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Chemistry > Topic 2 > Flashcards

Topic 2 Flashcards

1
Q

Relative atomic mass (Ar)

A

Weighted average mass of an atom compared to 1/12 the mass of an atom of carbon-12.

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2
Q

Relative formula mass (Mr)

A

The relative formula (or molecular) mass (Mr) of a compound is the weighted average mass of the compound compared to 1/12 the mass of an atom of carbon-12

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3
Q

Molar Mass

A

The molar mass (M) is the mass in grams of one mole of a substance. The unit for molar mass is grams per mole (g mol-1).

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4
Q

Difference between molar mass and relative atomic mass

A

The molar mass of a substance is numerically equal to its relative atomic mass (or relative formula / molecular mass), however, it is important to remember that molar mass has a unit (g mol-1) whereas relative atomic mass is a dimensionless quantity. For example, carbon has a relative atomic mass of 12.01; its molar mass is 12.01 g mol-1. Water (H2O) has a relative molecular mass of 18.02; its molar mass is 18.02 g mol-1

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5
Q

Empirical formula

A

An empirical formula is the lowest whole number ratio of atoms (or ions) in a compound. A molecular formula is the actual number of atoms in a compound.

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6
Q

Describe an experimental method to find empirical formula for magnesium oxide?

A

The first two methods involve the heating of a substance and measuring changes in mass. In the first example, a sample of magnesium is heated in a crucible until it reacts with the oxygen in the air using the apparatus in Figure 3. The crucible is opened periodically to allow oxygen to enter, however, care must be taken not to allow any product (a white powder) to escape.

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7
Q

Possible errors in the magnesium oxide empirical method?

A

The magnesium used was not pure

The product was something other than magnesium oxide (such as magnesium nitride, Mg3N2)

Some product was lost when the lid was removed to allow oxygen in

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8
Q

Definition of water crystallisation

A

The water of crystallisation is the number of water molecules present in one formula unit of a salt.

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9
Q

How can the water of crystallisation be determined?

A

The water of crystallisation can be determined experimentally by heating a hydrated salt until the water evaporates, leaving the anhydrous salt. By subtracting the mass of the anhydrous salt from the mass of the hydrated salt, the mass of water evaporated can be determined

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10
Q

Water crystallisation steps

A
  1. Determine the mass of water, H2O, evaporated from the hydrated salt
  2. Determine the amount (in mol) of H2O evaporated using the water mass calculated in step 1 and the mol of the remaining Na2CO3
  3. Divide each amount by the smallest
  4. Find lowest whole number ratio
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11
Q

What kind of mass is water heated to in water of crystallisation?

A

Constant mass - this involves heating the sample multiple times until a constant mass is obtained.

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12
Q

Assumptions made in the water of crystallisation method?

A

The mass change is only due to the water lost from the hydrated salt

The crucible does not absorb water

The anhydrous salt does not decompose

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13
Q

What is the other way of determining empirical formula

A

The final method of determining the empirical formula of a compound involves burning a sample of the substance in an excess of oxygen (complete combustion) and analysing the products of the combustion. Once the products of the combustion reaction have been collected and weighed, the empirical formula can be determined. Note that this technique usually involves the combustion of an organic compound containing the elements carbon, hydrogen and oxygen.

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14
Q

Determine the mass of carbon in 0.366 grams of CO2

A

0.366 x 12/44.01

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15
Q

Combustion empirical formula method steps?

A

Find moles of carbon and hydrogen - remember the moles of the carbon is the same as the moles of carbon dioxide due to the 1:1 ratio but moles of hydrogen in H20 would be twice the moles of water.

Find mass of oxygen by using percentage composition to first find CARBON and HYDROGEN mass

Subtract the mass of carbon + hydrogen from total amount combusted to find mass of oxygen

Now you know all masses, do the empirical table for carbon, hydrogen and oxygen

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16
Q

Volume = mol x 22.4

A

This equation helps you to convert between volume and mole

It also shows that 1 mole of ANY gas occupies 22.7 dm3 of volume at STP which is 273 K and 100 Kpa

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17
Q

If % yield is lower, give reasons for this?

A
  • Side reactions
  • Product lost during transferring during containers
  • gas produced
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18
Q

If % yield is higher, give reasons?

A
  • contamination

- water not dried

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19
Q

What does the ideal gas equation assume

A
  • assumes 0 volume of individual atoms and molecules gas, NOT the entire volume
  • elastic collisions, don’t stick to each other just bounce off
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20
Q

Equation linking pressure, volume and temperature

A

P1 x V1/ T1 = P2 x V2/ T2

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21
Q

Equation linking pressure, volume and temperature

A

P1 x V1/ T1 = P2 x V2/ T2

ASSUMES FIXED MASS

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22
Q

Describe the errors in the method to find molar mass of CO2 in cuco3 (4)

A

Gas collected may not be pure co2

Co2 may be soluble in water

Air in the tube is collected with the gas

Gas collected has not equilibrated to room temperature

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23
Q

Describe the errors in the method to find molar mass of CO2 in cuco3 (4)

Good
Cats
Admire
Goats

A

Gas collected may not be pure co2

Co2 may be soluble in water

Air in the tube is collected with the gas

Gas collected has not equilibrated to room temperature

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24
Q

How does decreasing the volume affect the pressure?

A

It increases the pressure as the concentration of the particles increases, resulting in more collisions with the container walls.

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25
Q

How does increasing the temperature affect the volume?

A

As the temperature increases, the average kinetic energy of the particles increases so the force with which the particles collide with the container walls increases. Hence pressure increases and is directly proportional to absolute temperature

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26
Q

Difference between real gases and ideal gases

A

Real gases have some intermolecular forces between them

Do occupy some space (their particles)

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27
Q

Gas constant units

A

JK^-1

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28
Q

What 2 things does the deviation from ideal gas behaviour graph show

A
  • challenges the fact that volume of particles is negligible

- challenges the no attractive forces

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29
Q

Mol dm3 to Gdm^3 - how do you do it?

A

Multiply MOL DM3 by Mr of that compound

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30
Q

Describe what happens to the number of moles and concentration when a solution is diluted

A

Number of moles stays the same

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31
Q

C1 x V1 = C2 x V2

A

initial and final - because number of moles stays the same

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32
Q

Describe the composition of an atom

A

Positive nucleus surrounded by negatively charged electrons

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33
Q

When are emission spectra produced

A

Emission spectra are produced when photons are emitted from atoms as excited electrons return to a lower energy level.

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34
Q

Why do the particles need to be ionised before placing them in a mass spectrometer?

A

So that they gain a positive charge and can be attracted to the negative plates and be accelerated

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35
Q

What types of ions are deflected the least

A

Ions with a higher mass to charge ratio are deflected less in the magnetic field than ions with a lower mass to charge ratio.

Higher mass - lower velocity - kinetic energy formula

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36
Q

Mass spectrum X and Y axis labels?

A

X axis = mass/charge ratio

Y axis = relative abundance %

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37
Q

Describe the bohr model of the atom

A

According to the Bohr model of the atom, electrons exist in energy levels, or principal energy levels

with n = 1 being closest to the nucleus and of lowest energy.

The further the energy level is from the nucleus, the higher its number (n) and the higher its ENERGY

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38
Q

What sub-levels exist in the n=1 energy level

A

S orbital

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39
Q

What sub-levels exist in the n=2 energy level

A

S and P orbital

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40
Q

What sub-levels exist in the n=3 energy level

A

S, P, D

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41
Q

What sub-levels exist in the n=4 energy level

A

S, P, D, F

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42
Q

How many electrons are in first energy level - n=1

A

2

43
Q

How many electrons are in 2nd energy level

A

2 + 6 = 8

44
Q

How many electrons are in 3rd energy level

A

2 + 6 + 10 = 18

45
Q

Electrons in n=4?

A

2 + 6 + 10 + 14 = 32

46
Q

How many P-orbitals are there

A

3

47
Q

How many S -orbitals

A

1

48
Q

How many D-orbitals

A

5

49
Q

How many F-orbitals

A

7

50
Q

Pauli’s exclusion principle

A

Two electrons can only occupy the same atomic orbital if they have opposite spins to reduce mutual repulsion.

One electron spins in one direction and the other electron spins in the other direction

51
Q

Afbau principle

A

Electrons fill atomic orbitals of lowest energy first.

The 1s sub-level has the lowest energy, therefore, it is filled first.

Within a given main energy level, s orbitals are of lower energy than p orbitals and therefore fill first.

The atomic orbitals within a sub-level are of equal energy (known as degenerate orbitals). This includes the three p orbitals in the 2p, 3p and 4p sub-levels and the five d orbitals in the 3d sub-level.

There is an overlap between the 3d and 4s sub-levels as energy levels get closer together further from the nucleus. This means that the 4s sub-level is of lower energy and fills before the 3d sub-level.

52
Q

Hund’s rule

A

Electrons fill orbitals in the same sub-level singly, before pairing up to reduce mutual repulsion

53
Q

Explain the formation of a line spectrum

A

According to the Bohr model of the atom, electrons can only exist at certain energy levels, which we call main energy levels or shells. By absorbing or emitting energy, electrons can transition between the energy levels. If a high voltage is passed through a gas, the electrons in the gaseous atoms become excited and transition to higher energy levels. As the electrons fall back down to lower energy levels, the transitions are accompanied by an emission of energy as a photon of light is emitted. This results in the formation of an emission line spectrum.

54
Q

Visible light spectrum?

A

In the case of a visible light emission line spectrum (Figure 6), the electrons transition from higher energy levels to the second main energy level (n = 2)

55
Q

What transition causes red light to be emitted in the visible spectrum

A

Transition from n = 3 to n = 2 emits energy that corresponds to red light

56
Q

What transition causes violet light to be emitted in the visible spectrum

A

Electron transition from n = 6 to n = 2 emits energy that corresponds to violet light and a violet line

57
Q

Which colour in the visible light spectrum has the greatest amount of energy

A

Violet light as it has a shorter wavelength

58
Q

What does the line convergence show

A

Note that the lines get closer together, or converge, as we move to the high-energy end of the spectrum

59
Q

Relationship between energy and frequency equation

A

E = hf

60
Q

Transition to n=1 energy level releases what?

A

UV

61
Q

Transition to n=2 energy level releases what?

A

Visible light

62
Q

Transition to n=3 energy releases what

A

Infrared - electron transition emits energy that corresponds to the wavelength or frequency of infrared (IR) radiation.

63
Q

How does line spectra give evidence for energy levels.

A

We have seen that the line spectra are at characteristic frequencies for a given element. This shows that a given atom only emits certain energies, and the conclusion is that there are only certain energy levels available for the electrons in the atom. Electrons can transition from one energy level to another, but not to somewhere ‘in between’. Examination of line spectra gives evidence for the model of the atom in which electrons are arranged in different energy levels.

64
Q

The energy of the convergence limit of the ultraviolet series corresponds to what?

A

The energy of the convergence limit of the ultraviolet series corresponds to the energy absorbed when an electron transitions from the n = 1 energy level to the n = ∞ energy level (Figure 2). Electrons in the n = ∞ energy level no longer experience the electrostatic attraction from the nucleus. At this point, the electron can be considered to have left the atom, resulting in the formation of a positive ion. This energy is the ionisation energy

65
Q

Equation to calculate ionisation energy

A

E = hv

E is energy in joules, h is Planck’s constant, 6.63 × 10−34 J s, and v is the frequency, in s−1

The units for E is joules and so the answer you get is in Joules. You MUST convert it into KJ mol- if it says in the question or even otherwise.

Multiply by 6.02 x 10^23

Divide answer by 1000

66
Q

Ground state is which energy level?

A

n = 1

67
Q

Are ionisation energies endo or exothermic

A

Endothermic because energy required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous 1+ ions.

68
Q

First and second ionisation energy equation example

A

Ca (g) → Ca+ (g) + e−

Ca+ (g) → Ca2+ (g) + e−

69
Q

Why do successive ionisation energies increase

A

We should expect the ionisation energies to increase progressively as we are removing negatively-charged electrons from increasingly positive ions which results in a stronger electrostatic attraction between the nucleus and the remaining electrons. More energy is needed therefore, to overcome this stronger attraction

70
Q

What do the large jumps in ionisation energy graphs indicate

A

Electron being removed from inner energy levels

71
Q

Exceptions to the afbau principle

A

chromium (Cr) and copper (Cu)

72
Q

Difference between continuum spectra and line spectra

A

A continuous spectrum shows all the wavelengths or frequencies of visible light. An absorption line spectrum differs from a continuous spectrum in that some of the wavelengths of visible light are missing, shown by the black lines on the coloured background. An emission line spectrum is characterised by having coloured lines on a black background.

73
Q

Explain why the melting points of the group 1 metals (Li → Cs) decrease down the group

A

atomic/ionic radius increases

smaller charge density

OR

force of attraction between metal ions and delocalised electrons decreases

metallic bond weakens

DO NOT accept discussion of attraction between valence electrons and
nucleus for M2.

74
Q

Reaction of P4O10 with water equation?

A

P4O10 (s) + 6H2O (l) → 4H3PO4

75
Q

Why does calcium have a greater density than potassium

A

Calcium has a smaller atomic radius
Stronger electrostatic force of attraction to nucleus
Stronger metallic bonding

76
Q

Describe the nature of orbitals

A

• Each orbital has a defined energy state for a given electronic configuration
and chemical environment and can hold two electrons of opposite spin.

77
Q

What are sub-levels

A

Sub-levels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron.

78
Q

• In an emission spectrum, the limit of convergence at higher frequency corresponds to what?

A

First ionization energy

79
Q

In an emission spectrum, the limit of convergence at higher frequency corresponds to what?

A

First ionization energy

80
Q

Explain what happens to atomic radius across a period

A

Atomic radius decreases

electrons added to same energy level

so no change in amount of shielding

Increased number of protons increases nuclear charge
stronger

Stronger electrostatic force of attraction

81
Q

Explain what happens to atomic radius across a period

A

Atomic radius decreases

electrons added to same energy level

so no change in amount of shielding

Increased number of protons increases nuclear charge
stronger

Stronger electrostatic force of attraction between nucleus and outer electron

82
Q

Describe the trend in the reactivity of group 1 elements down the group

A

Increases

atomic radius increases

shielding increases

greater number of energy levels

outer electron further away from nucleus

Weaker electrostatic force of attraction

Outer electron easily removed

83
Q

Compare the size of a cation to the parent atom

A 
Cations contain fewer electrons than
protons so the electrostatic attraction
between the nucleus and the outermost
electron is greater and the ion is smaller
than the parent atom.
84
Q

Compare the size of a cation to the parent atom

A

Cations contain fewer electrons than
protons so the electrostatic attraction
between the nucleus and the outermost
electron is greater and the ion is smaller
than the parent atom. Across the period the
ions contain the same number of electrons
( isoelectronic) , but an increasing number
of protons, so the ionic radius decreases.

85
Q

Compare the size of an anion to the parent atom

A

Anions contain more electrons than
protons so are larger than the parent atom.
Across a period the size decreases because
the number of electrons remains the same
but the number of protons increases.

86
Q

Discuss trend in melting points

A

At the left of the period, elements exhibit metallic bonding, which increases in strength across period as atomic radius decreases, charge density increases, metallic bonding gets stronger.

Carbon and silicon have high melting points in their respective periods because it is a macromolecular structure with strong bonds which take a lot of energy to overcome

Elements in group 15, 16, 17 e.g: P4, S4 and Cl2 are simple molecular structures with weak intermolecular forces

Noble gases exist as monoatomic molecules with extremely weak forces of attraction between the atoms

87
Q

Why does melting point increase down group 7?

A

Atoms get larger due to more energy shells

London dispersion forces increase between diatomic molecules

More energy needed to overcome

88
Q

Explain the decrease in ionisation energy of sulphur and oxygen

A

Electrons pair up in one of the d-orbitals

Increased repulsion

makes it easier to remove compared to if the electron was unpaired

89
Q

Why does ionisation energy slightly increase from beryllium to boron

A

Boron is in 2p sub-level which is at a higher energy level, further away from nucleus compared to beryllium in 2s sub-level which is closer to the nucleus

90
Q

Define electron affinity

A

Energy needed to add one mole of electrons to one mole of gaseous atoms to form one mole of ions

91
Q

Define electronegativity

A

The electrostatic force of attraction of the nucleus of an atom to a bonding pair of electrons

92
Q

Elements with high electronegativity form what type of bond

A

Ionic

93
Q

Describe trends in electronegativity across a period

A

Increases as atomic radius decreases

nuclear charge increases

same shielding

stronger electrostatic force of attraction between nucleus and outer electron

94
Q

Write the equations for lithium, sodium and potassium with water

A

2Li (s) + 2H2O (l) -> 2LiOH + H2 (g)

2Na (s) + 2H2O (l) -> 2NaOH + H2 (g)

95
Q

Test for halide ions?

A

Add silver nitrate solution

Silver ions react with halide to form silver halide

AgCl - white

AgBr - cream

AgI - yellow

96
Q

Metal oxides + water : write equations for:

Na2O + h20

MgO

SO3

P4O10

NO2

A

Na2O (s) + H2O (l) - 2NaOH (aq)

MgO (s) + H2O (l) - Mg(OH)2

SO3 (g) + H2O (l) -> H2SO4 (aq)

P4O10 (s) + 6H2O (l) -> 4H3PO4 (aq)

3NO2 (g) + H2O (l) -> 2HNO3 (aq) + NO (g)

97
Q

Define a transition element

A

An element that possesses an incomplete d sublevel in one or more of its oxidation states

98
Q

Why is zinc not a transition metal

A

It contains a full d-sub level in all its oxidation states

99
Q

Why is scandium not a typical transition metal

A

Its common ion Sc3+ has no d-electrons

100
Q

Describe the properties of transition metals

A

Variable oxidation states

Form complex ions with ligands

Have coloured compounds

Display catalytic and magnetic properties

All transition metals show an oxidation state of 2+ as 4s electrons are lost first

101
Q

Examples of transition metals as catalysts

A

Vanadium (v) oxide in the contact process

Iron in haber process

Nickel in hydrogenation process

102
Q

What is ferromagnetism?

A

Permanent magnetism - magnet stays a magnet because unpaired electrons align parallel to each other irrespective of whether an external magnetic field is present. Used in COMPASSES to point north

103
Q

What is paramagnetism?

A

When spinning unpaired electrons create a small magnetic field and line up in an applied magnetic field to make the transition metal weakly magnetic. The more unpaired electrons there are, the stronger the paramagnetism

104
Q

Diamagnetism?

A

When all electrons are paired - not magnetic

Chemistry (7 decks)

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