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Chemistry HL IB 2020-22 > 3.2 Periodic trends > Flashcards

3.2 Periodic trends Flashcards

1
Q

What is a covalent radius?

A

Half the distance between two neighboring nuclei

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2
Q

Does atomic radius decrease or increase across a period and why?

A

It decreases as:

  • number of protons increases=higher nuclear charge
  • number of electrons increase but electron shielding effect remains the same
  • Results in stronger attraction between nucleus and valence electrons so radius decreases
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3
Q

Does atomic radius increase or decrease down a group?

A

It increases as valence electrons occupy a main energy level further from nucleus and so less attraction

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4
Q

Do cations and anions lose or gain electrons?

A

Cations lose electrons while anions gain

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5
Q

Isoelectronic

A

Elements with same electronic configuration eg Na+ and Mg 2+

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6
Q

Does ionic radius increase or decrease across a period? Are there any exceptions?

A

It generally decreases because:

  • Proton number increases while electron number remains the same
  • There is still the same shielding effect so higher effective nuclear charge
  • Stronger attraction to nucleus
  • Exception: When it shifts from cation to anion as electron no.>proton no. so effective nuclear charge decreases and larger shielding effect. Starts decreasing after increase however.
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7
Q

Does ionic radius increase or decrease down a group?

A

Increases as shielding effect increases so there is less attraction.

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8
Q

Are cations bigger or smaller than parent atom?

A

Smaller as they occupy one less main energy level

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9
Q

Are anions bigger or smaller than parent atom?

A

Bigger as electrons are more than protons so more repulsion and hence bigger radii

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10
Q

Effective nuclear charge

A

Strength of attraction between positive nucleus and negative electrons

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11
Q

What does effective nuclear charge depend on?

A
  • Size of atomic radius and number of shielding electrons
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12
Q

What is effective nuclear charge of Sodium and Magnesium?

A

Sodium has ENC of +1 as it has 10 shielding electrons and Magnesium has ENC of +2

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13
Q

Trend of effective nuclear charge down a group

A

Remains constant as increase in protons and hence attraction of nucleus is offset by increase in occupied energy levels

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14
Q

Trend of effective nuclear charge across a period

A

Increases since shielding effect remains the same as it’s the same energy level

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15
Q

What is first ionization energy?

A

Energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions

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16
Q

Equation for first ionization energy

A

X (g)–> X+ (g) + e-

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17
Q

Trend of first ionization energy across a period

A

Increases as:

  • effective nuclear charge increases so electrons are more difficult to remove
  • atomic radius decreases and same shielding effect means more attraction
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18
Q

Trend of first ionization energy down a group

A

Decreases as:

  • atomic radii decreases
  • shielding effect increases reducing effective nuclear charge so less attraction makes it easier to remove
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19
Q

Electron affinity

A

Energy released when one mole of electrons are added to one mole of atoms to form one mole of gaseous 1- ions

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20
Q

Is first electron affinity positive or negative?

A

Negative as energy is released so it’s exothermic

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21
Q

Trend of electron affinity across a period

A

Increases because:
- Shielding effect is same
- Proton number increases
- Effective nuclear charge increases so attraction is stronger and more energy is released
Moving right means likelihood of pairing electrons increases as they try to get a more stable configuration. Metals often have more tendency to lose rather than gain

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22
Q

Trend of electron affinity down a group

A

Decreases as:

  • Electron gained is entering an energy level further (increase in atomic radii)
  • Increase in shielding effect + decrease in effective nuclear charge = weaker attraction so less release of energy
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23
Q

Electronegativity

A

Attraction of an atom for shared pair of electrons in a covalent bond

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24
Q

Which elements have the highest and lowest electronegativity?

A

Fluorine is highest at 4.0 while Francium is lowest at 0.7

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25
Q

Trend of electronegativity across a period

A

Increases because:

  • effective nuclear charge increases and shielding effect remains the same
  • atomic radii also decreases so attraction is stronger between nucleus and bonding electrons
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26
Q

Trend of electronegativity down a group

A

Decreases as attraction is less so it’s harder to gain bonding electrons

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27
Q

Trend of metallic character down a group and across a period

A

Increases as tendency to lose electrons increases while it decreases across a period

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28
Q

Trends of melting point across a period (eg. 3rd period)

A

Strength of metallic bond increases from Na to Al so melting point increases and peaks at Silicon due to giant covalent structure however at non-metals it decreases as intermolecular forces are weak and metallic character decreases

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29
Q

Why does Sulphur have higher melting point than Phosphorus?

A

S has bigger molecules as they exist as S8 so increased bonding = higher M.P.

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30
Q

Why are metals generally solids at STP?

A

They have strong electrostatic attractions

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31
Q

How does bonding depend on electronegativity?

A
  • If there is large difference in electronegativity between elements, it is an ionic compound
  • If there is a small difference in electronegativity, it is usually a covalent bond
32
Q

Product of basic oxides + acid

A

Salt + water

33
Q

Basic oxides + water

A

They dissolve to form basic solutions (hydroxides)

34
Q

Na2O (s) + H2O (l)

A

2NaOH (aq)

35
Q

Product of acidic oxides + bases

A

Salt + water

36
Q

Acidic oxides + waters

A

They dissolve to form acidic solutions

37
Q

P4O10 + 6H2O

A

4H3PO4

38
Q

SO3 + H2O

A

H2SO4

39
Q

3NO2 + H2O

A

2HNO3 + NO

40
Q

Cl2O + H2O

A

2HClO

41
Q

What property does silicon dioxide not share with other acidic oxides?

A

It doesn’t dissolve in water

42
Q

Why is silicon dioxide an acidic oxide?

A

It can react with NaOH to form Na2SiO3 + H2O. i.e Sodium silicate

43
Q

What type of oxide is aluminum oxide and why?

A

An amphoteric oxide because it can react with acids and bases

44
Q

Al2O3 + 2NaOH

A

3H2O + 2NaAl(OH)4 (Sodium aluminate)

45
Q

Al2O3 + 6HCl

A

2AlCl3 + 3H2O

46
Q

SO2 + 2NaOH

A

Na2SO3 + H2O Sodium sulphite

47
Q

SO3 + 2NaOH

A

Na2SO4 + H2O

Sodium sulfate

48
Q

Which nitrogen oxide is acidic?

A

Nitrogen dioxide NO2

49
Q

Which nitrogen oxides are neutral?

A

NO and N2O (Nitric and nitrous oxides)

50
Q

2NO2 + H2O

A

HNO3 + HNO2

51
Q

CO2 + H2O

A

H2CO3 (Carbonic acid)

52
Q

What is the ending when there is only one oxyanion?

A

-ate

53
Q

Ending for 2 oxyanions

A

> O2= -ate

54
Q

Ending for 4 oxyanions (Example: Chlorine)

A

Smallest no. of oxygens = hypo- -ite
2nd= -ite
3rd= -ate
Greatest no. = per- -ate

(Hypochlorite, chlorite, chlorate, perchlorate)

55
Q

Properties of alkali metals

A
  • Highly reactive
  • Soft
  • Low melting points
  • Increasing densities
56
Q

Why do alkali metals have similar chemical properties?

A

They all have one valence electron

57
Q

What are alkali metals held by?

A

Electrostatic attractions between positive ions in the lattice and delocalized electrons

58
Q

Why do metals have decreasing melting points as they go down the group?

A

Attractions become weaker as atomic radius increases, effective nuclear charge decreases so less energy is required to break the bonds

59
Q

Why are alkali metals highly reactive?

A

They only have one electron to lose which means lower ionization energies

60
Q

Metal + water

A

Metal hydroxide + hydrogen

61
Q

Describe reaction of sodium, potassium and caesium with water

A
  • Sodium: Melts, fizzes rapidly and moves around
  • Potassium: Bursts into flames (lilac)
  • Caesium: Explodes as it comes in contact
62
Q

Color and state at RTP of fluorine

A

Pale yellow gas

63
Q

Color and state at RTP of chlorine

A

Yellow-green gas

64
Q

Color and state at RTP of bromine

A

Deep-red/orange liquid

65
Q

Color and state at RTP of iodine

A

Grey, shiny solid (usually sublimates), + purple as vapors

66
Q

Why does melting point of halogens increase down the group?

A

As molar mass increases, London forces get stronger so ionization energy increases

67
Q

How many electrons do halogens have and what can they form from it?

A
  • They can gain an electron to become stable and form ionic compounds
  • Form covalent compounds
68
Q

Does reactivity increase or decrease down a halogen group?

A

Decreases

69
Q

What do you get from an alkali metal + halogen

A

White, colorless salts

70
Q

Halogens + water

A

Soluble in water and form colorless, neutral solutions

71
Q

Most and least vigorous reactions of alkali metal + halogen

A
  • Most vigorous: Caesium + Fluorine

- Less vigorous: Lithium + Iodine

72
Q

When can halogens be displaced in a solution and why?

A

When they react with a halogen that is more reactive as they have higher electron affinity (higher in the group)

73
Q

Reaction between chlorine and potassium bromide (Give ionic as well)

A

Cl2 + 2KBr –> 2KCl + Br2

Ionic: Cl2 + 2Br- –> 2Cl- + Br2

74
Q

Iodine + Potassium Bromide

A

No reaction

75
Q

What type of reactions are halogen-halide ion?

A

Redox because halogens are strong oxidizing agents. They oxidize a less reactive halide ion

76
Q

What determines physical properties?

A

Intermolecular forces

Chemistry HL IB 2020-22 (19 decks)

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