiling point of the aldehyde is lower than that of both the alcohol and the carboxylic acid

153
Q

Complete Oxidation of primary alcohols?

A

Carboxylic Acid

Conditions: heat under reflux + excess oxidising agent

If Carboxylic acid is wanted –> heat mixture

154
Q

Oxidation of Secondary Alcohols

A

Only produces 1 product for both partial and complete oxidation

Ketone

Conditions: Potassium Dichromate and heat under reflex

155
Q

Oxidation of tetiary alcohols

A

Tertiary alcohols have no hydrogen atoms bonded to the carbon atom that is directly bonded to the hydroxyl group. Therefore, they do not undergo oxidation reactions

156
Q

Examples of giant covalent structures?

A

Diamond, graphite Silicon, Silicon Dioxide

157
Q

Examples of Molecular covalent structures?

A

H2O, CO2, C6H12O6

158
Q

aRE TWO IDENTICAL NON-METALS POLAR?

A

No, two non-polar covalent bonds for two identical molecules (also for C-H)

159
Q

Scale for electronegativity?

A

0 - 0.4 (Pure covalent)
0.4-1.7 (Polar covalent)
1.7 - xyz (Iionic)

160
Q

Ncl3 VESPR

A

4 Electron Domain

Electron Domain Geometry Tetrahedral

3 Bond Pairs and 1 lone pairs

Molecular Geometry (no lone pairs) Trigonal Pyramid

Bond Angle - 107

161
Q

Cl2O VESPR

A

4 Electron Domain –> Electron Domain Geometry Tetrahedral

Bond pairs - 2

Lone Pairs - 2

Molecular Geometry - Bent

Bond Angle - 104.5

162
Q

Properties of Diamond

A

Conductivity: None - all electrons are used in localised bonds

Geometry: 3-dimensional

Hardness: Very hard - all atoms are held together in the lattice by strong covalent bonds

163
Q

Properties of Graphite

A

Conductivity - High only 3 electrons used in bonding - fourth found in a delocalised electron sea

Geometry - 2 dimensional

Hardness - Very soft - 2 dimeonsional sheets slide over each other because of weak van der Waals’ forces between the sheets

164
Q

Properties of Fullerene

A

Conductivity - No

Hardness - Soft imagine the molecular form of the ball room at Mcdonalds

165
Q

does sulfur trioxide have a simple molecular structure

A

Sulfur trioxide (SO3) has a simple molecular structure.

166
Q

What does the IHD mean?

A
  • Gives the number of molecules of the H2 needed to convert organic compounds to the corresponding saturated, non-cyclic molecule
167
Q

What does the IHD mean?

A
  • Gives the number of molecules of the H2 needed to convert organic compounds to the corresponding saturated, non-cyclic molecule
168
Q

IHD equation

A

IHD = 2X + 2-Y/2

169
Q

rules for IHD

A

O2 + Sulfur atoms are not counted

Halogen atoms are treated like H atoms

For each N atom –> add 1 C + 1 H

Each double bond = 1IHD

One Benzen Ring = 1

170
Q

Magnesium Oxide

Dot and cross for ionic compounds

A

SEE

171
Q

Maxwell BoltzmannCurve

A

see

172
Q

If you ae finding the IHD of a coumpound and you know the molecular formula, check the mr table and see if it corresponds with the exact mr value

A

If you ae finding the IHD of a coumpound and you know the molecular formula, check the mr table and see if it corresponds with the exact mr value

173
Q

How to calculate IHD

A

SEE

174
Q

Outline the concept of relative abundance.

A

Relative abundance is the percent abundance in nature of a particular isotope.
✓
✓

175
Q

What are the following requirements for a reaction to occur?

A
  • Correct orientation
  • Reactant particles must collide with sufficient energy –> Ea needed
176
Q

Diagram of Activation Energy on the effect of the Rate of Reaction

A

See

177
Q

Draw a Maxwell - Boltzmann Distrbution graph

A

SEE

178
Q

Draw a Maxwell - Boltzmann Distribution diagram when there is an increase in temperature?

A

SEE

179
Q

Describe how a catalyst increases the rate of reaction

A
  • Provides an alternative pathway
  • Lower Ae
  • Greater proportion of particles now have energy equal to, or greater, than the og AE
  • Greater frequency of succesful collisions
180
Q

What is the pH of unpolluted H2O

A

5.6

181
Q

What is the pH of acidic H2O

A

5.0

182
Q

What are the two types of Acid Deposition?

A

Wet Deposition and Dry Deposition

Wet - Acid rain, snow and fog
Dry - Acid, gases and solid particles

183
Q

How are SO2 and Oxides of Nitrogen formed?

A

1.) Lightnings
2.) Internal combustion in engines
3.) Volcanic Eruptions

184
Q

Draw the Equilibrium constant of

2NH3(g) + 2O2(g) ⇌ N2O(g) + 3H2O(g)

A

Squared always goes along the bottom

IB Chemistry (18 decks)

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