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A Level Chemistry (OCR B Salters) > Elements of Life > Flashcards

Elements of Life Flashcards

1
Q

What is the atomic number?

A

The number of protons in the nucleus of an atom

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2
Q

What is the mass number?

A

The total number of protons and nuetrons in the nucleus of an atom

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3
Q

What is relative atomic mass (Ar)?

A

The mass of one atom of an element relative to 1/12 the mass of carbon-12

Is an average of relative isotopic masses, taking into account abundance

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4
Q

What do the different numbers on the nuclear symbol (of an element etc.) tell you?

A

No. protons = Atomic no. (Bottom no.)

No. electrons = Atomic no. (Bottom no.)

No. neutrons = Mass no. - atomic no.
(Top no. - bottom no.)

Mass no. (Ar) = Top no.

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5
Q

How are models of the atom made + updated?

A

Tested using experimental investigations

Are revised when observations are made that aren’t predicted by model

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6
Q

What were the different steps/models in the development of the current atomic model?

A
  • Dalton model - simple ‘billiard ball’. Particles cannot be divided, created, or destroyed. Are unique.
  • ‘Plum pudding’ model - electrons embedded in sea of positive charge. Discovered by firing cathode rays (electrons) in air - discovered e-. Introduced idea that atoms made of smaller particles
  • Nuclear model - Geiger-Marsden experiment showed some alpha particles deflected at large angles by small, dense area of positive charge surrounded by e-
  • Bohr model - Evidence from atomic spectra + patterns of ionisation enthaply. e- arranged in shells - ‘planetry model’
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7
Q

What is nuclear fission?

A

The splitting of a large, unstable isotope triggered by bombarding it with smaller, high-speed particles (usually neutrons)

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8
Q

What conditions are needed for nuclear fission?

Why?

A

High temps and/or pressure to provide the energy needed to overcome the repulsion between the 2 positive nuclei

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9
Q

What is the nuclear symbol for a neutron?

A

10n

Except with the 1 and 0 in line above each other…

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10
Q

Write a nuclear equation for the fusion of a 11H nucleus with a 21H nucleus

A

21H + 31H → 42He + 10n

  • Ignore the bit in the picture about what reaction it is…*
  • Also neutrons only need to be included if numbers don’t add up*
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11
Q

What is the general formula for calculating Ar?

A

(% abundance of x X isotopic mass of x) + (% abundance of y X isotopic mass of y)

/ 100%

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12
Q

The Ar of potassium is 39.1. Calculate the relative abundance of 39K and 41K

A

Make 1 isotope abundance x so the other = 100-x
(39x + 41(100-x))/100% = 39.1

Multiply both sides by 100 + multiply out brackets
39x + 4100 - 41x = 3910
-2x = -190, x =95
39K = 95%, 40K = 5%

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13
Q

What are isotopes?

A

Atoms of the same element with a different number of neutrons

This causes mass number to be different

Their relative abundances are used to calculate Ar

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14
Q

What is Mr?

A

Relative molecular mass

The ratio of the average mass of one molecule of an element or compound to one twelfth of the mass of an atom of carbon-12.

(Ar but for molecules… (not elements))

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15
Q

What is the Avogadro constant (NA)?

A

The number of atoms/molecules in 1 mole of a substance

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16
Q

What does quantised mean?

A

Energy that can only take particular values (known as quanta)

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17
Q

What is the ground state?

A

The lowest energy level that an electron can occupy

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18
Q

What is a photon?

A

Quanta of energy in the form of electromagnetic radiation

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19
Q

Breifly describe Bohr’s model of the atom

A

Electrons in an atom occupy discrete, quantised energy levels/shells

Electrons in an energy level have a specific amount of energy

Hence the energy of the electron is said to quantised

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20
Q

What property does light have?

What does this mean?

A

Wave-particle duality

Means it can behave like both a wave and a particle…

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21
Q

What properties does light have the mean it can be described as a particle?

A

Made up of ‘tiny packets of energy’ called photons

The energy of a photon corresponds to its position in the EM spectrum

Increased freq. = increased energy + decreased wavelength

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22
Q

What equation links the wave + particle models of light?

A

ΔE = hv

ΔE = energy of photon (J)
h = Planck's Constant
v = frequence (Hz/s<sup>-1</sup>)
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23
Q

What equation expalins the wave properties of light?

A

c = vλ

c= speed of light (ms<sup>-1</sup>)
v = frequency (Hz/s<sup>-1</sup>)
λ = wavelength (m)
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24
Q

Describe the appearance of an emission spectrum

A

Consists of coloured lines on a black background

The lines become closer at higher frequencies

There are several series of lines (although some may fall outside visible part of spectrum)

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25
Q

What is spectroscopy?

A

The study of how light and matter interact

Uses IR, visible, and UV light

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26
Q

Explain the formation of an emission spectrum

A
  • Electrons in the ground state absorb energy
  • This promotes them to a higher energy level - excited state
  • Electrons then drop back down to lower energy levels. The energy lost (ΔE) us emitted as a photon of light
  • The frequency of the photon is related to the energy lost by ΔE = hv
  • Different energy gaps produce photons of different frequencies
  • This produces different coloured bands on the emission spectrum
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27
Q

Why can emission/absorption spectra be used to identify different atoms from a compound/mixture?

A

Because each element has a unique configuration of electrons, therefore has a unqiue emission/absorption spectrum

The energy levels of the electrons are discrete + quantised means only certain freqs. emitted/absorbed - it’s not continuous

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28
Q

What are flame tests?

A

Used to identify the presence of specific metals in a sample

Different metals give different coloured flames depending on their emission spectra

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29
Q

What is flame colour?

A

The light emitted by metal ions when a vaporised metal salt is heated up in a flame

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30
Q

What colour flame does Li+ give?

A

Bright red

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31
Q

What colour flame does Na+ give?

A

Yellow

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32
Q

What colour flame does K+ give?

A

Lilac

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33
Q

What colour flame does Ca2+ give?

A

Brick red

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34
Q

What colour flame does Ba2+ give?

A

Apple green

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35
Q

What colour flame does Cu2+ give?

A

Blue-green

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36
Q

Describe the appearance of an absorption spectrum

A

If white light is passed through a sample of vaporised atoms, an absorption spectrum is seen

Shown by black lines on a rainbow background (showing all colours of visible light)

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37
Q

How are atomic absorption spectra formed?

A
  • Electrons in the ground state absorb photons of light
  • The energy from these photons causes the electrons to be excited to higher energy levels
  • The electrons drop back down to the ground state and a photon/light is emitted
  • The energy of this photon is related to the frequency/energy of light initally absorbed as ΔE = hv
  • Light of the frequency doesn’t pass through the sample (as it’s absorbed) so a black line is seen in the spectrum
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38
Q

What are the similarities between emission and absorption spectra?

A

For a given element, lines appear at the same frequency

Lines converge at a higher frequency

Several series of lines are seen

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39
Q

What are the differences between atomic emission and absorption spectra?

A

Emission spectra show coloured lines on a black background

Absorption spectra show black lines on a coloured background

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40
Q

Why do the lines of emission/absorption spectra get closer together at higher frequencies?

A

Higher frequency lines are caused by translations of electrons with large ΔE values

These are produced from translations from higher energy levels

Higher energy levels are much closer together than lower energy levels

Translations from adjacent energy levels will have similar ΔE values and hence produce light of similar frequencies

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41
Q

Why are several series of lines seen on emission/absorption spectra?

A

Lines are produced when electrons drop to a lower energy level

Different series of lines are produced by electrons dropping to different ground states/electron energy levels

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42
Q

What is the principle quantum number?

A

Shell

Given as n (i.e. 1,2,3 etc. the number before the letter…)

The higher the value, the higher the energy

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43
Q

What are shell divided in to?

A

Sub-shells

Labelled s, p, d, and f

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44
Q

What is each sub-shell divided into?

What are its properties?

A

Atomic orbitals

Each can hold max of 2 electrons

These electrons must have opopsite (or paired) spinds

Represented by boxes. Arrows drawn in them represent electrons

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45
Q

How many orbitals does the s sub-shell contain?

A

1 s-orbital

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46
Q

Summarise the way in which electrons are organised in atoms, starting with the largest grouping.

A

Shell/PQN

Sub-shells

Atomic orbitals

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47
Q

Draw the different shapes of the p orbital

A

px-orbital

py-orbital

pz-orbital

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48
Q

How many orbitals does the p sub-shell have?

A

3

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49
Q

How many orbitals does the d sub-shell have?

A

5

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50
Q

How many orbitals does the f sub-shell have?

A

7

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51
Q

What are the rules that determine the distribution of electrons in atomic orbitals?

A

The orbitals are filled in order of increasing energy

Where there is more than one orbital at the same energy, the orbitals are first occupied by a single electron. When each orbital is singly occupied, the electrons pair up in the orbitals

Electrons in singly occupied orbitals have parallel spins

Electrons in doubly occupied orbitals have opposite (paired) spins

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52
Q

What are the 2 ways of representing electron distribution?

A

By writing out the electronic configuration in full
e.g. 1s22s22p5

By drawing the electronic configuration in boxes…
(see picture)

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53
Q

Draw a diagram to show the shape of the s-orbital

A

It’s a ball shape

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54
Q

On the periodic table what is a period?

A

A row

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55
Q

What is periodicity?

A

The occurence of a regular pattern in a property as you go across a period

The regular pattern is also repeated in other periods

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56
Q

How are elements in the periodic table arranged?

How did they used to be arranged?

A

Arranged by atomic number (no. protons)

Used to be arranged by Ar

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57
Q

What trend do melting/boiling points follow across a period?

(e.g. period 3)

A

Melting point increases then decreases across the period

This is because the metals on the left-hand side of a period are metalically bonded so have higher melting points due to the deloclaised electrons between nuclei. The further across the period, the more electrons and the more positive the nucleus becomes, so the stronger the bonds.

Silicone has a high melting point because it is a giant covalent structure which requires a lot of energy to break

The remaining non-metals are simple molecules. They are only held together by weak intermolecular forces (e.g. id-id). The melt these molecule you don’t need to break the strong covalent bonds, only the weak intermolecular bonds.

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58
Q

What is first ionisation enthalpy?

A

The energy needed to remove one electron from each atom in one mole of isolated gaseous atoms of an element

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59
Q

What is the general equation for first ionisation enthalpy?

A

X(g) → X+(g) + e-

Remember state symbols!

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60
Q

What is the trend in first ionisation enthalpies as you go across a period?

Why?

A

First ionisation enthalpy increases across a period.

Group 0 elements have the highest values because they have full outer shells, making it difficult to remove an electron

First ionisation enthalpy is lowest for Group 1 elements because they have only 1 outer shell electron which is relatively easy to remove/ionise

First ionisation enthalpy increases across a period because the number of protons in the nucleus increases, meaning the electrons are more strongly attracted to the nucleus so are harder to remove

61
Q

What is the trend in first ionisation enthalpy down a group?

Why?

A

First ionisation enthalpy decreases down a group

This is because electrons are in shells that are further away from the nucleus, thus the attraction between the two is less (electron shielding)

62
Q

What is the trend for atomic radii across a period?

A

Decreases due to the increased number of protons

This means there is greater attraction between the outer electrons and the nucleus

63
Q

What is a closed shell?

A

When all sub-shells are fully occupied by electrons

64
Q

Why are s-block elements more reactive than p-block elements?

A

Because the formation of M+ or M2+ ions only requires input of energy equivalent to the first/second ionisation enthalpy.

For p-block elements greater input of energy is needed to lose eletrons due to the greater electron affinity as a result of a more positive nucleus

65
Q

What is a dative covalent bond?

A

A type of covalent bond in which both electrons come from the same atom

Show by arrow pointing away from donor

66
Q

What is a lone pair?

A

A pair of electrons in the outer shell of an atom that are not involved in bonding

67
Q

What is ionic bonding?

A

Bond formed between metal + non-metal atom

Metal transfers/donates electron(s) to non-metal atom

This results in formation of charged ions, often with full outer shells. This makes them particularly stable

68
Q

Draw a dot-and-cross diagram to show the formation of NaCl from Na and Cl

A

Note: You only need to show outer shells + don’t have to draw circles

(Also don’t need the arrow, should be written like an equation)

69
Q

How are individual ions held together to form ionic compounds?

A

Cations + anions produced by ionic bonding held together by electrostatic attraction between each other

Results in the formation of a giant ionic lattice

70
Q

What is covalent bonding?

A

Bonding that occurs between 2 non-metal atoms

Formed by the atoms sharing one or more pairs of electrons

(If 2 pairs shared, double bond formed, etc.)

71
Q

Draw diagrams to show the different ways in which dative covalent bonds can be represented

A

Can be represented by:

Dot-and-cross diagrams

Arrow (in structural formulae)

72
Q

How are the atoms in a simle covalent bonds?

A

There is an electrostatic attraction between the positive nuclei of the 2 atoms and the shared pair of negative electrons in the bond

This is greater than the repulsion between the 2 nuclei

73
Q

What is electron pair repulsion theory?

(AKA VSEPR - Valance Shell Electron Pair Repulsion)

A

States that the shape adopted by a simple molecule is that which keeps repulsive forces to a minimum.

All bond angles must add up to = 360º

74
Q

Describe/explain how electron pair repulsion determines the shape of molecules

A

Electron pairs/groups repel each other

They will arrange themselves to get as far apart as possible

State the no. total pairs of electrons

State the no. bondind pairs/groups of electrons

(if applicable) state the no. lone pairs

(if applicable) lone pairs repel more than bonding pairs (decrease bond angle by 2.5º each)

This creates the shape […] with the angle(s) […]

75
Q

When describing/explaining electron pair repulsion in double/triple bonds, how should you describe electrons?

A

As groups not pairs

76
Q

Give the shape and bond angle of the molecule shown in the picture

State how many bonding pairs/pairs of electrons there are in the molecule

A

Linear

180º

2 pairs of electrons (all bonding)

77
Q

Give the shape and bond angle of the molecule shown in the picture

State how many bonding pairs/pairs of electrons there are in the molecule

A

Triangular planar

120º

3 pairs of electrons (all bonding pairs)

78
Q

Give the shape and bond angle of the molecule shown in the picture

State how many bonding pairs/pairs of electrons there are in the molecule

A

Tetrahedral

109.5º

4 pairs of electrons (all bonding pairs)

79
Q

Give the shape and bond angle of the molecule shown in the picture

State how many bonding pairs/pairs of electrons there are in the molecule

A

Trigonal Bipyrimidal

90º and 120º

5 pairs of electrons (all bonding pairs)

80
Q

Give the shape and bond angle of the molecule shown in the picture

State how many bonding pairs/pairs of electrons there are in the molecule

A

Octahedral

90º

6 electron pairs (all bonding)

81
Q

Give the shape and bond angle of the molecule shown in the picture

State how many bonding pairs/pairs of electrons there are in the molecule

A

Pyrimidal

107º

4 pairs of electrons

3 bonding pairs

1 lone pair

82
Q

Give the shape and bond angle of the molecule shown in the picture

State how many bonding pairs/pairs of electrons there are in the molecule

A

Bent (or V-shaped)

104.5º

4 electron pairs

2 bonding pairs

2 lone pairs

83
Q

Give the shape and bond angle of the molecule shown in the picture

State how many bonding pairs/pairs of electrons there are in the molecule

A

Bent (or V-shaped)

118º

3 electron groups

2 bonding groups

1 lone pair

84
Q

How do double/triple bonds affect the number of bonding electrons?

A

Can be thought of as a single group of electrons

e.g. CO2 has double 2 double bonds (4 electron pairs in total) which can be thought of as 2 electron groups

85
Q

What type of structure do ionic bonds have?

A

Always giant ionic (lattice)

86
Q

What type of structure do covalent bonds have?

A

Either simple molecular or giant covalent network

87
Q

What type of structure do metallic bonds have?

A

Always giant metallic lattice

88
Q

Describe the structure of a giant ionic lattice

A

Has a regular repeating pattern of postivitely and negatively charged ions in all 3 dimensions

The attraction between these oppositely charged ions outweighs the repulsion between ions with the same charge becayse the oppositely charged ions are closer

89
Q

What are the characteristic properties of giant ionic lattices?

A

High melting point because of strong electrostatic attractions between ions

Often soluble in water (due to charges of ions)

Conduct electricity when molten/in solution as charged ions able to move in response to voltage

90
Q

Describe the structure/bonding in simple molecular covalent bonding

A

Strong covalent bonds within molecules (between atoms) (strong intramolecular bonds)

But only weak intermolecular bonds between molecules

91
Q

What are the characteristic properties of simple covalent molecules?

A

Low melting point

Usually insoluble in water

Do not conduct electricity (or heat)

92
Q

Describe the structure/bonding in a giant covalent network

A

Contain billions of atoms with strong covalent bonds between them

93
Q

What are the characteristic properties of giant covalent networks?

A

High melting point because all bonds in structure are strong covalent bonds

Insoluble in water

Do not conduct electricity (apart from graphite)

94
Q

Describe the structure/bonding in a giant metallic lattice

A

All metals/metalic bonds have giant metallic lattice structure

Has a strong electrostatic attraction between the positive metal ions and the delocalised electrons between the ions

95
Q

What are the characteristic properties of giant metallic lattices?

A

High melting point because there is strong electrostatic attraction between ions + electrons

Insoluble in water

Conduct electricity when solid/molten because delocalised electrons are free to move in response to voltage

96
Q

What are delocalised electrons?

A

Electrons that are not associated with a particular atom

Instead are free to move over several atoms

97
Q

In the giant ionic lattice/crystalline structure of NaCl, 6 Cl- atoms arrange themselves around 1 smaller Na+ atom

The Cl- atoms arrange themselves to be as far apart as possible. Suggest the name for the 3D arrangement they will take up

A

Octahedral

98
Q

Group 1 metals have relatively low melting points

Use ideas about the charge and size of Group 1 ions to explain this

A

Group 1 ions have a small (+1) charge and have a relatively small ionic radius (compared with other ions in same period)

These 2 factors reduce the electrostatic attraction between the ions and delocalised electrons

They also only have 1 outer shell electron which can become delocalised, meaning electrostatic forces of attraction between ions and electrons are smaller

99
Q

What is electron affinity?

A

Energy change when 1mol gaseous atoms aquires 1mol electrons from 1mol gasous anions

100
Q

What is a complex ion?

A

Ion containing more than 1 atom

Charge is spread across whole ion

Contains covalent bonds

101
Q

What is a precipitate?

A

A suspension of solid particles fromed by a chemical reaction in solution

102
Q

What is a precipitation reaction?

A

Reaction between ions in solution that forms a precipitate

(suspension of solid particles/insoluble solid particles)

103
Q

What does the prefix -ate tell you about the composition of an ion?

A

There are 1+ non-metal ions/atoms bonded to oxygen

e.g. sulfate -SO42-

104
Q

What charge does the ClO ion have?

A

-1

105
Q

What is the formula + charge of a phosphate ion?

A

PO43-

106
Q

What charge does a lead ion have?

A

Pb2+

107
Q

What charge does a nitrate ion have?

A

NO3-

108
Q

What happens when ionic substances dissolve into solution/water?

A

The ions become surrounded by water + spread throughout the solution

They behave independantly of each other

109
Q

Which ionic substances are soluble?

A

All compounds containing…

Group 1 metals

Nitrate ions

Ammonium ions

… are soluble

110
Q

Which ionic substances are insoluble?

A

Sulfates of Ba, Ca, Pb, and Ag

Halides of Ag + Pb

All carbonates except those of Group 1/ammonium ions

Hydroxides containing some Group 2, Al, or d-block ions

111
Q

Give the colour of precipitate (if one forms) that forms from the reaction of OH- ions with:

Ca2+, Ba2+ Cu2+, Fe2+, Fe3+, Al3+, Pb2+, Zn2+, and Ag+

A

Ca2+ - White ppt

Ba2+ - White ppt

Cu2+ - Pale blue ppt

Fe2+ - Green ppt

Fe3+ - Brown ppt

Al3+ - White ppt

Pb2+ - White ppt

Zn2+ - White ppt

Ag+ - White ppt

112
Q

Give the colour of precipitate (if one forms) that forms from the reaction of SO42-- ions with:

Ca2+, Ba2+ Cu2+, Fe2+, Fe3+, Al3+, Pb2+, Zn2+, and Ag+

A

Ca2+ - White ppt

Ba2+ - White ppt

Cu2+ - Soluble

Fe2+ - Soluble

Fe3+ - Soluble

Al3+ - Soluble

Pb2+ - White ppt

Zn2+ - Soluble

Ag+ - White ppt

113
Q

Give the colour of precipitate (if one forms) that forms from the reaction of CO32- ions with:

Ca2+, Ba2+ Cu2+, Fe2+, Pb2+, Zn2+, and Ag+

A

Ca2+ - White ppt

Ba2+ - White ppt

Cu2+ - Green ppt

Fe2+ - Green ppt

Pb2+ - White ppt

Zn2+ - White ppt

Ag+ - White ppt

114
Q

Give the colour of precipitate (if one forms) that forms from the reaction of Cl- ions with:

Ca2+, Ba2+ Cu2+, Fe2+, Fe3+, Al3+, Pb2+, Zn2+, and Ag+

A

Ca2+ - Soluble

Ba2+ - Soluble

Cu2+ - Soluble

Fe2+ - Soluble

Fe3+ - Soluble

Al3+ - Soluble

Pb2+ - White ppt

Zn2+ - Soluble

Ag+ - White ppt

115
Q

Give the colour of precipitate (if one forms) that forms from the reaction of Cl- ions with:

Ca2+, Ba2+ Cu2+, Fe2+, Fe3+, Al3+, Pb2+, Zn2+, and Ag+

A

Ca2+ - Soluble

Ba2+ - Soluble

Cu2+ - Soluble

Fe2+ - Soluble

Fe3+ - Soluble

Al3+ - Soluble

Pb2+ - White ppt

Zn2+ - Soluble

Ag+ - White ppt

116
Q

Give the colour of precipitate (if one forms) that forms from the reaction of Cl- ions with:

Ca2+, Ba2+ Cu2+, Fe2+, Fe3+, Al3+, Pb2+, Zn2+, and Ag+

A

Ca2+ - Soluble

Ba2+ - Soluble

Cu2+ - Soluble

Fe2+ - Soluble

Fe3+ - Soluble

Al3+ - Soluble

Pb2+ - White ppt

Zn2+ - Soluble

Ag+ - White ppt

117
Q

Give the colour of precipitate (if one forms) that forms from the reaction of Br- ions with:

Ca2+, Ba2+ Cu2+, Fe2+, Fe3+, Al3+, Pb2+, Zn2+, and Ag+

A

Ca2+ - Soluble

Ba2+ - Soluble

Cu2+ - Soluble

Fe2+ - Soluble

Fe3+ - Soluble

Al3+ - Soluble

Pb2+ - White ppt

Zn2+ - Soluble

Ag+ - Cream ppt

118
Q

Give the colour of precipitate (if one forms) that forms from the reaction of I- ions with:

Ca2+, Ba2+ Cu2+, Fe2+, Fe3+, Al3+, Pb2+, Zn2+, and Ag+

A

Ca2+ - Soluble

Ba2+ - Soluble

Cu2+ - Soluble

Fe2+ - Soluble

Fe3+ - Soluble

Al3+ - Soluble

Pb2+ - Yellow ppt

Zn2+ - Soluble

Ag+ - Yellow ppt

119
Q

The presence of which ions can be tested for by adding barium chloride solution?

(Solution containing Ba2+ ​ions)

A

Sulfate (SO42-) ions - white ppt formed

120
Q

The presence of which ions can be tested for by adding silver nitrate solution (containing Ag+ ions)?

A

Shows presence of halide ions

White ppt forms for Cl-
Cream ppt forms for Br-
Yello ppt forms for I-

121
Q

What are the 4 ways/reactions that can be used to make an ionic salt?

A

Acid + base/alkali → Salt + Water

Acid + Carbonate → Salt + Water + CO2

Acid + Metal → Salt + Hydrogen (MASH)

122
Q

Give three examples of practical applications of precipitation reactions

A

Water treatment

Production of coloured pigments for paints/dyes

Identification of certain metal ions in solutions

123
Q

What is relative isotopic mass?

A

The mass of one atom of an isotope compared to 1/12 of the mass of a 12C atom

(The same as Ar but for isotopes…)

124
Q

What is empirical formula?

A

The simplest ratio of atoms in a compound

125
Q

What is water of crystallisation?

A

Number of water molecules contained in an ionic lattice per molecule of salt

(i.e. how hydrates the salt is)

126
Q

Hydrates cobalt (II) chloride has the formula CoCl2•xH2O

1.173g of hydrated cobalt chloride is heated to drive off the water of crytallisation. The mass remaining is 0.641g.

Calculate the formula of the hydrated salt

A

Calculate the mass of water removed:
1.173 - 0.641 = 0.532g

Calculate the moles of water removed:
0.532 / 18 = 0.02956mol

Calculate the moles of anhydrous salt:
0.641 / 129.9 = 0.00493mol

Calculate the ratio

  • *moles water / moles anhydrous salt** =
    0. 02956 / 0.00493 = 5.996 = 6
127
Q

Calculate the percentage by mass of nitrogen in ammonium sulfate (NH4)2SO4

A

Calculate the Mr of (NH4)2SO4 = 132.1

Calculate the mass of N in 1 mol (NH4)2SO4:
14 x 2 = 28

% by mass = mass element in 1mol / Mr compound x 100
(28/132.1) x 100 = 21.2%

128
Q

What is the formula for percentage yield?

A

(Experimental yield / theoretical yield) x 100

Yield given by moles of reactants/products from balanced equation

129
Q

Why might percentage yield be lower than expected?

A

Loss of product from reaction vessels (when transfering)

Side reactions (may create by-products)

Impurities in reactants

Changes in temp + pressure (may effect equilibrium)

If the reaction is an equilibrium system

130
Q

What are the 3 main factors that affect 1st ionisation enthalpy?

A

Atomic radii - larger = lower

Nuclear charge - more protons = higher

Electron shielding - outer shells feel less electrostatic attraction to nucleus

131
Q

How do group 2 metals react with water?

A

Group 2 metal + Water → Metal hydroxide + Hydrogen

M(s) + 2H2O(l) → M(OH)2(s) + H2(g)

132
Q

How does the vigourousnes of Group 2 metals change as you go down the group?

A

The reactions become more vigorous as you go down the group

133
Q

How do the hydroxides of Group 2 metals change as you go down the group?

A

Become increasing soluble and more alkaline

134
Q

How do Group 2 metals react with oxygen (when heated)?

A

Metal + Oxygen → Metal oxide

2M(s) + O2(g) → 2MO(s)

135
Q

What substances do Group 2 metal oxides react with?

What property does this give them?

A

They react with acids, so can act as bases

Metal oxide + Acid → Salt + Water

MO(s) + H2SO4(aq) → MSO4(aq) + H2O(l)

136
Q

How do metal hydroxides react with acids?

A

Metal hydroxide + Acid → Salt + Water

M(OH)2(s/aq) + 2HCl(aq) → MCl2(aq) + 2H2O(l)

137
Q

What happens to Group 2 metal carbonates when they are heated?

Give the general equation

A

Undergo thermal decomposition

Metal carbonate → Metal oxide + Carbon dioxide

MCO3(s) → MO(s) + CO2(g)

138
Q

How does the thermal stability of Group 2 carbonates change going down the group?

A

As you go down the group, thermal stability increases

Means that metals further down decompose at higher temperatures than those further up the group

139
Q

Why does the thermal stability of Group 2 carbonates increase as you go down the group?

A

M2+ ions get larger as you go down the group, so their charge density is lower

Because ions higher up the group have greater charge densities, they polarise the carbonate ion more

The more polarised the carbonate ion, the more likely it is to break up + form an oxide ion and CO2

140
Q

What is a polarised ion?

A

A large (complex) ion that can have its electron distribution altered by small, highly-charged ions

This is known as polarisation

141
Q

What is charge density?

A

The charge of an ion relative to its size

Mg2+ has a greater charge density than Ba2+ because, although they both have the same overall charge, Mg2+ is smaller

142
Q

10.0cm3 of a 1.00moldm-3 solution of HCl was transfered to a volumetric flask and made up to 250cm3 with water.

What is the concentration of the diluted solution?

A

Calculate the dilution factor:
10/250 = 0.04

Use this to calculate the new conc. of the diluted solution:
0.04 x 1.00 = 0.0400moldm-3

143
Q

What is an acid?

A

A substance that produces/donates H+ ions in a solution

144
Q

What is a base?

A

A copound that reacts with an acid to produce water and a salt

Is a proton acceptor

145
Q

What is an alkali?

A

A soluble base

Dissolves in water to produce hydroxide (OH-) ions

146
Q

What is an oxonium ion?

What is its formula?

A

H3O+(aq)

Hydrogen ion bonded to a molecule of water

Created due to presence of H+ in acidic solutions (hence present in all acidic solutions)

147
Q

Briefly describe how a soluble salt can be made

A

By reacting the appropriate acid and alkali together

The solid salt can then be produced by evaporating the excess solution/water

148
Q

Breifly describe how an insoluble salt can be made

A

By a precipitation reaction

E.g. silver iodide can be made by reacting silver nutrate + potassium iodide

A Level Chemistry (OCR B Salters) (12 decks)

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