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A Level Chemistry > Periodic Table > Flashcards

Periodic Table Flashcards

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1
Q

What is periodicity?

A

The trends in physical and chemical properties of elements as you go across the periodic table

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2
Q

Which parts of the periodic table are s-block, p-block, d- block and f-block

A

Groups 1 & 2 - s-block, Groups 3,4,5,6,7 and 0 - p-block, transition elements d-block, f-block at the bottom

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3
Q

What is first ionisation energy?

A

The energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms

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4
Q

Is ionisation energy endothermic or exothermic?

A

Endothermic - energy has to be put in

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5
Q

Does a lower ionisation energy make it easier or harder to form an ion?

A

Easier

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6
Q

What affects ionisation energy?

A
  1. Nuclear charge: more protons = more positively charged nucleus = stronger attraction for the electrons
  2. Atomic radius: electron close to nucleus will be more strongly attracted
  3. Shielding: as no. of electrons between outer electrons and nucleus increases the outer electrons feel less attraction towards the nuclear charge
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7
Q

What is the trend in ionisation energy down a group? Why.

A

As you go down ionisation energies generally fall (easier to remove outer electrons). This is because elements further down the group have extra electron shells so have a larger atomic radius and more shielding. Both of these factors make it easier to remove outer electrons, resulting in a lower ionisation energy. Positive charge of nucleus does increase as you go down but this effect is overridden by the effect of the extra shells

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8
Q

What is the trend in ionisation energy as you move across a period? Why?

A

Ionisation energies increase (harder to remove outer electrons). This is because number of protons is increasing, so higher nuclear charge, so smaller atomic radius and a stronger attraction. The extra electrons are roughly at the same energy level so there is generally little extra shielding effect to lessen the attraction from the nucleus

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9
Q

Why is there a drop in ionisation energy between groups 2 and 3?

A

Outer electrons in group 3 elements are in a P-orbital rather than an S-orbital. P orbitals are slightly higher energy than S orbital, so the electron on average, is found further from the nucleus. P orbital also has additional shielding provided by the s electrons. These factors override the effect of the increased nuclear charge

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10
Q

Why is there a drop in ionisation energy between groups 5 and 6?

A

Due to electron repulsion. In group 5 the electron is being removed from a singly-occupied orbital. In group 6 the electron is being removed from an orbital containing 2 electrons. The repulsion between 2 electrons in an orbital means electrons are easier to remove from shared orbitals.

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11
Q

What is second ionisation energy?

A

The energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions

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12
Q

What is a successive ionisation energy?

A

The energy needed to remove 1 mole of each subsequent electron from each ion in 1 mole of positively charged gaseous ions.

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13
Q

How do successive ionisation energies change within an element?

A

Within each shell successive ionisation energies increase. This is because electrons are being removed from an increasingly positive ion (less repulsion amongst remaining electrons, so more energy needed to remove the next electron). Big jumps in ionisation energy happen when a new shell is broken into ( an electron is being removed from a shell loser to the nucleus.

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14
Q

What are allotropes?

A

Different forms of the same element in the same state are called allotropes

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15
Q

What are allotropes?

A

Different forms of the same element in the same state are called allotropes

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16
Q

What are 3 carbon allotropes?

A

Graphite, diamond and graphene

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17
Q

What is the structure of graphite?

A

Sheets of flat hexagons covalently bonded with 3 bonds each. The 4th outer electron of each carbon atom is delocalised. The sheets of hexagons are bonded together by weak induced dipole-dipole forces.

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18
Q

What are 5 properties of graphite?

A
  1. Slippery - weak forces between the layers are easily broken so sheets can slide over each other.
  2. Electric current can flow - delocalised electrons are free to move along the sheets
  3. Low density - layers are quite far apart compared to the length of the covalent bonds
  4. Very high melting point - because of strong covalent bonds
  5. Insoluble - covalent bonds are too difficult to break
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19
Q

What is the structure of diamond?

A

Each carbon atom is covalently bonded to 4 other carbon atoms. Atoms arrange themselves into a tetrahedral shape - crystal lattice structure

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20
Q

What are 5 properties of diamond?

A
  1. Very high melting point- due to strong covalent bonds
  2. Extremely hard - due to strong covalent bonds
  3. Good thermal conductor - vibrations travel easily through the stiff lattice
  4. Can’t conduct electricity - all outer electrons held in localised bonds
  5. Won’t dissolve in any solvent - covalent bonds to difficult to break
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21
Q

Silicon structure

A

Each silicon atom forms 4 strong covalent bonds and forms a crystal lattice structure. Similar properties to diamond

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22
Q

What are the properties of graphene

A

1.Best known electrical conductor - delocalised electrons free to move. Without layers, they can move quickly above and below the sheet
2. Extremely strong - delocalised electrons strengthen covalent bonds
3. Transparent and incredibly light

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23
Q

What is metallic bonding?

A

Outermost electrons of a metal atom are delocalised. This leaves a positive metal cation which is electrostatically attracted to the delocalised electrons

24
Q

How does metallic bonding affect Mp and Bp?

A
  1. Mp & Bp - the more delocalised electrons per atom the stronger the bonding will be & the higher the Mp and Bp. The size of the metal ion and lattice structure also affect Mp and Bp. Smaller ionic radius will hold the delocalised electrons closer to the nuclei
25
Q

How does metallic bonding affect the ability to be shaped?

A

No binds holding specific ions together, metal ions can slide over each other when the structure is pulled, so metals and malleable and ductile

26
Q

What does ductile mean?

A

Can be drawn into a wire

27
Q

How does metallic bonding affect conductivity?

A

Delocalised electrons can pass kinetic energy to each other, making metals good thermal conductors. Delocalised electrons can also carry a charge making them good electrical conductors

28
Q

How does metallic bonding affect solubility?

A

Metals are insoluble, except in liquid metals because of the strength of the metallic bonds

29
Q

Why do MPs and Bps increase from group 1 to group 2?

A

Because the metal-metal bonds get stronger because the metal ions have a greater charge, increasing the number of delocalised electrons and a decreasing ionic radius. This leads to a higher charge density which attracts the ions together more strongly

30
Q

What are the Bp and Mp of simple molecular structures like?

A

Low melting & boiling points because it depends on the strength of the induced dipole-dipole forces between the molecules which are weak and easily overcome. More atoms in a molecule mean stronger induced dipole-dipole forces.

31
Q

Why do noble gases have the lowest melting & boiling points?

A

Because the exist as individual atoms resulting in very weak induced dipole-dipole forces

32
Q

How does reactivity change as you go down group 2? Why?

A

Reactivity increases down the group. Because ionisation energy decreases down the group. It becomes easier to lose electrons making the element more reactive.

33
Q

Group 2 + water -> ??? (What happens to one of the products?)

A

Metal hydroxide + hydrogen (metal hydroxide dissolves in water to produce OH- ions which make the solution strongly alkaline)

34
Q

What happens to ionisation energies as you go down group 2? Why?

A

Decrease because of increasing size of atomic radii and increasing shielding effect

35
Q

What are group 2 elements also known as?

A

Alkaline earth metals

36
Q

What happens when group 2 metals react with oxygen?

A

Burn in oxygen to form solide white oxides

37
Q

Group 2 oxide + water -> ??? (What does the product do? Are there any exceptions? What does this tell you about the pattern?)

A

Metal hydroxides which dissolve, OH- ions make these solutions strongly alkaline. MgO is an exception, it only reacts slowly and the hydroxide isn’t very soluble. Oxides form more strongly alkaline solutions as you go down because hydroxides get more soluble

38
Q

Dilute acids + group 2 element –>

A

salt + hydrogen

39
Q

What are 2 examples of uses of group 2 compounds?

A

Ca(OH)2 : agriculture to neutralise acid soils
Mg(OH)2 & CaCO3 : indigestion tablets as antacids (neutralise excess stomach acid)

40
Q

What are group 7 elements also known as?

A

Halogens

41
Q

What is the colour and physical state of F2, Cl2, Br2, I2

A

F2 : pale yellow gas
Cl2 : green gas
Br2 : red-brown liquid
I2 : grey solid

42
Q

What happens to the Mp and Bp of group 7 elements as you go down the group? Why?

A

Increase - increasing strength of induced dipole-dipole (London) forces (as no. of electrons increases when size and relative mass of the atoms increases

43
Q

Halogen atoms are (reduced/oxidised), when this happens they (reduce/oxidise) another substance so they are (reducing/oxidising) agents

A
  1. Reduced 2. Oxidise 3. Oxidising
44
Q

What happens to reactivity as you go down group 7?

A

Get less reactive as you go down the group because:
1. Atomic radius increases so outer electrons further from nucleus
2. Outer electrons shielded more because there are more inner electrons
So harder for larger atoms to attract electron needed (despite increased charge on nucleus) so larger atoms are less reactive

45
Q

What reactions can show the halogens’ relative oxidising strengths? How does this apply to Cl, Br and I? show ionic equations. What are the colour changes?

A

Displacement reactions with halide ions. A halogen will displace a halide from a solution if the halide is below it in the periodic table.
Chlorine will displace bromide and iodide:
1. Cl2 + 2Br- -> 2Cl- + Br2 : colourless -> orange/red (with organic solvent)
2. Cl2 + 2I- -> 2Cl- + I2 : colourless -> violet/pink (with organic solvent)
Bromine will displace iodide:
1. Br2 + 2I- -> 2Br- + I2 : orange -> violet/pink (with organic solvent)
Iodine will have no reaction with F-, Cl- or Br- : brown
(All ionic equation elements/ions are aq)
Organic solvent: e.g. hexane - halogen present will readily dissolve in organic solvent which settles out as a distinct layer above aqueous solution (very pale yellow/green shows chlorine)
KCl,KBr and KI(aq) used for halide ions

46
Q

How do you identify halogens in a solution?

A

Use halogen displacement reactions: mix unknown solution with some known halogen solutions and see which colour changes take place.
1. Violet/pink shows presence of iodine
2. Orange/red shows presence of bromine
3. Very pale yellow/green shows chlorine

47
Q

How do test for halides?

A
  1. Dissolve solid in distilled water
  2. Add 3cm cubed of solution to separate test tubes using a pipette
  3. Label each test tube
  4. Prepare table to record results
  5. Add few drops of dilute nitric acid (HNO3) using pipette (this removes any unwanted ions)
  6. Add few drops of aqueous silver nitrate (AgNO3) using pipette
  7. If chloride, iodide or bromide ions are present a precipitate of silver halide will be formed. Chloride = white precipitate, Bromide = cream precipitate, Iodide = yellow precipitate
  8. Add dilute ammonia (NH3) using pipette (around same amount as amount of precipitate)
  9. Stir with glass rod, Cl- will dissolve
  10. If doesn’t dissolve add concentrated ammonia (same amount as dilute ammonia was added)
  11. Stir with glass rod Br- will dissolve. I- will still not have dissolved
48
Q

What is disproportionation?

A

When a single element is simultaneously oxidised and reduced.

49
Q

When do halogens undergo disproportionation?

A

Halogens undergo disproportionation when they react with a cold dilute alkali solutions (e.g. NaOH or KOH)

50
Q

What are 2 examples of useful disproportionation reactions?

A
  1. Making bleach: 2NaOH + Cl2 -> NaClO (bleach) + NaCl + H2O, used in water treatment, to bleach paper & cleaning toilets
  2. Chloric(I) acid: Cl2 + H2O + HClO (Chloric(I) acid) this ionises to form chlorate(I) ions (ClO-), these ions kill bacteria so can make water safe to drink or swim in, prevent algae growth, eliminates bad smells and tastes, removes discolouration caused by organic compounds. However risks include: chlorine gas is toxic, liquid chlorine on skin & eyes causes severe chemical burns, chlorine reacts with organic compounds to form chlorinated hydrocarbons many of which are carcinogenic
51
Q

What alternatives to chlorine are there to purify drinking water?

A
  1. Ozone - strong oxidising agent so great at killing microorganisms, expensive to produce, short half-life in water so treatment isn’t permanent
  2. Ultraviolet light - damages DNA of microorganisms, ineffective in cloudy water, won’t stop water being contaminated further down the line
52
Q

To prevent false positives, what order should you carry out ion tests?

A

Carbonates, then sulfates, then halides

53
Q

How do you test for carbonates?

A

Add a dilute strong acid (e.g. dilute nitric of hydrochloric acid), if carbonates present CO2 will be released. Limewater will go cloudy if CO2 produced

54
Q

How do you test for sulphates?

A

Add dilute acid to remove any carbonate or sulfites. Add a few drops of barium nitrate solution (Ba(NO3)2). If white precipitate formed sulphate is present.

55
Q

How do you test for ammonium ions

A

NH3 : Damp red litmus paper: paper will turn blue
NH4 + : add few drops of aqueous sodium hydroxide & warm mixture. Then test litmus paper

A Level Chemistry (8 decks)

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