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Chemistry A-Level > Module 3 - Periodic Table and Energy > Flashcards

Module 3 - Periodic Table and Energy Flashcards

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1
Q

How is the periodic table organised by atomic number?

A

The elements are arranged in order of increasing atomic number. Each successive element has atoms with one extra proton.

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2
Q

How is the periodic table organised by groups

A

Arranged in vertical columns by groups. Each elements in a group has atoms with the same number of outer shell electrons and similar properties.

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3
Q

How is the periodic table organised by periods and periodicity

A

The elements are arranged in horizontal rows called periods. The number of the period gives the number of the highest energy electron shell in an element’s atoms.

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4
Q

What is periodicity

A

Repeating trend in properties of the elements across a period

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5
Q

What change happens across each period?

A

Change from metals to non-metals

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6
Q

Where are s-block elements

A

In groups 1 and 2

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7
Q

Where are p-block elements

A

Groups 3 to 0

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8
Q

Where are d-block elements

A

Transition metals

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9
Q

What are f-block elements

A

Lanthanides and actinides

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10
Q

What is the electron configuration trend across period 2

A

The 2s sub-shell fills with two electrons, followed by the 2p sub-shell with six electrons

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11
Q

What is the electron configuration trend across period 3

A

Filling for the 3s and 3p sub-shells

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12
Q

What is first ionisation energy

A

The minimum energy required to remove one mole of electrons from one mole of atoms in a gaseous state. It is measured in kJmol-1

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13
Q

Equation for first ionisation energy of sodium

A

Na(g) > Na+(g) + e-

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14
Q

What is successive ionisation energy

A

When further electrons are removed after the first ionisation energy

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15
Q

Why do successive ionisation energies cause the ionisation energy to increase

A

As electrons are removed the electrostatic force of attraction between the positive nucleus and the negative outer electron increases. Therefore more energy is needed to overcome this attraction.

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16
Q

How does the atomic radius of an atom affect ionisation energy

A

The greater the distance between the nucleus and the outer electrons the less the nuclear attraction. The force of attraction falls off sharply with increasing distance.

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17
Q

How does nuclear charge affect ionisation energy

A

The more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons. Therefore more energy is needed to overcome the attraction.

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18
Q

How does electron shielding affect ionisation energy

A

More energy levels = more shielding between nucleus and outer electrons.
This means less energy is needed to remove an electron.

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19
Q

What is the general trend in ionisation energy as you go down a group

A

As you go down the energy required to remove an electron decreases as it is further from the nucleus and there is more electron shielding

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20
Q

What is the general trend in ionisation energy as you go across a period

A

-IE increases as the electron being removed is attracted more strongly by the nucleus as it has extra protons but the distance and shielding have both remained the same
-Slight drop in atomic radius as we go across the period due to increasing nuclear charge

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21
Q

What causes the exception in ionisation energies between Mg > Al (Period 3)

A

-In Al and Mg you are removing an electron from a p-sub shell in Al and an s-sub shell in Mg
-It is easier to remove an electron from the p-sub shell
-This is because the 3s sub shell is shielding the p-electron and the p-sub shell is further away so there is less attraction

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22
Q

Why is there an exception in ionisation energy trends between N > O (Period 2)

A

Nitrogen - 1s2 2s2 2p3
Oxygen - 1s2, 2s2, 2p4
The repulsion between the two electrons in the same sub-orbital means that the electron is easier to remove than it would be otherwise be so less energy is required

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23
Q

Why is there an exception in ionisation energy trends between Be > B (Period 2)

A

Boron has an electron that is being removed from a higher energy level compared to Be that is further from the nucleus so the electron is easier to remove

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24
Q

Why is there an exception to the ionisation energy rule between P > S (Period 3)

A

Sulphur has repulsion within the 3p orbital when two electrons with opposite spins are placed in the same orbital, destabilising the atom and allowing the electron to be removed more easily.

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25
Outline metallic bonding
-Consists of a lattice of positively charged ions surrounded by a sea of delocalised electrons. There are very strong electrostatic forces of attraction between the oppositely charged particles. -The electrons are mobile however the cations are in a fixed positing to maintain the structure of the metal
26
How does the charge of a metal affect its bonding
The greater the charge, stronger the attractive forces as more electrons are released into the ‘sea.’
27
How does the size of a metal ion affect the metallic bonding
Ions that are larger in size produce a weaker attraction due to their greater atomic radius, therefore decreasing the charge density.
28
Outline the conductivity of metals
-Good conductors as the sea of delocalised electrons allow them to carry charge
29
Outline the melting and boiling points of metals
-Most have high MPs and BPs -They depend on the strength of the metallic bonds holding together the atoms. -For most metals high temperatures are needed to overcome the strong electrostatic forces of attraction.
30
Outline all properties of metals
-High MP and BPs -Good electrical conductors -Malleability -Ductility
31
What is a ductile metal?
Metal can be stretched
32
What is a malleable metal?
The metal can be shaped into different forms
33
Outline the solubility of metals
-Do not dissolve
34
What are giant covalent lattices?
Non-metals boron, carbon and silicon have different structures. Many atoms are held by a network of strong covalent bonds
35
Outline diamond and silicon structure
-They are in group 4 so outer shells have four electrons. -They use these electrons to form covalent bonds to other carbon or silicon atoms. -This results in a tetrahedral structure with bond angles of 109.5°
36
Outline melting and boiling points for giant covalent structures
-High as covalent bonds are strong so high temperatures are necessary to provide the large quantity of energy needed to break them.
37
Outline solubility of giant covalent structures
-Insoluble to almost all solvents -The covalent bonds holding together the atoms in the lattice are far too strong to be broken by interaction with solvents.
38
Outline electrical conductivity in giant covalent structures
-Non-conductors of electricity apart from graphite and graphene
39
Outline graphene
-Single layer of graphite composed of hexagonally carbon atoms linked by covalent bonds -Delocalised electrons are used to allow this giant covalent structure to conduct electricity -120° bond angle
40
Outline graphite
-Parallel layers of graphene which are bonded by weak London Forces -Spare delocalised electrons allow conduction of electricity
41
Outline trend of melting point across period 2 and 3
-Increaes from group 1-4 (giant structures) and drastically decreases between 4-5 and stays low 5-0 (simple molecules)
42
Outline group 2 metals as a reducing agent
-Each metal atom is oxidised losing two electrons to form a 2+ ion -Another species will gain these two electrons and be reduced
43
Outline group 2 reactions with oxygen
-Group 2 metals react with oxygen to form oxide in a redox reaction. It is a vigorous reaction. -Strontium and barium can react with excess oxygen and energy to form metal peroxides.
44
Outline group 2 reactions with water
-A redox reaction to produce a metal hydroxide and hydrogen. The metal hydroxide forms an alkaline solution. -Reaction becomes more vigorous with metals further down the group
45
Outline group 2 reactions with dilute acids
Metal + acid > salt + hydrogen -Reactivity increases down the group
46
How does atomic radius change down group 2
Increases down group 2 due to additional electron shells
47
Outline how reactivity changes down Group 2
-Increased electron shielding and atomic radius makes the outer electrons easier to lose. -Reactivity increases down the group
48
Outline how ionisation energy changes down group 2
-Decrease down the group due to a greater atomic radius and increased amounts of shielding. This makes it easier for an electron to be removed.
49
How do group 2 OXIDES react with water
-Releases hydroxide ions and forms alkaline solutions of the metal hydroxide -The group 2 hydroxides are only slightly soluble in water. When the solution becomes saturated any further metal and hydroxide ions form a solid precipitate
50
How does solubility of group 2 HYDROXIDES change down the group
-Increases down the group = more OH- ions and are more alkaline -Mg(OH)2 is only slightly soluble -Ba(OH)2 is much more soluble in water
51
How does the solubility of group 2 SULFATES change down the group
-Decreases in solubility down the group
52
How is Calcium hydroxide used
-Added to field to increase pH of acidic soils -Neutralises acid in soil forming neutral water
53
Outline the use of barium sulphate
-Useful in medicine on barium meals. It allows internal tissues and organs to be imaged. It is safe as it cannot be absorbed into the blood (as it is insoluble)
54
How is magnesium hydroxide used
-Used in medicine as antacids for treating acid indigestion so can neutralise stomach acid
55
What are group 7 elements
-Highly reactive non-metals. They exist as diatomic molecules with single covalent bonds -In order to achieve a full outer shell the halogen gains an electron and forms a 1- ion
56
How is barium chloride used?
Test for sulphate ions as it reacts to form barium sulphate which forms as a white precipitate when sulphate ions are present
57
Outline how atomic radius changes down group 7
Increases down the group due to additional electron shells
58
Outline how electronegativity changes down group 7
-Atomic radius and electron shielding increases therefore electrons in the outer shells are less strongly attracted to the nucleus, and so are more easily removed. -Therefore electronegativity decreases down group 7
59
Outline trend in melting and boiling points of group 7
-Simple covalent molecules with weak van der waals forces. The strength of the IMFs increase down the group as there are more electrons. -Therefore more energy is needed to overcome them
60
How does fluorine exist at RTP
Pale yellow gas F2
61
How does chlorine exist at RTP
Pale green gas Cl2
62
How does bromine exist at RTP
Red-brown liquid Br2
63
How does iodine exist at RTP
Shiny grey-black solid I2
64
Outline halogens as an oxidising agent
-Each halogen atom is reduced, gaining one electron to form a 1- halide ion. -It oxidises another species
65
Outline trend in reactivity of halogens
Halogens need to gain an electron in order to react, as atomic radius increases this becomes harder as the positive attraction of the nucleus is weakened by additional shielding. Therefore reactivity decreases down the grown as it is harder to attract an electron.
66
Outline halogen-halide displacement reactions
-A solution of each halogen is added to aqueous solutions of the other halides. -If the halogen added is more reactive than the halide present a reaction takes place, the halogen displacing the halide and the solution changes colour
67
What are the colours of the halogens in solution in water
Cl2 is pale green Br2 is orange I2 is brown
68
In halogen-halide displacement reactions what can be added to further tell results apart?
Cyclohexane is an organic non-polar solvent that can be added. The non-polar halogens dissolve more readily in cyclohexane than in water.
69
What colours do the halogens turn in solution in cyclohexane
Cl2 stays pale green Br2 stays orange I2 turns violet
70
Which halides will react with Cl2
-Both Br- and I- as Cl2 is more reactive
71
Which halides will react with Br2
I- as Br2 is more reactive
72
Which halides will react with I2
None
73
What is a disproportionation reaction
Redox reaction in which the same element is both oxidised and reduced
74
Outline halogen disproportionation reaction between chlorine and water
-When small amounts of chlorine are added to water, one chlorine atom is oxidised and the other is reduced Cl2(aq) + H2O(l) > HClO(aq) + HCl(aq) -The bacteria is killed by chloric acid and ClO- ions
75
How can you test that Chloric acid acts as a weak bleach
Add indicator to a solution of chlorine in water. The indicator turns red from the presence of two acids and then the colour disappears as the bleaching action of chloric acid takes effect.
76
Outline halogen disproportionation reaction between chlorine and cold dilute NaOH(aq)
chlorine and water reaction is limited by the low solubility of chlorine in water. If the water contains dissolved NaOH, much more chlorine dissolves and another disproportionation reaction takes place. Cl2(aq) + 2NaOH(aq) > NaClO(aq) + NaCl(aq) + H2O(l) -This results in large concentration of ClO- ions from the NaClO formed. This is used as a household bleach.
77
Outline the reaction is chlorine is reacted with hot, concentrated alkali
-Chlorine is disproportionated even further to form one species with an oxidation number of -1 and another with an oxidation number of +5 3Cl2 + 6NaOH > NaClO3 + 5NaCl + 3H2O
78
Outline the benefits of using chlorine
-Kills bacteria so is a disinfectant and is used to treat drinking water and swimming pool water -If we didn’t treat water with chlorine it would be compromised by diseases such as typhoid and cholera
79
Outline the risk of using chlorine
-Toxic gas. -In small concentrations is respiratory irritant and can be fatal in larger concentrations -In water chlorine can react with organic hydrocarbons forming chlorinated hydrocarbons which are suspected of causing cancer
80
How does chlorine react with group 1 and 2 metals
-Form metal chloride which are all white precipitates
81
How are halide ions tested for
-Aqueous silver nitrate is added to form precipitates of silver halides
82
What colour are the silver halides
AgCl - White precipitate AgBr - Cream precipitate AgI - Yellow precipitate
83
How can you further test for halide ions after adding silver nitrate?
It may not always be clear to distinguish the colour of the precipitate so they can be tested further using NH3
84
How do the silver halides react to dilute NH3 being added
AgCl- precipitate dissolves AgBr- No change AgI- No change
85
How do the silver halides react to concentrated NH3 being added
AgCl- precipitate dissolves AgBr- Precipitate dissolves AgI- No change
86
How do hydrogen halides react with ammonia gas
Form ammonia salts. The hydrogen halides are strong acids in solution and react with ammonia in an acid-base reaction to form a salt.
87
How do hydrogen halides react with water
-They form dilute acids -In solution these strong acids dissociate to release their halide ions and hydrogens ions. The hydrogen ions form a hydroxonium ion with H2O molecules e.g. HCl + H2O > Cl- + H3O+
88
How do you test for carbonate anions
-Carbonates react with acids to form CO2 1. In a test tube, add dilute nitric acid to the solid or solution to be tested 2. If you see bubbles the unknown compound could be a carbonate 3. To prove that the gas is CO2 bubble the gas through limewater Ca(OH)2 and the lime water turn cloudy as a white precipitate of CaCO3 is formed.
89
How do you test for sulphate ions
-BaSO4 is very insoluble and forms a white precipitate. Barium ions are therefore added into a solution of unknown compound as either barium chloride or barium nitrate. (Use barium nitrate for halide test)
90
How do you test for halides?
-Silver halides are insoluble 1. Add AgNO3 to an aqueous solution of a halide 2. The silver halide precipitates are different colours 3. Add aqueous ammonia to test the solubility of the precipitates
91
Outline results of AgNO3 + halides and their solubilities in NH3 (reminder)
-AgCl is white and is soluble in dilute NH3 -AgBr is cream and is soluble in concentrated NH3 -AgI is yellow and insoluble in concentrated NH3
92
What is the correct sequence of tests for anions
1. Carbonate (CO3)2- 2. Sulphate (SO4)2- 3. Halides Cl-, Br-, I-
93
Why should you test for carbonates first in the anions tests
-You add a dilute acid and are looking for CO2 -Neither sulphate nor halide ions produce bubbles with dilute acid, so the test can be carried out without possibility of an incorrect conclusion, if the test produces no bubbles no carbonate us present and you can proceed to the next test.
94
What should you test for sulphates second in the anions test
-You add a solution containing Ba2+ ions and are looking for white precipitate of BaSO4 -BaCO3 is white and insoluble in water so if you carry out this test on a carbonate you will form a white precipitate, therefore you need to do the carbonate test first and proceed to this test when you know no carbonate is present.
95
Why is it important to test for halides last in the anions tests
-Silver carbonate (Ag2CO3) and Ag2SO4 are insoluble in water and will form a precipitate in this test. Therefore you need to have previously ruled out any carbonates or sulphate to get only correct halides with a positive result.
96
If you are using a mixture of chemicals how do you carry out the anions test
-In the carbonate test if you see bubbles add HNO3 till the bubbles stop (HNO3 so you don’t use H2SO4 or HCL) -In sulphate test add excess of Ba(NO3)2 and any sulphate ions present will precipitate as BA(SO4). Filter solution to remove the Ba(SO4) -In the halide test any carbonate or sulphate ions have been removed so any precipitate formed must be a halide
97
Outline test for ammonium ions (NH4+)
-When heated together NH4+ and OH- ions react to form NH3 gas NH4+(aq) + OH-(aq) > NH3(g) + H2O(l) 1. NaOH(aq) is added to solution of ammonium ion 2. NH3(g) produced, unlikely to see bubbles as is very soluble in water 3. Mixture is warmed and NH3 gas is released 4. To test presence use moist indicator paper which turns blue in NH3 presence

Chemistry A-Level (9 decks)

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