Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

Chemistry > EL > Flashcards

EL Flashcards

1
Q

Why do isotopes have the same reactivity and chemical properties?

A

They have the same configuration of electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Why do isotopes have different physical properties e.g rate of diffusion?

A

Different atomic mass

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

How do most elements exist naturally?

A

As a mixture of isotopes

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

What determines the chemical properties of an element?

A

The number and arrangement of electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

What often determines the physical properties of an element?

A

The mass of the atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Describe Bohr’s Theory.

A

The electron in hydrogen only exists in certain energy levels.
A photon is emitted when an electron changes energy level.
Energy of a photon is equal to the difference between two energy levels.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Describe an s subshell.

A

One s orbital, two electrons.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Describe a p subshell.

A

Three p orbitals, six electrons.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Describe a d subshell.

A

Five d orbitals, ten electrons.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Describe an f subshell.

A

Seven f orbitals, fourteen electrons.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Define dative covalent bonds.

A

Both bonding electrons come from the same atom.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

What is the bond angle in a linear molecule?

A

180°.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

What is the bond angle in a planar molecule?

A

120°.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

What is the bond angle in a bipyramidal molecule?

A

90° and 120°.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

What is the bond angle in an octahedral molecule?

A

90°.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

What is the bond angle in a tetrahedral molecule?

A

109.5°.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

How many degrees does a lone pair remove?

A

2.5°.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

What colour precipitate is formed when potassium iodide is added to Pb2+

A

yellow.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

What colour precipitates are formed when silver nitrate is added to Cl-, Br-, and I-.

A

Cl- white.
Br- cream.
I- yellow.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Describe ionic bonding.

A

Strong electrostatic attraction between oppositely charged ions in a metal and a non metal

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

Describe group 1 and 2 trends.

A

More metallic down a group, less metallic across a period.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

How did the structure of the atom develop over time?

A

First was plum pudding model

Disproved by Geiger marsden

Bohr model - protons, neutrons and electrons

Evidence for shells - ionisation enthalpies and spectra

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

How do covalent bonds work?

A

There is a balance between repulsive forces between the nuclei and the attractive forces between the nuclei and electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

What is the bond angle of a pyramidal molecule with a lone pair?

A

107

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

What is the bond angle of a bent molecule with two lone pairs?

A

104.5

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

How do giant ionic compounds work?

A

Overall attraction in a lattice made of attraction between 1 of different charge and repulsion of ions of the same charge

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q

Explain the trend in melting points in period 2 and 3

A

In period 2 melting point increases until carbon/silicon (increased delocalised electrons in metallic structure) then drops suddenly and decreases as molecules have less atoms so weaker intermolecular bonds

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

Negative ions with -1 charge

A

Nitrate NO3

Hydroxide

Hydrogencarbonate

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

Ions with -2 charge

A

Sulphate

Carbonate

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

Ions with + charge

A

Ammonium

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
31
Q

Ions with +2 charge

A

Copper

Zinc

Iron (II)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
32
Q

Ion with +3 charge

A

Iron (III)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
33
Q

What happens and what is the trend when group 2 metals react with water?

A

Form metal hydroxides and hydrogen

Increasing reactivity down group as outer electrons are more easily lost

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
34
Q

What happens when group 2 metals react with oxygen?

A

Form metal oxides

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
35
Q

What is the trend in solubility of group 2 hydroxides?

A

Increase in solubility down the group

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
36
Q

What is the trend in solubility of group 2 carbonates?

A

Decreases down group

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
37
Q

Define ionisation enthalpy?

A

The enthalpy needed to remove the 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous ions

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
38
Q

What is the trend in ionisation enthalpies and why?

A

Increase across the period
And decreases down group

Atomic radius decreases

Nuclear charge Increases

Electron shielding

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
39
Q

Solubility of nitrates?

A

Soluble

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
40
Q

Sulfates are soluble except?

A

Barium

Calcium

Lead

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
41
Q

Solubility of carbonates?

A

Insoluble except lithium potassium sodium and ammonium

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
42
Q

Solubility of hydroxides?

A

Insoluble expect lithium sodium potassium strontium calcium barium and ammonium

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
43
Q

Test for calcium?

A

White precipitate with NaOH

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
44
Q

Test for copper?

A

Blue precipitate with NaOH

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
45
Q

Test for iron (ii)?

A

Green precipitate with NaOH

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
46
Q

Test for iron (iii)?

A

Brown precipitate with NaOH

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
47
Q

Test for aluminium?

A

White dissolvable precipitate with NaOH

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
48
Q

Test for carbonates?

A

react with HCl - forms CO2 so turns lime water cloudy

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
49
Q

Test for sulfates?

A

Barium chloride

Forms a white precipitate

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
50
Q

Test for ammonia

A

Litmus paper

Red –> blue

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
51
Q

Test for halides?

A

Add silver nitrate

White

Cream

Yellow

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
52
Q

What happens when you add group 2 oxides and hydroxides to water?

A

They form alkaline solutions

Oxides form stronger alkalines down the group

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
53
Q

Group 2 hydroxides/oxides and acids?

A

Neutralisation

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
54
Q

What is the order of the electromagnetic spectrum?

A

Infrared visible ultraviolet

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
55
Q

What is the melting point, solubility in water and conductivity of giant metallic structures?

A

High melting point

Insoluble in water, some react

Conductive when solid or liquid

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
56
Q

What is the melting point, solubility in water and conductivity in giant covalent ?

A

Very high melting point

Insoluble

Unconductive - except graphite

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
57
Q

What is the melting point, solubility in water and conductivity of simple covalent?

A

Low melting point

Insoluble unless the have polar group

Inconductive

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
58
Q

What is an acid?

A

proton (H+) donor

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
59
Q

What is a base?

A

proton (H+) acceptor

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
60
Q

What is an alkali?

A

Base that dissolved in water to produce OH-

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
61
Q

What do solid and dashed wedges mean?

A

Solid wedges coming out of plane of paper, dashed are going behind

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
62
Q

Define activation energy

A

minimum energy needed in a collision to cause a reaction

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
63
Q

What is meant by the term mole?

A

The amount of substance that has the same number of particles as there are atoms in 12g of carbon 12

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
64
Q

What is a salt?

A

A compound where a H+ in an acid has been replaced by a metal ion

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
65
Q

What is the maximum number of electrons in ‘s’ sub-shell?

A

2

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
66
Q

What is maximum number of electrons in ‘p’ sub shell?

A

6 (2 in each of 3 p orbitals)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
67
Q

What is the maximum number of electrons in ‘d’ sub shell?

A

10 (5 x 2)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
68
Q

List order in which sub-shells fill with electrons

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
69
Q

What are the four pieces of info needed to describe an electron (its movement around the nucleus)?

A
  • electron shell its in
  • its sub shell
  • its orbital within the sub shell
  • its spin
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
70
Q

Why do electrons fill orbitals singularly before pairing up?

A

Keeps electrons as far away from each other as possible

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
71
Q

What are the 3 p orbitals called?

A

Px, Py, Pz

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
72
Q

How are the elements arranged in Mendeleev’s version of the periodic table?

A

by increasing relative atomic mass

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
73
Q

How are the elements arranged in the modern version of the periodic table?

A

by the atomic number (number of protons)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
74
Q

The elements are broken into 4 main blocks, which letters are used?

A

s, p, d and f block

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
75
Q

What is periodicity?

A

The occurrence of a pattern that is repeated as you go across periods

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
76
Q

What is a closed shell arrangement?

A

These are elements which have a full out shell, these are very stable.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
77
Q

Why do bonds push each other as far apart as possible?

A

To reduce the repulsive forces between their pairs of electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
78
Q

Why are single bonds equally spaced?

A

Because the repulsion between each bond are equal

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
79
Q

How many bond pairs does a linear molecule have?

A

2

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
80
Q

Define a linear molecule

A

An atom/ion where two groups of electrons surround its centre

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
81
Q

How many bond pairs does a triagonal planar/planar triangular molecule have?

A

3

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
82
Q

Define a triagonal planar/planar triangular molecule

A

An atom/ion where three groups of electrons surround its centre

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
83
Q

How many bond pairs does a tetrahedral molecule have?

A

4

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
84
Q

Define a tetrahedral molecule

A

An atom or ion where four groups of electrons surround its centre

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
85
Q

How many bond pairs does a bipyramidal molecule have?

A

5

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
86
Q

Define a bipyramidal molecule

A

An atom/ion where five groups of electrons surround its centre

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
87
Q

How many bond pairs does an octahedral molecule have?

A

6

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
88
Q

Define an octahedral molecule

A

An atom/ion where six groups of electrons surround its centre

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
89
Q

Why are shapes sometimes slightly distorted away from their regular shapes in a molecule/ion?

A

If the molecule/ion has lone pairs on the central atom, extra repulsion caused by them results in distortion of the regular shape

90
Q

How many groups of electrons does a triagonal based pyramid molecule have?

A

4

91
Q

What 2 types of electron groups does a triagonal based pyramid molecule have?

A
  • 3x single covalent bonds

- 1x lone pair of electrons

92
Q

What causes ammonia to have a triagonal based pyramid molecular shape?

A

The lone pair in nitrogen repulses the hydrogen-nitrogen bond pairs, pushing them closer together

93
Q

What shape is the triagonal based pyramid and bent/v-shaped shape based on?

A

Tetrahedral shape

94
Q

How many groups of electrons does a bent/v-shaped molecule have?

A

4

95
Q

What types of groups of electrons does a bent/v-shaped molecule have?

A
  • 2x single covalent bonds

- 2x lone pair of electrons

96
Q

Why does water have a bent/v-shaped molecular shape?

A

The lone pair to lone pair repulsions of the two lone pairs in the oxygen pushes the Oxygen-Hydrogen bonds together

97
Q

Define a “lone pair”

A

Pairs of electrons uninvolved in the bonding of a covalent structure

98
Q

Why do substances want a full outer shell?

A

As this is the most energetically stable arrangement

99
Q

Define “electronegativity”

A

The ability of an atom to attract the electron pair in a covalent bond to itself

100
Q

Explain why a non polar bond is non polar.

A

A bond with similar atoms - so they have the same electronegativity so they will both pull on the electrons to the same extent and be equally shared

101
Q

Explain why a polar bond is polar.

A
  • A bond with different atoms - so they have different electronegativity
  • Therefore one will pull the electron pair closer to its end
  • It will be slightly more negative than the overall bond (δ−)
  • The other atom will be slightly less negative - more positive (δ+)
  • A dipole is formed and the bond is said to be polar
102
Q

How does the electronegativity of the atoms in a polar bond relate to its polarity?

A

The greater the difference in electronegativity between the two atoms in the bond, the greater the bond’s polarity

103
Q

What is a non-polar covalent bond between?

A

The same/very similar non metals

104
Q

What is a polar covalent bond between?

A

Different non-metals

105
Q

Order the following in order of increasing polarisation:

  • Polar covalent
  • Metallic
  • Non - polar covalent
  • Ionic
A
  • Non - polar covalent
  • Polar covalent
  • Ionic
  • Metallic
106
Q

What are the electrical properties of covalent substances and why?

A

They don’t conduct electricity because they have no mobile ions/electrons

107
Q

Are covalent substances soluble in water and why?

A
  • No they tend to be more so in organic solvents

- The polar water molecules are more attracted to each other than the molecular substance and so do not react with it

108
Q

Why do covalent bonds have a low melting point?

A

Because they have weak intermolecular forces which means little energy is needed to separate molecules from each other

109
Q

Why does CH4 have a lower melting point than C2H6?

A

Because the molecule has a smaller surface area

110
Q

Define relative atomic mass

A

Mass of atom relative to carbon 12

111
Q

Define relative (molecular mass)

A

Sum of relative atomic masses of elements making up compound

112
Q

Define molecular formula

A

Actual number of atoms of each element in a molecule

113
Q

Are ionic substances soluble?

A
  • dissolve in polar solvents (H2O)

- do NOT dissolve in non-polar solvents, no regions of positive/negative charge, no new bonds form

114
Q

Do ionic substance conduct electricity?

A
  • solids do NOT

- molten/aqueous solution DOES; electrons can move and carry charge

115
Q

Define giant metallic lattice

A

3D structure of positive metal ions and delocalised electrons held together by strong metallic bonds

116
Q

Is mp and bp of metallic lattice high or low?

A
  • high

- lots of energy needed to break strong metallic bonds

117
Q

Are metallic lattices soluble?

A
  • insoluble in polar and non-polar
118
Q

Can metallic lattices conduct electricity?

A

Yes; delocalised electrons carry charge

119
Q

Give properties of metallic lattices

A
  • ductile; can be stretched
  • malleable; hammered into shapes
  • can be alloyed; mixed with other metals
120
Q

A substance has a high boiling point, poor solubility in polar and non polar substances and does not conduct electricity. What structure may this substance have?

A

Giant covalent

121
Q

Why can ionic substances conduct electricity when dissolved in water

A

The ions are mobile and can be attracted to charged electrodes.

Not the movement of free electrons!!!

122
Q

What is the difference between an alkali and a base?

A

An alkali is a soluble base

123
Q

Which sulfates will NOT dissolve in water?

A

barium, calcium, lead, and silver sulfates

124
Q

Which halides will NOT dissolve in water?

A

silver and lead halides

125
Q

Which carbonates will NOT dissolve in water?

A

All metal carbonates are insoluble

126
Q

Which hydroxides will NOT dissolve in water?

A

All except group 1 hydroxides and ammonium hydroxide.

127
Q

Define ionisation energy

A

Energy required to remove an electron from an atom, forming a positive ion

128
Q

Define first ionisation energy

A

Energy required to remove 1 electron from each atom in 1 mole of gaseous atoms

129
Q

How does atomic radius affect ionisation energy?

A
  • as atomic radius increases, I.E decreases, due to weaker nuclear attraction
130
Q

How does nuclear charge affect ionisation energy?

A
  • greater the nuclear charge, greater nuclear attraction, I.E increases
131
Q

How does electron shielding affect ionisation energy?

A
  • inner shells of electrons repel outer electrons, therefore more inner shell means greater shielding, less nuclear attraction therefore lower I.E
132
Q

Why do successive ionisation energies increase?

A

As electrons are removed, nuclear charge increases, ionic radius decreases, nuclear attraction increases

133
Q

How can a graph showing successive ionisation energies allow you to predict the group of an element?

A
  • there are large jumps between each energy level
134
Q

Does electronegativity increase or decrease across a period? And why

A
  • increases

- atomic radius decreases, stronger nuclear attraction

135
Q

Does electronegativity increase or decrease down a group? And why

A

Decrease

- atomic radius increases, weaker nuclear attraction

136
Q

Do elements become more or less metallic across a period?

A
  • less metallic

- reactivity decreases across period

137
Q

What is calcium hydroxide used for?

A
  • neutralising acidic soil
138
Q

What is magnesium hydroxide used for?

A

Neutralising stomach acid

139
Q

Does thermal decomposition become easier or harder down group and why?

A
  • harder
  • smaller ions have greater charge density, therefore greater polarising effect on negatively charged cloud around carbonate ion, therefore carbonate is less stable
140
Q

Acid + alkali =

A

Salt + water

141
Q

Acid + base =

A

Salt + water

142
Q

Acid + carbonate =

A

Salt + water + carbon dioxide

143
Q

Acid + metal =

A

Salt + hydrogen

144
Q

Acid + metal oxide =

A

Salt + water

145
Q

What is a salt?

A

Compound formed when hydrogen ion in an acid is replaced with a metal ion or ammonium ion

146
Q

Give some properties of acids

A
  • turns litmus paper red

- neutralised by bases

147
Q

What do organic acids contain?

A

Carboxyl group (COOH)

148
Q

What occurs when an acid and a base react?

A

They neutralise each other

149
Q

What is titration used for?

A

To determine concentration of a solvent

150
Q

Give method of volumetric analysis (titration)

A
  1. Standard solution is measured with volumetric pipette, put in conical flask with indicator
  2. Calculate number of moles in solution
  3. Using burette, add unknown solution until colour changes
  4. Calculate number of moles that reacted (using volume used) and calculate concentration
151
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

152
Q

What is nuclear fission?

A

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

153
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

154
Q

What are isotopes?

A

Atoms of the same element with a different number of neutrons

This causes mass number to be different

155
Q

What is the Avogadro constant (NA)?

A

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

156
Q

What does quantised mean?

A

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

157
Q

What is the ground state?

A

The lowest energy level that an electron can occupy

158
Q

What is a photon?

A

Quanta of energy in the form of electromagnetic radiation

159
Q

Briefly 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

160
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…

161
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

162
Q

What equation links the wave + particle models of light?

A

ΔE = hv

ΔE = energy of photon (J)
h = Planck's Constant
v = frequency (Hz/s-1)
163
Q

What equation explains the wave properties of light?

A

c = vλ

c= speed of light (ms-1)
v = frequency (Hz/s-1)
λ = wavelength (m)
164
Q

What is spectroscopy?

A

The study of how light and matter interact

Uses IR, visible, and UV light

165
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 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

166
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 unique emission/absorption spectrum

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

167
Q

What colour flame does Li+ give?

A

Bright red

168
Q

What colour flame does Na+ give?

A

Yellow

169
Q

What colour flame does K+ give?

A

Lilac

170
Q

What colour flame does Ca2+ give?

A

Brick red

171
Q

What colour flame does Ba2+ give?

A

Apple green

172
Q

What colour flame does Cu2+ give?

A

Blue-green

173
Q

Describe the appearance of an absorption spectrum

A

Black lines on a colourful background (showing all colours of visible light)

174
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

175
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

176
Q

What are the differences between atomic emission and absorption spectra?

A

Emission: coloured lines on a black background

Absorption: black lines on a coloured background

177
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

178
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

179
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

180
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.

181
Q

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

A

Melting point increases then decreases across the period

This is because the metals on the left-hand side of a period are metallically bonded so have higher melting points due to the delocalised 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.

182
Q

What is the general equation for first ionisation enthalpy?

A

X(g) → X+(g) + e-

183
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

184
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)

185
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

186
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 electrons due to the greater electron affinity as a result of a more positive nucleus

187
Q

What is a lone pair?

A

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

188
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. bonding 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) […]

189
Q

What type of structure do ionic bonds have?

A

Always giant ionic (lattice)

190
Q

What type of structure do covalent bonds have?

A

Either simple molecular or giant covalent network

191
Q

What type of structure do metallic bonds have?

A

Always giant metallic lattice

192
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

193
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

194
Q

What are the characteristic properties of simple covalent molecules?

A

Low melting point

Usually insoluble in water

Do not conduct electricity (or heat)

195
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)

196
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

197
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

198
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

199
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

200
Q

What is a precipitate?

A

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

201
Q

What is a precipitation reaction?

A

Reaction between ions in solution that forms a precipitate

suspension of solid particles/insoluble solid particles

202
Q

Which ionic substances are soluble?

A

All compounds containing…

Group 1 metals

Nitrate ions

Ammonium ions

… are soluble

203
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

204
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

205
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

206
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

207
Q

What is empirical formula?

A

The simplest ratio of atoms in a compound

208
Q

What is water of crystallisation?

A

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

209
Q

What is the formula for percentage yield?

A

(Experimental yield / theoretical yield) x 100

210
Q

Why might percentage yield be lower than expected?

A

Loss of product from reaction vessels (when transferring)

Side reactions (may create by-products)

Impurities in reactants

Changes in temp + pressure (may effect equilibrium)

If the reaction is an equilibrium system

211
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

212
Q

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

A

Become increasing soluble and more alkaline

213
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

214
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

215
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

216
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

217
Q

What is charge density?

A

The charge of an ion relative to its size

218
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)

219
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

220
Q

Briefly describe how an insoluble salt can be made

A

By a precipitation reaction

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

221
Q

Explain how the lines on spectra can be used to identify elements

A

• atoms of each element have their own specific
energy levels
• thus different gaps
• thus different frequency lines
• frequencies of lines can be checked against a
database

222
Q

Describe the appearance of emission spectra

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)

Chemistry (6 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions