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Chemistry Year 12/13 > Module 3.1 - The Periodic Table > Flashcards

Module 3.1 - The Periodic Table Flashcards

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1
Q

What is meant by the term ‘periodicity’?

A

Arrangements of the elements in periods showing repeating trends in physical and chemical properties

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2
Q

How are the elements arranged in the periodic table?

A

By increasing atomic numbers

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3
Q

What is each horizontal row of the periodic table called?

A

period

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4
Q

What is each vertical column of the periodic table called?

A

group

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5
Q

Why do elements of the same group have similar properties?

A

Have the same number of outer shell electrons (repeating pattern of electron configuration) (have the same type of orbitals)

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6
Q

What does a higher principal quantum number show?

A

Higher energy level and shell further from the nucleus

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7
Q

Compare when the 4s orbital fills and empties with that of the 3p orbital.

A
  • 4s energy level is lower than 3d energy level
  • 4s orbital fills before 3d orbital
  • 4s orbital would be emptied before 3d orbital during ionisation
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8
Q

What is meant by the term ‘first ionisation energy’?

A

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

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9
Q

What is the equation for the first ionisation energy of neon?

A

Ne(g) –> Ne+(g) + e-

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10
Q

How is an atom oxidised to 1+ ions?

A

Energy supplied to overcome attraction of positive nucleus to outer electrons (outer electrons as have least attraction so least ionisation energy)

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11
Q

What does nuclear attraction (and therefore ionisation energy) depend on and why?

A
  • atomic radius (larger radius = less attraction to outermost electrons as further away)
  • nuclear charge (high charge = greater nuclear attraction)
  • electron shielding (inner shell repel outer shell as all -ve, repelling effect = electron shielding, more inner shells = larger shielding effect so smaller nuclear attraction by outermost electrons)
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12
Q

What is successive ionisation energy?

A

Values that measure the energy required to remove each electron in turn

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13
Q

What is the trend in successive ionisation energies and why?

A
  • each successive ionisation energy is higher than the one before
  • as each electron removed, less repulsion between remaining electrons and each shell drawn slightly closer to nucleus
  • distance of each electron from nucleus decreases slightly so nuclear attraction increases so more energy to remove each successive electron
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14
Q

How do successive ionisation energies provide evidence for shells?

A

Large increases in ionisation energy show next shell in as atomic radius decreases massively

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15
Q

Why do noble gases have a higher first ionisation energy than the rest of the atoms in their period?

A

Full outer shell of electrons and high positive attraction from nucleus so ionisation energy values are large

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16
Q

What are the trend across the periodic table that can affect ionisation energy?

A

-increasing number of protons so higher attraction to electrons
-same shell so outer shell drawn inwards slightly
-same number of inner shells, so electron shielding hardly changes
∴ attraction between nucleus and outer electrons increases so more energy needed to remove an electron so 1st ionisation energy increases across a period

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17
Q

Why is there a small decrease in first ionisation energy between group 2 and 13 elements?

A
  • group 13’s outermost electron in p orbital but group 2’s in s orbital
  • p orbitals have slightly higher energy than s orbital so slightly further from nucleus
  • electrons in theres orbitals are slightly easier to remove so elements have lower 1st ionisation energy
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18
Q

Why is there a small decrease in first ionisation energy between group 15 and 16 elements?

A
  • electrons in p orbitals
  • 13, 14, 15: each p orbital only has a single electron
  • 16: electron now paired in a p orbital
  • spin-paired electrons have some repulsion so slightly easier to remove so lower 1st ionisation energy
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19
Q

Why is there a sharp decrease in first ionisation energy between the noble gas if one period and the group 1 element of the next period?

A
  • new shell added so outermost electrons further from nucleus
  • increase in distance of outermost shell from then nucleus
  • increase in electron shielding of outermost shell by inner shells
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20
Q

What is the trend of first ionisation energy down a group and why?

A
  • 1st ionisation energy decreases
  • no. of shells increases so outer electrons further from nucleus so weaker force of attraction on outer electrons
  • more inner shells so shielding effect on outer electrons from nuclear charge increases so weaker attraction
  • nuclear charge (no. of protons) increases but increased attraction outweighed by increasing distance/shielding
  • attraction (nucleus and outermost electron) decreases so less energy to remove electron so lower 1 IE (causes larger atomic radius as electrons not pulled as close to nucleus)
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21
Q

What is the structure of a metallic lattice?

A
  • metal cations in fixed position in lattice
  • outer shell electrons are delocalised (sea of delocalised electrons) (shared between all atoms in metallic structure) and spread throughout structure
  • electrons can move in structure (can’t match electron to cation it came from)
  • charge balanced over whole structure
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22
Q

What are the properties of giant metallic lattices?

A
  • high melting point and boiling point
  • good electrical conductivity
  • malleability and ductility (delocalised electrons give a degree of give)
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23
Q

Why do metals have high melting and boiling points?

A
  • electrons free to move throughout structures but +ve ions remain where they are
  • attraction between +ve ions and -ve delocalised electrons is v strong
  • high temp needed to overcome metallic bonds and dislodge ions from their rigid position in lattice
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24
Q

Why are metals good electrical conductors?

A
  • delocalised electrons can move freely anywhere in lattice

- so metals conduct electricity even as a solid

25
Q

What does ductile mean?

A

Can be drawn-out or stretched

26
Q

What can metals be used for because of their ductility?

A

Can be drawn into wires

27
Q

What does malleable mean?

A

Can be hammed into different shapes

28
Q

What can metals be used for due to their malleability?

A

Can be pressed into shapes or hammered into thin sheets

29
Q

Why is it difficult to classify silicon as a metal or a non metal?

A
  • shiny appearance like metal but brittle

- conducts electricity but v poorly

30
Q

Describe and explain the trends in melting point across a period.

A
  • 1-14: mp increases steadily as elements have giant structure. Each successive group w metallic lattice mp increases as nuclear charge/no. of outer shell electrons increases so stronger attraction. If giant covalent lattice, each successive group has more electrons to form covalent bonds
  • 14-15: sharp decrease in mp as elements have simple molecular structure so only need to overcome weak intermolecular forces
  • 15-18: mp remain low as all simple structures
31
Q

What is the structure of graphite?

A

Forms 2 dimensional giant lattice, 1 C atom thick, of interlocking hexagonal rings. V strong and light and can conduct electricity. Forms layers which are graphite

32
Q

What are the physical properties of group 2 elements?

A
  • reasonably high melting and boiling points
  • light metals with low densities
  • form colourless (white) compounds
33
Q

Describe the trend in reactivity of group 2 elements as you go down the group.

A
  • reactivity increases
  • each successive element has its outer electrons in a higher energy level so has larger atomic radius so more shielding from positive nucleus
34
Q

How can group 2 elements be reducing elements?

A

Are oxidised to 2+ ions
M(g) –> M+(g) + e-
M+(g) –> M2+(g) + e-

35
Q

How do group 2 elements react with oxygen?

A
  • react vigorously with oxygen
  • form an ionic oxide, MO
  • redox
    e. g. 2Ca(s) + O2(g) –> 2CaO(s)
36
Q

How do group 2 elements react with water?

A
  • all (except Be) react to form hydroxides, M(OH)2 and H2 gas
  • Mg reacts slowly; further down group metal reacts more vigorously w water
  • redox
    e. g. Ca(s) + 2H2O(l) –> Ca(OH)2(aq) + H2(g)
37
Q

How do group 2 elements react with dilute acid?

A
  • all (except Be) react to form salt and H2 gas
  • more vigorous down group
    e. g. Ca(s) + 2HCl(aq) –> CaCl2(aq) + H2(g)
38
Q

How do group 2 oxides react with water?

A

-form metal hydroxides
MO(s) + H2O(l) –> M(OH)2(aq)
-metal hydroxides are soluble in water and form alkaline solutions by releasing OH- ions

39
Q

Describe the solubility of group 2 metal hydroxides with water.

A
  • solubility of hydroxides increases down the group
  • more soluble = more OH- released = more alkaline
  • Be: BeO insoluble in water
  • Mg forms Mg(OH)2(s) that’s only slightly soluble in water; resulting solution id dilute w comparatively low OH- conc
  • Ba(OH)2(s) is much more soluble in water than Mg(OH)2 so more OH- conc so more alkaline
40
Q

What can calcium hydroxide be used for?

A
  • neutralise acidic soils

- reduce acidity of soils

41
Q

What can magnesium hydroxide be used for?

A

-indigestion from build up of stomach acid (HCl)
-‘milk of magnesia’ w Mg(OH)2 to neutralise excess stomach acid
Mg(OH)2 + 2HCl –> MgCl2 + 2H2O

42
Q

What can calcium carbonate be used for?

A
  • useful building material
  • present in limestone and marble
  • used in manufacture of glass and steel
  • but reacts readily w acids
    e. g. CaCO3(s) + 2HCl(aq) –> CaCl2(aq) + H2O(l) + CO2(g)
  • most rainwater has an acidic pH causing gradual erosion of objects made using limestone or marble e.g. buildings or statues
43
Q

What are the properties of the halogens?

A
  • low melting and boiling points

- exist as diatomic molecules

44
Q

What is the trend in boiling point of halogens?

A
  • as you go down the group, bp increases and physical state changes from gas to liquid to solid
  • each successive element has extra shells of electrons so higher level of London forces between molecules
45
Q

Describe the reactivity of the halogens.

A
  • v reactive and v electronegative
  • strong oxidising agents (take elections)
  • atomic radius increases (nuclear full further away from incoming electrons)
  • electron shielding increases
  • ability to gain an electron to form 1- ion decreases
  • more reactive halogens displace a less reactive halogen (halogen displaced can be checked)
46
Q

What colour is chlorine in water?

A

pale green

47
Q

What colour is chlorine in cyclohexane?

A

pale green

48
Q

What colour is bromine in water?

A

orange

49
Q

What colour is bromine in cyclohexane?

A

orange

50
Q

What colour is iodine in water?

A

brown

51
Q

What colour is iodine in cyclohexane?

A

violet

52
Q

What is meant by the term ‘disproportionation’?

A

The reduction and oxidation of the same element in a redox reaction

53
Q

Describe the reaction between chlorine and water.

A

-chlorine kills bacteria to make water safe to drink
-disproportionation reaction occurs
Cl2(aq) + H2O(l) –> HClO(aq) + HCl(aq)

54
Q

Describe the reaction between chlorine and cold dilute sodium hydroxide.

A

-forms bleach

Cl2(aq) + 2NaOH(aq) –> NaCl(aq) + NaClO(aq) + H2O(l)

55
Q

How do you test for carbonate ions?

A

-add dilute strong acid to suspected carbonate
-collect any gas formed and pass through limewater
-fizzing/colourless gas produced
-gas turns limewater cloudy
CO3 2-(aq) + 2H+(aq) –> H2O(l) + CO2(g)

56
Q

How do you test for sulphate ions?

A

-add dilute hydrochloric acid and add barium chloride to suspected sulphate
-white ppt of barium sulphate is produced
Ba2+(aq) + SO4 2-(aq) –> BaSO4(aq)

57
Q

How do you test for halide ions (Cl-, I-, Br-)?

A

-dissolve suspected halide in water
-add aqueous solution of silver nitrate
-note colour of any ppt formed
-if colour hard to distinguish, add aqueous ammonia (first dilute then concentrated)
-note solubility of ppt in aqueous ammonia
-silver chloride: white ppt, soluble in dilute NH3
-silver bromide: cream ppt, soluble in conc NH3 only
-silver iodide: yellow ppt, insoluble in dilute and conc NH3
Ag+(aq) + X-(aq) –> AgX(s)

58
Q

How do you test for ammonium ions?

A
  • add sodium hydroxide solution to suspected ammonium compound and warm v gently
  • test any gas evolved w red litmus paper
  • ammonia gas turns red litmus paper blue
  • ammonia gas has a v distinctive smell

Chemistry Year 12/13 (18 decks)

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