EL Flashcards
(222 cards)
1
Q
Why do isotopes have the same reactivity and chemical properties?
A
They have the same configuration of electrons
2
Q
Why do isotopes have different physical properties e.g rate of diffusion?
A
Different atomic mass
3
Q
How do most elements exist naturally?
A
As a mixture of isotopes
4
Q
What determines the chemical properties of an element?
A
The number and arrangement of electrons
5
Q
What often determines the physical properties of an element?
A
The mass of the atom
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.
7
Q
Describe an s subshell.
A
One s orbital, two electrons.
8
Q
Describe a p subshell.
A
Three p orbitals, six electrons.
9
Q
Describe a d subshell.
A
Five d orbitals, ten electrons.
10
Q
Describe an f subshell.
A
Seven f orbitals, fourteen electrons.
11
Q
Define dative covalent bonds.
A
Both bonding electrons come from the same atom.
12
Q
What is the bond angle in a linear molecule?
A
180°.
13
Q
What is the bond angle in a planar molecule?
A
120°.
14
Q
What is the bond angle in a bipyramidal molecule?
A
90° and 120°.
15
Q
What is the bond angle in an octahedral molecule?
A
90°.
16
Q
What is the bond angle in a tetrahedral molecule?
A
109.5°.
17
Q
How many degrees does a lone pair remove?
A
2.5°.
18
Q
What colour precipitate is formed when potassium iodide is added to Pb2+
A
yellow.
19
Q
What colour precipitates are formed when silver nitrate is added to Cl-, Br-, and I-.
A
Cl- white.
Br- cream.
I- yellow.
20
Q
Describe ionic bonding.
A
Strong electrostatic attraction between oppositely charged ions in a metal and a non metal
21
Q
Describe group 1 and 2 trends.
A
More metallic down a group, less metallic across a period.
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
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
24
Q
What is the bond angle of a pyramidal molecule with a lone pair?
A
107
25
What is the bond angle of a bent molecule with two lone pairs?
104.5
26
How do giant ionic compounds work?
Overall attraction in a lattice made of attraction between 1 of different charge and repulsion of ions of the same charge
27
Explain the trend in melting points in period 2 and 3
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
28
Negative ions with -1 charge
Nitrate NO3
Hydroxide
Hydrogencarbonate
29
Ions with -2 charge
Sulphate
Carbonate
30
Ions with + charge
Ammonium
31
Ions with +2 charge
Copper
Zinc
Iron (II)
32
Ion with +3 charge
Iron (III)
33
What happens and what is the trend when group 2 metals react with water?
Form metal hydroxides and hydrogen
Increasing reactivity down group as outer electrons are more easily lost
34
What happens when group 2 metals react with oxygen?
Form metal oxides
35
What is the trend in solubility of group 2 hydroxides?
Increase in solubility down the group
36
What is the trend in solubility of group 2 carbonates?
Decreases down group
37
Define ionisation enthalpy?
The enthalpy needed to remove the 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous ions
38
What is the trend in ionisation enthalpies and why?
Increase across the period
And decreases down group
Atomic radius decreases
Nuclear charge Increases
Electron shielding
39
Solubility of nitrates?
Soluble
40
Sulfates are soluble except?
Barium
Calcium
Lead
41
Solubility of carbonates?
Insoluble except lithium potassium sodium and ammonium
42
Solubility of hydroxides?
Insoluble expect lithium sodium potassium strontium calcium barium and ammonium
43
Test for calcium?
White precipitate with NaOH
44
Test for copper?
Blue precipitate with NaOH
45
Test for iron (ii)?
Green precipitate with NaOH
46
Test for iron (iii)?
Brown precipitate with NaOH
47
Test for aluminium?
White dissolvable precipitate with NaOH
48
Test for carbonates?
react with HCl - forms CO2 so turns lime water cloudy
49
Test for sulfates?
Barium chloride
| Forms a white precipitate
50
Test for ammonia
Litmus paper
| Red --> blue
51
Test for halides?
Add silver nitrate
White
Cream
Yellow
52
What happens when you add group 2 oxides and hydroxides to water?
They form alkaline solutions
| Oxides form stronger alkalines down the group
53
Group 2 hydroxides/oxides and acids?
Neutralisation
54
What is the order of the electromagnetic spectrum?
Infrared visible ultraviolet
55
What is the melting point, solubility in water and conductivity of giant metallic structures?
High melting point
Insoluble in water, some react
Conductive when solid or liquid
56
What is the melting point, solubility in water and conductivity in giant covalent ?
Very high melting point
Insoluble
Unconductive - except graphite
57
What is the melting point, solubility in water and conductivity of simple covalent?
Low melting point
Insoluble unless the have polar group
Inconductive
58
What is an acid?
proton (H+) donor
59
What is a base?
proton (H+) acceptor
60
What is an alkali?
Base that dissolved in water to produce OH-
61
What do solid and dashed wedges mean?
Solid wedges coming out of plane of paper, dashed are going behind
62
Define activation energy
minimum energy needed in a collision to cause a reaction
63
What is meant by the term mole?
The amount of substance that has the same number of particles as there are atoms in 12g of carbon 12
64
What is a salt?
A compound where a H+ in an acid has been replaced by a metal ion
65
What is the maximum number of electrons in 's' sub-shell?
2
66
What is maximum number of electrons in 'p' sub shell?
6 (2 in each of 3 p orbitals)
67
What is the maximum number of electrons in 'd' sub shell?
10 (5 x 2)
68
List order in which sub-shells fill with electrons
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f
69
What are the four pieces of info needed to describe an electron (its movement around the nucleus)?
- electron shell its in
- its sub shell
- its orbital within the sub shell
- its spin
70
Why do electrons fill orbitals singularly before pairing up?
Keeps electrons as far away from each other as possible
71
What are the 3 p orbitals called?
Px, Py, Pz
72
How are the elements arranged in Mendeleev's version of the periodic table?
by increasing relative atomic mass
73
How are the elements arranged in the modern version of the periodic table?
by the atomic number (number of protons)
74
The elements are broken into 4 main blocks, which letters are used?
s, p, d and f block
75
What is periodicity?
The occurrence of a pattern that is repeated as you go across periods
76
What is a closed shell arrangement?
These are elements which have a full out shell, these are very stable.
77
Why do bonds push each other as far apart as possible?
To reduce the repulsive forces between their pairs of electrons
78
Why are single bonds equally spaced?
Because the repulsion between each bond are equal
79
How many bond pairs does a linear molecule have?
2
80
Define a linear molecule
An atom/ion where two groups of electrons surround its centre
81
How many bond pairs does a triagonal planar/planar triangular molecule have?
3
82
Define a triagonal planar/planar triangular molecule
An atom/ion where three groups of electrons surround its centre
83
How many bond pairs does a tetrahedral molecule have?
4
84
Define a tetrahedral molecule
An atom or ion where four groups of electrons surround its centre
85
How many bond pairs does a bipyramidal molecule have?
5
86
Define a bipyramidal molecule
An atom/ion where five groups of electrons surround its centre
87
How many bond pairs does an octahedral molecule have?
6
88
Define an octahedral molecule
An atom/ion where six groups of electrons surround its centre
89
Why are shapes sometimes slightly distorted away from their regular shapes in a molecule/ion?
If the molecule/ion has lone pairs on the central atom, extra repulsion caused by them results in distortion of the regular shape
90
How many groups of electrons does a triagonal based pyramid molecule have?
4
91
What 2 types of electron groups does a triagonal based pyramid molecule have?
- 3x single covalent bonds
| - 1x lone pair of electrons
92
What causes ammonia to have a triagonal based pyramid molecular shape?
The lone pair in nitrogen repulses the hydrogen-nitrogen bond pairs, pushing them closer together
93
What shape is the triagonal based pyramid and bent/v-shaped shape based on?
Tetrahedral shape
94
How many groups of electrons does a bent/v-shaped molecule have?
4
95
What types of groups of electrons does a bent/v-shaped molecule have?
- 2x single covalent bonds
| - 2x lone pair of electrons
96
Why does water have a bent/v-shaped molecular shape?
The lone pair to lone pair repulsions of the two lone pairs in the oxygen pushes the Oxygen-Hydrogen bonds together
97
Define a "lone pair"
Pairs of electrons uninvolved in the bonding of a covalent structure
98
Why do substances want a full outer shell?
As this is the most energetically stable arrangement
99
Define "electronegativity"
The ability of an atom to attract the electron pair in a covalent bond to itself
100
Explain why a non polar bond is non polar.
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
Explain why a polar bond is polar.
- 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
How does the electronegativity of the atoms in a polar bond relate to its polarity?
The greater the difference in electronegativity between the two atoms in the bond, the greater the bond's polarity
103
What is a non-polar covalent bond between?
The same/very similar non metals
104
What is a polar covalent bond between?
Different non-metals
105
Order the following in order of increasing polarisation:
- Polar covalent
- Metallic
- Non - polar covalent
- Ionic
- Non - polar covalent
- Polar covalent
- Ionic
- Metallic
106
What are the electrical properties of covalent substances and why?
They don't conduct electricity because they have no mobile ions/electrons
107
Are covalent substances soluble in water and why?
- 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
Why do covalent bonds have a low melting point?
Because they have weak intermolecular forces which means little energy is needed to separate molecules from each other
109
Why does CH4 have a lower melting point than C2H6?
Because the molecule has a smaller surface area
110
Define relative atomic mass
Mass of atom relative to carbon 12
111
Define relative (molecular mass)
Sum of relative atomic masses of elements making up compound
112
Define molecular formula
Actual number of atoms of each element in a molecule
113
Are ionic substances soluble?
- dissolve in polar solvents (H2O)
| - do NOT dissolve in non-polar solvents, no regions of positive/negative charge, no new bonds form
114
Do ionic substance conduct electricity?
- solids do NOT
| - molten/aqueous solution DOES; electrons can move and carry charge
115
Define giant metallic lattice
3D structure of positive metal ions and delocalised electrons held together by strong metallic bonds
116
Is mp and bp of metallic lattice high or low?
- high
| - lots of energy needed to break strong metallic bonds
117
Are metallic lattices soluble?
- insoluble in polar and non-polar
118
Can metallic lattices conduct electricity?
Yes; delocalised electrons carry charge
119
Give properties of metallic lattices
- ductile; can be stretched
- malleable; hammered into shapes
- can be alloyed; mixed with other metals
120
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?
Giant covalent
121
Why can ionic substances conduct electricity when dissolved in water
The ions are mobile and can be attracted to charged electrodes.
Not the movement of free electrons!!!
122
What is the difference between an alkali and a base?
An alkali is a soluble base
123
Which sulfates will NOT dissolve in water?
barium, calcium, lead, and silver sulfates
124
Which halides will NOT dissolve in water?
silver and lead halides
125
Which carbonates will NOT dissolve in water?
All metal carbonates are insoluble
126
Which hydroxides will NOT dissolve in water?
All except group 1 hydroxides and ammonium hydroxide.
127
Define ionisation energy
Energy required to remove an electron from an atom, forming a positive ion
128
Define first ionisation energy
Energy required to remove 1 electron from each atom in 1 mole of gaseous atoms
129
How does atomic radius affect ionisation energy?
- as atomic radius increases, I.E decreases, due to weaker nuclear attraction
130
How does nuclear charge affect ionisation energy?
- greater the nuclear charge, greater nuclear attraction, I.E increases
131
How does electron shielding affect ionisation energy?
- inner shells of electrons repel outer electrons, therefore more inner shell means greater shielding, less nuclear attraction therefore lower I.E
132
Why do successive ionisation energies increase?
As electrons are removed, nuclear charge increases, ionic radius decreases, nuclear attraction increases
133
How can a graph showing successive ionisation energies allow you to predict the group of an element?
- there are large jumps between each energy level
134
Does electronegativity increase or decrease across a period? And why
- increases
| - atomic radius decreases, stronger nuclear attraction
135
Does electronegativity increase or decrease down a group? And why
Decrease
| - atomic radius increases, weaker nuclear attraction
136
Do elements become more or less metallic across a period?
- less metallic
| - reactivity decreases across period
137
What is calcium hydroxide used for?
- neutralising acidic soil
138
What is magnesium hydroxide used for?
Neutralising stomach acid
139
Does thermal decomposition become easier or harder down group and why?
- harder
- smaller ions have greater charge density, therefore greater polarising effect on negatively charged cloud around carbonate ion, therefore carbonate is less stable
140
Acid + alkali =
Salt + water
141
Acid + base =
Salt + water
142
Acid + carbonate =
Salt + water + carbon dioxide
143
Acid + metal =
Salt + hydrogen
144
Acid + metal oxide =
Salt + water
145
What is a salt?
Compound formed when hydrogen ion in an acid is replaced with a metal ion or ammonium ion
146
Give some properties of acids
- turns litmus paper red
| - neutralised by bases
147
What do organic acids contain?
Carboxyl group (COOH)
148
What occurs when an acid and a base react?
They neutralise each other
149
What is titration used for?
To determine concentration of a solvent
150
Give method of volumetric analysis (titration)
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
How are models of the atom made + updated?
Tested using experimental investigations
Are revised when observations are made that aren't predicted by model
152
What is nuclear fission?
The splitting of a large, unstable isotope triggered by bombarding it with smaller, high-speed particles (usually neutrons)
153
What conditions are needed for nuclear fission?
Why?
High temps and/or pressure to provide the energy needed to overcome the repulsion between the 2 positive nuclei
154
What are isotopes?
Atoms of the same element with a different number of neutrons
This causes mass number to be different
155
What is the Avogadro constant (NA)?
The number of atoms/molecules in 1 mole of a substance
156
What does quantised mean?
Energy that can only take particular values (known as quanta)
157
What is the ground state?
The lowest energy level that an electron can occupy
158
What is a photon?
Quanta of energy in the form of electromagnetic radiation
159
Briefly describe Bohr's model of the atom
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
What property does light have?
What does this mean?
Wave-particle duality
Means it can behave like both a wave and a particle...
161
What properties does light have the mean it can be described as a particle?
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
What equation links the wave + particle models of light?
ΔE = hv
```
ΔE = energy of photon (J)
h = Planck's Constant
v = frequency (Hz/s-1)
```
163
What equation explains the wave properties of light?
c = vλ
```
c= speed of light (ms-1)
v = frequency (Hz/s-1)
λ = wavelength (m)
```
164
What is spectroscopy?
The study of how light and matter interact
Uses IR, visible, and UV light
165
Explain the formation of an emission spectrum
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
Why can emission/absorption spectra be used to identify different atoms from a compound/mixture?
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
What colour flame does Li+ give?
Bright red
168
What colour flame does Na+ give?
Yellow
169
What colour flame does K+ give?
Lilac
170
What colour flame does Ca2+ give?
Brick red
171
What colour flame does Ba2+ give?
Apple green
172
What colour flame does Cu2+ give?
Blue-green
173
Describe the appearance of an absorption spectrum
Black lines on a colourful background (showing all colours of visible light)
174
How are atomic absorption spectra formed?
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
What are the similarities between emission and absorption spectra?
For a given element, lines appear at the same frequency
Lines converge at a higher frequency
Several series of lines are seen
176
What are the differences between atomic emission and absorption spectra?
Emission: coloured lines on a black background
Absorption: black lines on a coloured background
177
Why do the lines of emission/absorption spectra get closer together at higher frequencies?
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
Why are several series of lines seen on emission/absorption spectra?
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
What are the rules that determine the distribution of electrons in atomic orbitals?
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
What are the 2 ways of representing electron distribution?
By writing out the electronic configuration in full
e.g. 1s22s22p5
By drawing the electronic configuration in boxes.
181
What trend do melting/boiling points follow across a period?
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
What is the general equation for first ionisation enthalpy?
X(g) → X+(g) + e-
183
What is the trend in first ionisation enthalpies as you go across a period?
Why?
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
What is the trend in first ionisation enthalpy down a group?
Why?
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
What is the trend for atomic radii across a period?
Decreases due to the increased number of protons
This means there is greater attraction between the outer electrons and the nucleus
186
Why are s-block elements more reactive than p-block elements?
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
What is a lone pair?
A pair of electrons in the outer shell of an atom that are not involved in bonding
188
Describe/explain how electron pair repulsion determines the shape of molecules
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
What type of structure do ionic bonds have?
Always giant ionic (lattice)
190
What type of structure do covalent bonds have?
Either simple molecular or giant covalent network
191
What type of structure do metallic bonds have?
Always giant metallic lattice
192
What are the characteristic properties of giant ionic lattices?
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
Describe the structure/bonding in simple molecular covalent bonding
Strong covalent bonds within molecules (between atoms) (strong intramolecular bonds)
But only weak intermolecular bonds between molecules
194
What are the characteristic properties of simple covalent molecules?
Low melting point
Usually insoluble in water
Do not conduct electricity (or heat)
195
What are the characteristic properties of giant covalent networks?
High melting point because all bonds in structure are strong covalent bonds
Insoluble in water
Do not conduct electricity (apart from graphite)
196
Describe the structure/bonding in a giant metallic lattice
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
What are the characteristic properties of giant metallic lattices?
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
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
Octahedral
199
Group 1 metals have relatively low melting points
Use ideas about the charge and size of Group 1 ions to explain this
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
What is a precipitate?
A suspension of solid particles formed by a chemical reaction in solution
201
What is a precipitation reaction?
Reaction between ions in solution that forms a precipitate
| suspension of solid particles/insoluble solid particles
202
Which ionic substances are soluble?
All compounds containing...
Group 1 metals
Nitrate ions
Ammonium ions
... are soluble
203
Which ionic substances are insoluble?
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
The presence of which ions can be tested for by adding barium chloride solution?
(Solution containing Ba2+ ions)
Sulfate (SO42-) ions - white ppt formed
205
Give three examples of practical applications of precipitation reactions
Water treatment
Production of coloured pigments for paints/dyes
Identification of certain metal ions in solutions
206
What is relative isotopic mass?
The mass of one atom of an isotope compared to 1/12 of the mass of a 12C atom
207
What is empirical formula?
The simplest ratio of atoms in a compound
208
What is water of crystallisation?
Number of water molecules contained in an ionic lattice per molecule of salt
209
What is the formula for percentage yield?
(Experimental yield / theoretical yield) x 100
210
Why might percentage yield be lower than expected?
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
What are the 3 main factors that affect 1st ionisation enthalpy?
Atomic radii - larger = lower
Nuclear charge - more protons = higher
Electron shielding - outer shells feel less electrostatic attraction to nucleus
212
How do the hydroxides of Group 2 metals change as you go down the group?
Become increasing soluble and more alkaline
213
What substances do Group 2 metal oxides react with?
What property does this give them?
They react with acids, so can act as bases
Metal oxide + Acid → Salt + Water
214
What happens to Group 2 metal carbonates when they are heated?
Give the general equation
Undergo thermal decomposition
Metal carbonate → Metal oxide + Carbon dioxide
215
How does the thermal stability of Group 2 carbonates change going down the group?
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
Why does the thermal stability of Group 2 carbonates increase as you go down the group?
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
What is charge density?
The charge of an ion relative to its size
218
What is an oxonium ion?
What is its formula?
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
Briefly describe how a soluble salt can be made
By reacting the appropriate acid and alkali together
The solid salt can then be produced by evaporating the excess solution/water
220
Briefly describe how an insoluble salt can be made
By a precipitation reaction
E.g. silver iodide can be made by reacting silver nitrate + potassium iodide
221
Explain how the lines on spectra can be used to identify elements
• 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
Describe the appearance of emission spectra
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)
