Periodic Table Flashcards
(177 cards)
1
Q
What is the modern periodic table arranged in?
A
By order of proton number
2
Q
What are Döbereiner triads?
A
The groups that were made when, in 1817, Johann Döbereiner attempted to group similar elements
3
Q
What did the English chemist, John Newlands, notice when her arranged the elements in order of mass?
A
That elements with similar chemical and physical properties occurred at regular intervals, but every 8th element was different
4
Q
What did Newlands call his discovery?
A
The law of octaves
5
Q
What does the Periodic law state?
A
If you arrange the elements in order of increasing atomic number then their chemical and physical properties will repeat in a systematic way
6
Q
What was different about Mendeleev’s table?
A
He left gaps, so that elements with similar properties were in the same group
7
Q
Give an example of a case where 2 elements in the periodic table are the wrong way around?
A
Ar —-> K are not in order of atomic mass
8
Q
All the elements within the same period have the same number of…
A
Shells
9
Q
What is periodicity?
A
A repeating pattern of physical and chemical properties across a period
10
Q
Why is the sub-shell 4s before 3d?
A
The 4s sub-shell has a lower energy level
11
Q
What is the electron configuration of cobalt (27 electrons)
A
1s2 2s2 2p6 3s2 3p6 4s2 3d7
12
Q
What is the first ionisation energy?
A
The energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms
13
Q
What type of process is the first ionisation energy?
A
Endothermic, as you have to put energy in to remove the 1st electron
14
Q
Write out the 1st ionisation energy equation of oxygen
A
O —–> O+ + E-
15
Q
Why do you have to use the gas state symbol when writing out the equation for the first ionisation energy?
A
Because ionisation energies are measured for gaseous atoms
16
Q
The lower the ionisation energy…
A
The easier it is to form an ion
17
Q
What is an endothermic process?
A
One that takes in heat
18
Q
What does a high ionisation energy mean?
A
There’s a strong electrostatic attraction between the electron and nucleus, so more energy is needed to remove the first electron
19
Q
What 3 things affect ionisation energy?
A
Nuclear charge
Atomic radius
Shielding
20
Q
How does nuclear charge affect ionisation energy?
A
The more protons there are in the nucleus, the more positively charged it is and so the stronger the attraction for the electrons
21
Q
How does atomic radius affect ionisation energy?
A
Attraction falls off rapidly with distance, so an electron closer to the nucleus will be more strongly attracted than one that’s further away
22
Q
How does shielding affect ionisation energy?
A
As the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less attraction
23
Q
What is the trend in ionisation energy down a group?
A
As you go down a group, ionisation energies decrease (as it gets easier to remove an electron)
24
Q
Why does ionisation energy decreases down a group?
A
Elements further down a group have extra electron shells, so the atomic radius is larger (reducing attraction to the nucleus). Shielding also increases
25
What is the trend of ionisation energies as you go along a period?
They increase (gets harder to remove the outer electrons)
26
Why do the ionisation energies increase as you go along a period?
Number of protons increases, meaning a larger nuclear charge, so a smaller atomic radii and a stronger attraction
27
Is there any shielding change as you go along a period?
No, as all the extra electrons are at roughly the same energy level
28
As you go along a period, why is there a drop of ionisation energy between Group 2 and 3?
Due to the outer electrons in Group 3 being in a p-orbital rather than an s-orbital
29
As you go along a period, why is there a drop of ionisation energy between Group 5 and 6?
Due to electron repulsion
30
Why does electron repulsion cause a drop of ionisation energy between Group 5 and 6?
For the Group 5, the outer electron is being removed from an orbital containing 1 electron, whereas for the Group 6, the outer electron is being taken from an orbital containing 2 electrons (which repel each other). This repulsion makes it easier to remove an electron
31
What is successive ionisation energy?
The energy needed to remove 1 mole of each subsequent electron from each ion in 1 mole of positively charged gaseous ions
32
What is the equation for the second ionisation of Oxygen?
O+ ----> O2+ + e-
33
Explain the trend in successive ionisation energies
Within each shell, successive ionisation energies increase because electrons are being removed from an increasingly positively charged ion (so there's less repulsion between the remaining electrons). There are big jumps in ionisation energy every time a new shell is broken into
34
What are macromolecular structures?
Another name for giant covalent lattices
35
Give 2 examples of elements that can form macromolecular structures
Carbon and Silicon
36
Why can carbon and silicon form giant covalent structures?
They both form 4 strong, covalent bonds
37
What is an allotrope?
A different form of the same element in the same state
38
What are the 3 carbon allotropes?
Graphite
Diamond
Graphene
39
How are the carbon atoms arranged in graphite?
Sheets of flat hexagons covalently bonded with 3 bonds each (the 4th outer electron of each carbon atom is delocalised)
40
What are the sheets of hexagons in graphite bonded together by?
Weak induced dipole-dipole forces
41
What does delocalised mean?
An electron isn't attached to a particular atom
42
Why is graphite used as a dry lubricant?
The weak forces between layers can easily be broken, so the sheets can slide over each other
43
Can graphite conduct electricity?
Yes, as the delocalised electrons are free to move along the sheets
44
Describe the density of graphite?
A low density because the layers are far away from each other compared to the length of the covalent bonds
45
Does graphite have a high or low melting point?
High melting point (around 3700K) because of the strong covalent bonds
46
Is graphite soluble or insoluble?
Insoluble in any solvent, as the covalent bonds in the sheets are too difficult to break
47
What is the shape formed by the carbon atoms in diamond?
Tetrahedral, as each carbon is covalently bonded to 4 others
48
Describe the melting/boiling point of Diamond
Very high melting point
49
Describe the strength of Diamond
Extremely hard (which is why it's used on the tips of diamond-tipped drills)
50
Describe the thermal conductivity of Diamond
It's a good thermal conductor because vibrations travel easily through it
51
Can Diamond conduct electricity?
No because all the outer electrons are held in localised bonds
52
Does Diamond dissolve in solvent?
No
53
What are macroscopic properties?
Properties that can directly be perceived
54
What are the properties of Silicon similar to?
That of diamond
55
What is graphene?
One layer of graphite (a sheet of carbon atoms in hexagonal shapes)
56
Why is graphene a 2D compound?
It's one atom thick
57
Why is graphene thought of as the best electrical conductor?
The delocalised electrons are free to move, but because graphene doesn't have any layers they can move quickly above and below the sheet making good at conducting electricity
58
Describe the strength of graphene
Very strong because the delocalised electrons strengthen the covalent bonds between the covalent bonds
59
Give 2 uses of graphene
High-speed electronics
| Aircraft technology
60
What properties make graphene useful in high-speed electronics?
It's high strength, low mass and good electrical conductivity
61
Why is graphene used in the screens of phones?
It's flexible and transparent
62
What is metallic bonding?
A lattice of metal cations in a sea of delocalised electrons
63
What 2 things affect the melting/boiling points of metals
The number of delocalised electrons per atom
| Size of metal ion and lattice structure
64
How does the number of delocalised electrons per atom affect the melting/boiling points of metals?
The more there are, the stronger the metallic bonds and so the higher the melting and boiling point is
65
Which has a higher melting point between Na+ and Mg2+?
Mg2+ as it has 2 delocalised electrons per atom compared to Na+ having only 1
66
How does the size of the ion and lattice structure affect the boiling/melting points of metals?
A smaller ionic radius will hold the delocalised electrons closer to the nucleus (meaning a larger melting/boiling point)
67
Are metals malleable or non-malleable?
They're malleable because there an no bonds holding specific ions together, and so the metal ions can slide over each other when the structure is pulled
68
Describe the thermal conductivity of metals
They are good thermal conductors because the delocalised electrons can pass kinetic energy to each other
69
Describe the electrical conductivity of metals
They are good electrical conductors because the delocalised electrons can carry a current
70
Describe the solubility of metals
They are insoluble (except in liquid metals)
71
Why are metals insoluble?
Because of the strength of the metallic bonds
72
What is charge density?
Amount of charge in relation to the size of the ion
73
Describe the trends of boiling/melting points across Periods 2 and 3
A general increase between the 1st and 4th elements in the period, but then decrease from 4th to the 8th
74
Explain the trends in boiling/melting points across Periods 2 and 3
As you go along the period, the types of structure changes from giant metallic to giant covalent and then to a simple molecular lattice
75
For metals across a period, why does the boiling/melting point increase?
The metal-metal bonds get stronger because the metal ions have a greater charge, an increasing number of delocalised electrons and a smaller atomic radii
76
Which has a greater charge density between a small and a large ions with the same charge?
The small ion
77
Why do B, C and Si have the highest melting and boiling point in their period?
They have strong covalent bonds linking their atoms together, which need a lot of energy to break
78
Why do simple molecular structures have low melting and boiling points?
The melting and boiling points depend on the strength of the London forces (not how strong the covalent bonds are), so these forces are weak and are easily overcome
79
More........in a molecules mean stronger London forces
Atoms
80
Why do the noble gases have the lowest boiling and melting point?
They are monatomic, so the London forces are very weak
81
What block of the periodic table are the alkaline earth metals in?
S-block
82
Where are the alkaline earth metals found?
Group 2
83
What is the trend in reactivity as you go down Group 2?
Increases
84
Why does the reactivity increase down Group 2?
When Group 2 elements react, they lose electrons forming positive ions (cations). The easier it is to remove an electron, the more reactive the element. So because ionisation energy decreases down a group, the reactivity increases
85
What is the electronic structure of Mg?
1s2 2s2 2p6 3s2
86
What is the electronic structure of Mg2+?
1s2 2s2 2p6
87
Write an ionic equation for when Mg reacts
Mg ------> Mg2+ + 2e-
88
What is the product when Group 2 elements react with water?
Metal Hydroxide and Hydrogen
89
Write the equation for the reaction of Mg with H2O
Mg + 2H2O -----> Mg(OH)2 + H2
90
What is the pH of the reaction of a Group 2 element with water?
pH 12 or 13, as the metal hydroxide that forms dissolves in water releasing Hydroxide ions (making it strongly alkaline)
91
Why do ionisation energies decrease as you go down a Group?
Because of the increasing size of the atomic radii and the increasing shielding effect
92
What is the general formula of the reaction between the Group 2 metal 'M' and Water?
M + 2H20 ----> Ca2+ + 2OH- + H2
93
What is the general formula of the reaction between the Group 2 metal 'M' and Oxygen?
2M + 02 -----> 2MO
94
What is the product when Group 2 metals burn in oxygen?
A solid white oxide
95
Why are Group 2 metals known as Alkaline Earth Metals?
Because their oxides form alkaline solutions (pH12-13)
96
What is the product of the reaction between the oxides of Group 2 metals and water?
They form metal hydroxides, which are very alkaline and dissolve
97
What is an alkali?
A base that is soluble in water
98
Why is Magnesium Oxide an exception to the other oxides of Group 2 metals?
It only reacts slowly and the Hydroxide isn't very soluble
99
What is the trend down Group 2 for the strength of the alkaline solutions formed?
As you go down the group, the oxides form more strongly alkaline solutions
100
Why do the oxides form more strongly alkaline solutions as you go down Group 2?
Because the hydroxides get more soluble
101
What is the general formula of the reaction between the Group 2 metal 'M' and dilute acids (this case being HCl)?
M + 2HCl ----> MCl2 + H2
102
What are many compounds of group 2 metals used for?
For neutralising acids
103
What is CaOH used for?
In agriculture to neutralise acid soils
104
What are Mg(OH)2 and CaCO3 used for?
Used in indigestion tablets (used to neutralise stomach acids)
105
What do halogens exist as?
Diatomic molecules
106
What colour is fluorine at room temperature?
Pale yellow
107
What colour is chlorine at room temperature?
Green
108
What colour is bromine at room temperature?
Red-Brown
109
What colour is iodine at room temperature?
Grey
110
What state is fluorine at room temperature?
Gas
111
What state is chlorine at room temperature?
Gas
112
What state is bromine at room temperature?
Liquid
113
What state is iodine at room temperature?
Solid
114
What is the trend in boiling/melting points you go down Group 7?
Increase
115
Why does the boiling/melting point increase as you go down Group 7?
Due to the increasing strength of the London forces, as the number of electrons increases when the size and relative mass of the atom increases
116
Why does the states change from gas to solid as you go down the top 4 elements in Group 7?
Because of the increasing boiling point
117
What does it mean if a substance is volatile?
It has a low boiling point
118
Are halogen atoms oxidised or reduced?
Reduced
119
What forms when halogens are reduced?
Halide ions
120
Why are halogens oxidising agents?
Because when they are reduced, they oxidise another substance
121
What is the trend in atomic radii as you go down Group 7?
The atomic radii increases, so the outer electrons become further away from the nucleus
122
What is the trend in reactivity as you go down Group 7?
They become less reactive
123
Why does the reactivity decreases as you go down Group 7?
The atomic radii increases so the outer electrons get further away from the nucleus. The outer electrons are also shielded more from the attraction of the nucleus, because there are more inner electrons
124
What 2 ions will Chlorine displace?
Br- and I-
125
What ion will Bromine displace?
I-
126
What can you add to a Group 7 displacement reaction to make the colour change stand out more?
Any organic compound such as Hexane
127
What colour is iodine in the presence of an organic compound?
A violet/pink
128
What colour is bromine in the presence of an organic compound?
Orange/red
129
What colour is chlorine in the presence of an organic compound?
Very pale yellow/green
130
What are the 6 steps for testing for halide ions?
- If you're testing solid substances, dissolve each of them in separate water
- Use a pipette to add about 3cm^3 of each solution to 3 different test tubes
- Label each test tube
- Prepare a table to record your results, leaving room to record what precipitate forms
- Use a pipette to add a few drops of nitric acid to remove unwanted ions
- Add a few drops of silver nitrate solution and observe the colour of the precipitate formed to see which ions are present
131
What colour is the precipitate when silver nitrate is added to chloride ions?
A white precipitate
132
What colour is the precipitate when silver nitrate is added to bromide ions?
A cream precipitate
133
What colour is the precipitate when silver nitrate is added to iodide ions?
A yellow precipitate
134
Describe the 5 steps of the further test if you can't tell which ion is present
- Using a pipette, add dilute ammonia to each of the silver halide solutions
- Stir the mixture using a glass rod and record whether any precipitates dissolve in the ammonia
- If they haven't dissolved, add concentrated ammonia to the solution
- Stir the mixture using a glass rod and record whether any of the precipitates dissolve
135
What happens to the silver chloride precipitate when dilute ammonia is added?
Precipitate dissolves
136
What happens to the silver bromide precipitate when dilute ammonia is added?
It doesn't dissolve
137
What happens to the silver bromide precipitate when concentrated ammonia is added?
Precipitate dissolves
138
What happens to the silver iodide precipitate when dilute ammonia or concentrated ammonia is added?
It doesn't dissolve
139
What is disproportionation?
When a single element is oxidised and reduced simultaneously
140
The halogens undergo disproportionation when they react with what?
Cold dilute alkali solutions such as KOH or NaOH
141
What is the ionic equation of the disproportionation reaction between a halogen (X) and NaOH?
X2 + 2OH- ----> XO- + X- + H2O
142
In the reaction X2 + 2OH- ----> XO- + X- + H2O (where X is a halogen) what is the oxidation number of X in XO-?
+1
143
What is the full equation of the disproportionation reaction between a halogen (X) and NaOH?
X2 + 2NaOH ----> NaXO + NaX + H2O
144
What is Sodium Chlorate (I) used for?
Household bleach
145
How do you form Sodium Chlorate (I)?
Mix chlorine gas and cold, dilute aqueous sodium chloride
146
What is the full equation for the reaction to make bleach?
2NaOH(aq) + Cl2(g) ------> NaClO(aq) + NaCl(aq) + H2O(l
147
What is chlorine's oxidation state in NaClO?
As ClO- is a chlorate ion, the chlorine has a +1 oxidation number
148
What type of reaction is it when you react chlorine and water?
A disproportionation reaction
149
What are the 2 products when you react chlorine and water?
Hydrochloric acid and Chloric (I) acid
150
Why is chlorine used water treatment?
It kills disease causing microorganisms as well as preventing algae growth and removes discolouration
151
What is the main risk of using chlorine in water treatment?
Chlorine is very toxic
152
What is the affect on your health if you breath in chlorine?
It irritates the respiratory system
153
What are the health risks with liquid chlorine?
Liquid chlorine on the skin or in eyes causes severe chemical burns
154
When chlorine reacts with the organic compounds in the water, it forms chlorinated hydrocarbons. What are the health risks with this?
The chlorinated hydrocarbons are carcinogenic (cancer causing)
155
Compare the health risks of treated water and untreated water
The risk of causing cancer is small compared to a cholera epidemic from untreated water
156
What are 2 alternatives to chlorine for treating water?
Ozone
| UV light
157
Name an advantage of using ozone in water treatment
It's a strong oxidising agent, meaning it's great at killing microorganisms
158
What are the 2 disadvantages of using ozone in water treatment?
It's very expensive to produce
| It's short half-life in water means the treatment isn't permanent
159
How does UV light kill microorganisms in water treatment?
It damages their DNA
160
What is the chemical formula of ozone?
O3
161
What is a false positive result?
A result that suggests an ion is present, when it actually isn't
162
How do false positive results occur when testing for ions?
If there are ions present that will interfere in the test
163
What order should you test for halides, carbonates and sulfates to prevent false positive results?
Carbonates ---> Sulfates ---> Halides
164
How do you test for carbonates?
Add a dilute strong acid (e.g. nitric acid) to unknown solution. If carbonates are present, CO2 will be released (which turns limewater cloudy)
165
What is the ionic equation for the reaction between carbonates and acid?
CO3 2- + 2H+ ----> CO2 + H2O
166
How do you test for sulfates?
Add a few drops of Barium Nitrate (Ba(NO3)2) to a solution of the unknown substance. If you get a white precipitate, then it's Barium Sulfate so it contains sulfate ions
167
Are sulfates soluble or insoluble in water?
Soluble, except for Barium Sulfate
168
What reaction do you have to avoid when testing for sulfates?
Sodium Sulfide reacting with Barium Nitrate solution
169
Why do you have to avoid the reaction of Sodium Sulfide reacting with Barium Nitrate solution when testing for sulfates?
It forms Barium Sulfite and Sodium Nitrate. Barium Sulfite is also a white precipitate, which will confuse results for test for sulfates
170
When testing for carbonates, why do you add a dilute strong acid?
The acid will react with any carbonates and sulfites present, making sure they don't interfere with the test for sulfates and cause a false positive result
171
What is the ionic equation for the test for sulfates?
SO4 - + Ba 2+ ---> BaSO4
Sulfate Anion + Barium Cation ---> Barium Sulfate
172
Why do you test for halides after sulfates?
Because silver sulfate is insoluble. If you tested for halides before sulfates, then any precipitate could be a false result as it could be silver sulfate rather than silver halide
173
How do you test for halides?
Add nitric acid, and then silver nitrate solution
174
What will happen in the test if chloride, bromide or iodide are present?
A precipitate will form:
```
Chloride = White
Bromide = Cream
Iodide = Yellow
```
175
How do you test of ammonium ions?
Use a damp piece of red litmus paper. If there's ammonia present, the paper will turn blue
176
Ammonium Chloride + Sodium Hydroxide ---> ?
Ammonia + Water + Sodium Chloride
177
When testing for ammonium ions, why does the litmus paper have to be damp?
So the ammonium gas can dissolve and make the colour change
