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Chemistry A level > Periodic table > Flashcards

Periodic table Flashcards

1
Q

Definition of a group?

A

Vertical column in the periodic table. Elements in the same group have similar chemical properties and their atoms have the same number of outer electron shells

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2
Q

Definition of a period?

A

A horizontal row in the periodic table, Elements show trends in properties across a period

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3
Q

What are group 1 metals also known as?

A

Alkali metals

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4
Q

What are the group 2 metals also known as?

A

Alkaline Earth metals

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5
Q

Which is the S block

A

Block on left

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6
Q

Which is the D block?

A

Block in centre

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7
Q

Which is the P block?

A

Block on right

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8
Q

Which is the F block?

A

Block at bottom

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9
Q

What are the group 7 elements also known as?

A

The halogens

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10
Q

What are the group 8 elements also known as?

A

The noble gasses

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11
Q

Why do elements in a period show a trend in physical and chemical properties?

A

Due to different electronic configurations

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12
Q

What’s the regular repeating pattern in properties of the elements across different periods known as?

A

Periodicity

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13
Q

What’s the difference between a subshell and a orbital?

A

Subshell is a group of the same type of orbitals, where as an orbitals are regions of an atom containing up to 2 electrons with opposite spin

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14
Q

Definition of first ionisation energy?

A

The energy required to remove one electron from each atom in one mole gaseous atoms, to form 1 mole of gaseous +1 ions

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15
Q

Algebraic formula for first ionisation energy?

A

X(g) = X+(g) + e-

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16
Q

Describe the general trend in the first ionisation energy change across a period?

A

It increases

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17
Q

How does the first ionisation energy change down a group?

A

It decreases

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18
Q

Why do the first ionisation energies drop sharply when you go down a group?

A

The electrons are being taken away from a new outer shell, meaning the shell is further away from the nucleus, and experiences more shielding. Despite the fact the nuclear charge has increased, the nuclear attraction is still weaker, therefore less energy required to remove the outer electron

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19
Q

4 factors that must be used every time you explain an ionisation energy question?

A

Nuclear charge (amount of protons)
Sheilding (number of shells)
Atomic radius
Nuclear attraction

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20
Q

Write general trend in first ionisation energies across a period?

A

As you go across a period, the nuclear charge increases, and the atomic radius decreases.
As the electrons are in the same shell they experience the same shielding by inner electrons.
So the overall effect is an increased nuclear attraction for the outer electron as you go across a period, therefore more energy is required to overcome this attraction, so the first ionisation energy increases

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21
Q

Which are the points which don’t follow the general trend in first ionisation energy going across a period?

A

Between group 2 and 3

Between group 5 and 6

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22
Q

Explain the difference in first ionisaiton energy between Be (900) and B (801)? (group 2 and 3)

A

The fall between groups 2 and 3 marks the filling of the p subshell. In Be the outer electron is in a 2s subshell, where as in B the outer electron is in a 2p orbital
The 2p orbital in B is of higher energy, therefore experiences more shielding from inner electrons and ones from 2s subshell

So even though the nuclear charge increases by 1, the difference in shielding and energy means the nuclear attraction is weaker in B, so less energy is required to remove it’s outer electron

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23
Q

Explain the difference in first ionisation energy between Nitrogen (1402) and Oxygen (1314)? (group 5 and 6)

A

The fall between groups 5 and 6 marks the start of electron pairing in the p orbitals in the p subshell. In N the outer electron us unpaired in a 2p orbital, in O the outer electron is paired in a 2p orbital.

The paired electrons in oxygen repel slightly, making it easier to remove and electron than in Nitrogen

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24
Q

3 types of bonding elements in periods 2 and 3 could have?

A

Giant metallic lattice, giant covalent lattice, simple molecular structures

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25
What's the 4th type of bonding which can't occur in periods 2 and 3?
Ionic bonding, requires 2 elements to form cations and anions
26
Type of structure in all period 2 and 3 elements?
Giant metallic structure Li, Be Na, Mg, Al Giant covalent structure: B,C Si Simple molecular structure: N2,O2,F2,Ne P4,S8,Cl2,Ar
27
Describe the trend in boiling points from group 1 to 8?
Increases from 1 to 4, but then decreases with groups 5 and 8 being much lower
28
Why do group 4 elements have highest melting points?
They have giant covalent structures, in which each element is covalently bonded to 4 other atoms In order to melt them all 4 of these covalent bonds need to be broken requiring a lot of energy. Higher than Boron as it only makes 3 covalent bonds
29
Why do elements with diatomic molecules or single atoms have low melting/boiling points?
They only have weak intermolecular forces (London forces) holding the molecules together. Which can easily be broken The strength of London forces relates to the number of electrons in the instantaneous dipole-induced dipole interactions, so the smaller the molecule, the lower the melting/boiling point
30
Why does the melting point increase from group 1 to 2 to 3?
Metallic bonding increases in strength with the number of electrons that can be present in the delocalised sea of electrons, and the charge of the metal ion. So in group 1 there will be +1 ions, where as in group 3 there are +3 ions, and there will be more delocalised electrons per ion in the structure. So there will be an increase in electrostatic attraction between the positive ions and sea of electrons, so more energy required to overcome these interactions
31
Why can metals conduct electricity?
Delocalised electrons can carry charge
32
Why does electrical conductivity increase from group 1 to 3?
Increased amount of delocalised electrons in sea of electrons
33
Why do elements with molecular structures have poor electrical conductivity?
Have no mobile charge carriers
34
What is unusual about Carbon considering electrical conductivity?
Graphite is ion of the allotropes of carbon, it has a giant covalent structure but still can conduct electricity This occurs due to in graphite inly 3 of the 4 outer shell electrons are involved in covalent bonding, so the 4th electron is free to move and carry charge
35
How does the atomic radius change as you go across a period?
General decrease ( period 3 elements have larger atomic radius than period 2 elements)
36
Explain the decrease in atomic radius as you go across a period?
The electrons are all in the same quantum shell, with very little increase of electron shielding, however there is a regular increase in nuclear charge. The overall effect is that the outer electrons are pulled more strongly towards the nucleus, reducing the atomic radius
37
How does atomic radius repeated word change going down a group?
It increases, as you go down the group there will be an increasing amount of shells of electrons, with each shell getting larger and the outer electrons being shielded by more inner electron shells, increasing the atomic radius, as the nuclear attraction is weaker
38
Formula for a group 2 metal with Oxygen when M= the group 2 metal?
2M(s) + O2 (g) = 2MO (s) | Redox reaction
39
Formula for a group 2 metal with water when M= the group 2 metal?
M(s) + 2H2O (l) = M(OH)2 (aq) + H2 (g) | Effervesence and metal disappears/dissolves
40
Formula for a group 2 metal with an acid when M= the group 2 metal?
M(s) + 2HNO3 (aq) = M(NO3)2(aq) + H2 (g) | Effervesence and metal disappears/dissolves
41
Reaction of group 2 metal Oxide with water, and observations?
MO(s) + H2O (l) = 2M(OH) (aq) | At first the metal oxide dissolves, then a white solid forms, the solution formed will be alkaline
42
Explain the traffic light reaction?
Water added to metal oxide, a metal hydroxide is formed which initially dissolved in the water But then cloudy white solid formed as Mg2+ (aq) + 2OH- (aq) = Mg(OH)2 (s) is formed Universal indicator added, turns purple as there is a very high concentration OH- ions (alkaline) Then an acid is added drop by drop, and the solution changed from purple to blue, to green, to yellow to red, as it became more acidic The tube was then shaken, causing the indicator to go back to purple, this is because the shaking caused some of the solid Mg(OH)2 to redissolve, increasing the concentration of OH- ions
43
Definition of solubility?
A measure of the ability of a solute to dissolve in a solvent (measured in moldm^3, shows the maximum amount of moles of a solute which will dissolve in 1dm^3 of water)
44
Definition of an alkali?
A type of base that dissolves in water forming hydroxide ions, OH-
45
Definition of alkaline?
A term used to describe a solution containing OH- ions
46
Definition of alkalinity?
A measure of the concentration of OH- ions present in an alkaline solution
47
Definition of pH?
A measure of the concentration of H+ ions present in a solution
48
What's the trend in solubility as you go down the group 2 metals?
It increases (more soluble), and therefore more alkaline, as there more OH-ions are present in the solution
49
Chlorine's symbol, state at rtp, appearance at rtp and harzards?
Cl2, gas, green, toxic
50
Bromine's symbol, state at rtp, appearance at rtp and harzards?
Br2, Liquid, Brown, toxic and corrosive
51
Iodine's symbol, state at rtp, appearance at rtp and harzards?
I2, solid, Grey, toxic vapour formed when heated
52
What happens to the boiling points as you go down group 7?
Boiling point increases
53
What's the relationship between boiling point, and the energy required to boil a substance?
The higher the boiling point, the more energy required
54
What type of structure and bonding do halogen elements have?
Covalently bonded diatomic molecules, with simple molecular structure
55
What type of attraction exists between halogen molecules?
They are non polar, and can't form hydrogen bonds, so the only forces are intermolecular London forces
56
Why as you go down the group do the boiling points increase?
As you go down the molecules increase in size, meaning there are an increased amount of electrons which are further from the nucleus, meaning instantaneous dipoles form more easily and the dipoles are larger, so the London forces are stronger. On boiling these stronger forces require more energy to break, resulting in higher boiling points as you move down the group
57
What is the trend in atomic radius as you go down group 7?
Atomic radius increases due to increasing numbers of shells of electrons
58
Which halogen molecule out of Cl2, Br2 and I2 has the strongest covalent bond and why?
Chlorine, the covalent bond is formed by a pair of electrons between the atoms attracting the nuclei of these atoms, the closer the electrons are to the nucleus, the stronger the attraction. Also smaller atoms have a more condensed electron cloud, with a higher charge density, leading to stronger bonding
59
Compare the size of halogen atoms compared to their ions?
Negative halide ions will have a larger radius, because they've accepted a negative electron which will lead to a greater repulsion between the electrons
60
Colour of chlorine in water and in Hexane?
Water: Pale green/colourless Hexane: Pale green
61
Colour of Bromine in water and in Hexane?
Water: orange Hexane: orange
62
Colour of Iodine in water and in hexane?
Water: Orange Hexane: purple
63
Colour of halide ion in water and in Hexane?
Colourless for both
64
Why do two layers form when hexane is mixed with a solution containing a halogen?
Hexane is non polar, so only forms London forces between molecules, where as water forms hydrogen bonds between molecules So the layers are immiscible
65
Which layer is hexane when hexane is mixed with a solution containing a halogen?
Top layer, it's least dense
66
Why does the colour of the organic layer become more vivid when the halogen containing solution and it are shaken?
The halogen solutions are much more soluble in the organic layer, shaking assists them moving to that layer, so their concentration in the top layer will increase with more shaking
67
Why does hexane need to be added to identify the halogen in solution?
Iodine and Bromine are the same colour in water
68
When added to the halide solution which one of the halogen solutions always took part in the reaction?
Chlorine, it's the most reactive
69
Write the full equation of chlorine reacting with potassium bromide?
Cl2 (aq) + 2KBr (aq) = Br2 (aq) + 2KCL (aq) Displacement reaction
70
Write the ionic equation of chlorine reacting with potassium bromide?
Cl2 (aq) + 2Br- (aq) = 2Cl- (aq) + Br2 (aq)
71
Explain why chlorine out of Chlorine, Bromine and Iodine is the most reactive?
Because chlorine's nucleus is closer to the electrons it gains, and they experience less shielding therefore despite having the lowest nuclear charge the nuclear attraction to out electron is the greatest
72
Which halide ion is most likely to LOSE an electron?
An iodide ion, as the nucleus is further from the outer electrons and they experience more shielding, therefore despite having the largest nuclear charge, the nuclear attraction on the outer electron is the weakest and requires the least energy to overcome
73
What's a disproportionation reaction?
The oxidation and reduction of the same element in a redox reaction
74
disproportionation reaction of chlorine with water?
Chorine + water= Chloric acid + hydrochloric acid Cl2(g) + H2O (l) = HClO (aq) + HCl (aq)
75
Why when universal indicator is added to the products of chlorine + water does it go red and then colourless?
H+ ions are present in solution so it's acidic turning it red, but then the ClO- ions bleach the solution turning it colourless
76
Why is chlorine added to drinking water?
Kills bacteria
77
Why might the addition of chlorine to water cause concern?
Chlorine can react with hydrocarbons, such as those present in decaying vegetation, producing products with suspected carcinogens
78
Show the reaction between Chlorine and cold dilute Sodium Hydroxide?
Chlorine + Sodium Hydroxide = Sodium Chlorate + Sodium Chloride + water Cl2(g) + NaOH (aq) = NaClO (aq) + NaCl (aq) + H2O (l) Disproportionation reaction
79
Why is Sodium chlorate a more effective bleach than chloric acid?
Doesn't decompose as readily
80
What happens when Silver chloride is reacted with aqueous silver nitrate, then dilute ammonia solution, and then concentrated ammonia solution?
White precipitate, dissolves, dissolves
81
What happens when Silver bromide is reacted with aqueous silver nitrate, then dilute ammonia solution, and then concentrated ammonia solution?
Cream precipitate, doesn't dissolve, dissolves
82
What happens when Silver Iodide is reacted with aqueous silver nitrate, then dilute ammonia solution, and then concentrated ammonia solution?
Yellow precipitate, doesn't dissolve, doesn't dissolve

Chemistry A level (16 decks)

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