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

Module 3.1 - The Periodic Table Flashcards

(58 cards)

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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?

Study These Flashcards
A
  • delocalised electrons can move freely anywhere in lattice

- so metals conduct electricity even as a solid

25
What does ductile mean?
Can be drawn-out or stretched
26
What can metals be used for because of their ductility?
Can be drawn into wires
27
What does malleable mean?
Can be hammed into different shapes
28
What can metals be used for due to their malleability?
Can be pressed into shapes or hammered into thin sheets
29
Why is it difficult to classify silicon as a metal or a non metal?
- shiny appearance like metal but brittle | - conducts electricity but v poorly
30
Describe and explain the trends in melting point across a period.
- 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
What is the structure of graphite?
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
What are the physical properties of group 2 elements?
- reasonably high melting and boiling points - light metals with low densities - form colourless (white) compounds
33
Describe the trend in reactivity of group 2 elements as you go down the group.
- 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
How can group 2 elements be reducing elements?
Are oxidised to 2+ ions M(g) --> M+(g) + e- M+(g) --> M2+(g) + e-
35
How do group 2 elements react with oxygen?
- react vigorously with oxygen - form an ionic oxide, MO - redox e. g. 2Ca(s) + O2(g) --> 2CaO(s)
36
How do group 2 elements react with water?
- 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
How do group 2 elements react with dilute acid?
- 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
How do group 2 oxides react with water?
-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
Describe the solubility of group 2 metal hydroxides with water.
- 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
What can calcium hydroxide be used for?
- neutralise acidic soils | - reduce acidity of soils
41
What can magnesium hydroxide be used for?
-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
What can calcium carbonate be used for?
- 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
What are the properties of the halogens?
- low melting and boiling points | - exist as diatomic molecules
44
What is the trend in boiling point of halogens?
- 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
Describe the reactivity of the halogens.
- 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
What colour is chlorine in water?
pale green
47
What colour is chlorine in cyclohexane?
pale green
48
What colour is bromine in water?
orange
49
What colour is bromine in cyclohexane?
orange
50
What colour is iodine in water?
brown
51
What colour is iodine in cyclohexane?
violet
52
What is meant by the term 'disproportionation'?
The reduction and oxidation of the same element in a redox reaction
53
Describe the reaction between chlorine and water.
-chlorine kills bacteria to make water safe to drink -disproportionation reaction occurs Cl2(aq) + H2O(l) --> HClO(aq) + HCl(aq)
54
Describe the reaction between chlorine and cold dilute sodium hydroxide.
-forms bleach | Cl2(aq) + 2NaOH(aq) --> NaCl(aq) + NaClO(aq) + H2O(l)
55
How do you test for carbonate ions?
-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
How do you test for sulphate ions?
-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
How do you test for halide ions (Cl-, I-, Br-)?
-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
How do you test for ammonium ions?
- 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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