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Chemistry > The Periodic Table > Flashcards

The Periodic Table Flashcards

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1
Q

Antoine-Laurent de Lavoisier’s periodic table

A
  • distinguished between metals and non metals
  • included some compounds and mixtures
  • included terms such as light which he believed to be substances
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2
Q

Berzelius’ periodic table

A
  • published table of atomic weights
  • determined the composition by mass of many compounds
  • introduced letter based symbols for elements
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3
Q

Dobereiner’s periodic table

A
  • noticed certain groups of 3 elements (triads) ordered by atomic weight would have a middle element with a middle weight and propoerties that were roughly an average of the other two.
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4
Q

Newlands’ periodic table

A
  • elements arranged in order of relative atomic weights

- suggested elements show similar properties to the element 8 places after it (law of octaves)

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5
Q

Mendeleev’s periodic table

A
  • elements ordered by atomic weight
  • elements arranged in vertical columns
  • gaps left where no element fitted repeating pattern, later the missing elements have been found to match Mendeleev’s predictions
  • arrangement of elements in order of atomic weights corresponded to their distinctive chemical properties
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6
Q

Moseley’s input to the periodic table

A
  • determined the atomic number for all known elements
  • elements vary periodically with atomic numbers, rather than atomic weight
  • this corrected the order of some elements
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7
Q

Seaborg’s input to the periodic table

A
  • discovered the transuranic elements

- placed the actinide series below the lanthanide series at the bottom of the table

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8
Q

How to shorten an electronic configuration (examples for Li, Na and K)

A

The inner shell configuration is based on the noble gas that comes before the element in the periodic table.v
Li: 1s2 2s1 or [He] 2s1
Na: 1s2 2s2 2p6 3s1 or [Ne] 3s1
K: 1s2 2s2 2p6 3s1 3p6 4s1 or [Ar] 4s1

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9
Q

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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10
Q

Equation for the first ionisation energy of sodium

A

Na(g) –> Na+(g) + e-

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11
Q

Factors affecting ionisation energy

A
  1. Atomic radius - the larger the atomic radius the smaller the nuclear attraction on outer electrons
  2. Nuclear charge - the higher the nuclear charge, the larger the attractive force on outer electrons
  3. Electron shielding - inner shells of electrons repel the outer shells, the more inner shells the larger the shielding and the smaller the nuclear attraction on the outer electrons
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12
Q

Equation for second ionisation energy of lithium

A

Li+(g) –> Li2+(g) + e-

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13
Q

Why is each successive ionisation energy higher than the one before?

A
  • As each electron is removed there is less repulsion between remaining electrons so each shell will be drawn closer to the nucleus
  • Positive nuclear charge outweighs the negative charge each time an electron is removed
  • As distance of outer electrons decreases, nuclear attraction increases so more energy is required to remove each electron
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14
Q

What is the trend in ionisation across a period and why?

A

The attraction between the nucleus and outer electrons increases across a period so more energy is needed to remove the outer electron.

  • number of protons in nucleus increases so there is a higher attraction on electrons as increased nuclear charge pulls the electrons closer to the nucleus. This factor has the biggest effect on ionisation energy
  • electrons added to same shell, outer shell is drawn inwards slightly
  • same number of shells so shielding doesn’t change
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15
Q

Why is there a decrease in ionisation energy between groups 2 and 13?

A

Group 13 elements have their outermost electron in a p-orbital whereas group 2 elements have outer electron in s-orbital. p-orbitals have a slightly higher energy than s-orbitals so a slightly further from the nucleus so electrons in p-orbitals are easier to remove.

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16
Q

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

A

In groups 13, 14 and 15, each of the p-orbitals contain a single electron. In group 16, the outer electron is now spin-paired in the first p-orbital. Electrons that are spin paired experience some repulsion which makes the outer electron easier to remove.

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17
Q

Why is there a sharp decrease in ionisation energy between the end of one period and the beginning of the next?

A

Another shell has been added which is further from the nucleus which leads to an increase in distance of the outer electron from the nucleus and an increase in electron shielding by inner shells.

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18
Q

What is the trend in ionisation energy moving down a group and why?

A

Moving down a group, first ionisation energy decreases.

  • number of shells increases so the distance of outer electrons to nucleus increases so there is a weaker force of attraction on outer electrons
  • more inner shells so shielding effect increases, this leads to a weaker attraction on outer electron.

Even though the number of protons increases, the increased attraction is outweighed by the increased nuclear radius and shielding.

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19
Q

Metallic bonding

A
  • Positive ions occupy fixed positions in lattice
  • outer shell electrons are delocalised and are shared between all atoms in metallic structure
  • metal held together by strong attractions between positive ions and negative electrons
20
Q

In a giant metallic lattice…

A
  • delocalised electrons spread throughout the structure and can move within the structure
  • over the whole structure the charges must balance

A giant metallic lattice is described as a lattice of positive ions in fixed positions surrounded by a sea of delocalised electrons.

21
Q

Why do metals have high melting and boiling points?

A
  • electrons free to move throughout structure but positive ions remain where they are
  • attraction between positive ions and negative delocalised electrons is very strong
  • High temperature needed to overcome metallic bonds and dislodge ions from rigid positions
22
Q

Why are metals good electrical conductors?

A
  • delocalised electrons can move freely anywhere in the metallic lattice
23
Q

What does it mean to be malleable and ductile and why do metals have these properties?

A

Ductile - can be drawn out or stretched, allows metals to be drawn into wires

Malleable - can be hammered into different shapes

Delocalised electrons allow the metallic structure to have a degree of ‘give’ which allows atoms and layers to slide past each other. This allows metals to have these properties.

24
Q

What is the trend in melting points across periods 2 and 3 for groups 1 - 14?

A
  • Between groups 1 and 14, melting points steadily increase because the elements have giant structures. - - The elements with giant metallic lattices (Li and Be for period 2 and Na, Mg and Al for period 3) the nuclear charge and number of outer shell electrons increases. This causes a stronger attraction.
  • Elements with a giant covalent lattice (B and C for period 2 and Si for period 3). Each succesive group has more electrons with which to form more covalent bonds
25
Q

What is the trend in melting points across periods 2 and 3 for groups 14 - 15?

A

There is a sharp decrease in melting point because elements have simple molecular structures. Molecules are attracted to others by weak intermolecular forces of attraction.

26
Q

What is the trens in melting points across periods 2 and 3 for groups 15-18?

A

Melting point remains relatively low because the elements have simple molecular structures where molecules are held together by weak intermolecular forces of attraction.
Period 2 - N2 O2 F2 Ne
Period 3 - P4 S8 Cl2 Ar

27
Q

Physical properties of group 2 elements

A
  • reasonably high melting and boiling points
  • light metals with low densities
  • form white colourless compounds
28
Q

Reaction between group 2 elements and oxygen (Ca example)

A
  • react vigorously with oxygen
  • redox reaction
  • product is an ionic oxide, general formula is MO
  • 2Ca(g) + O2(g) –> 2CaO(s)
  • calcium is oxidised
  • oxygen is reduced
29
Q

Reaction between group 2 elements and water (Ca example)

A
  • form hydroxides with general formula M(OH)2
  • hydrogen gas is also formed
  • reactivity with water increases going down the group
  • redox, metal is oxidised and hydrogen is reduced
  • Ca(s) + 2H2O(l) –> Ca(OH)2(aq) + H2(g)
30
Q

Reactions between group 2 elements and dilute acids (Ca example)

A
  • all group 2 elements except Be react to for a salt and hydrogen gas
  • reaction becomes more vigorous moving down the group
  • Ca(s) + 2HCl(aq) –> CaCl2(aq) + H2(g)
31
Q

Reactions between group 2 oxides and water

A
  • form metal hydroxides
  • MO(s) + H2O –> M(OH)2(aq)
  • metal hydroxides are soluble in water and form alkaline solutions with water because they release OH- ions
32
Q

Solubility of group 2 metal hydroxides

A

Solubility decreases down the group:

  • Beryllium oxide is insoluble in water
  • Mg(OH)2 is slightly soluble in water
  • Ba(OH)2 is much more soluble in water so has a higher [OH-]
33
Q

Uses of group 2 compounds

A
  1. Neutralising acidic soils - Calcium hydroxide used
  2. Indigestion remedies - magnesium hydroxide neutralises excess stomach acid Mg(OH)2 + 2HCl –> MgCl2 + 2 H2O
  3. Building and construction uses - calcium carbonate used in limestone and marble, however there readily react with acids
34
Q

What is the trend in boiling points of the halogens and why?

A

Boiling point increases moving down the group as the state changes from gas to liquid to solid. This is because each successive element has an extra shell of electrons so there is a higher level of London forces between the molecules

35
Q

What is the trend in reactivity of the halogens and why?

A

Reactivity and oxidising power decreases moving down the group. This is because:

  • atomic radius increases so the nuclear pull is further away from incoming electrons
  • electron shielding increases so more repulsion
36
Q

What colour do halogens (Cl2, Br2 and F2) form in water

A

Cl2 - pale green
Br2 - orange
F2 - brown

37
Q

What colour do halogens (Cl2, Br2 and F2) form in cyclohexane

A

Cl2 - pale green
Br2 - orange
F2 - violet

38
Q

Which ions can chlorine oxidise? Include equations

A

Br- and I-
Cl2(aq) + 2Br-(aq) –> 2Cl-(aq) + Br2(aq)
Cl2(aq) + 2I-(aq) –> 2Cl-(aq) + I2(aq)

39
Q

Which ions can bromine oxidise?

A

Only I-

Br2(aq) + 2I-(aq) –> 2Br-(aq) + I2(aq)

40
Q

What is a disproportionation reaction?

A

A redox reaction where the same element is oxidised and reduced

41
Q

Water purification reaction

A

Cl2(aq) + H2O(l) –> HClO(aq) + HCl(aq)

Chlorine is oxidised and reduced

42
Q

Bleach formation reaction

A

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

Chlorine is oxidised and reduced

43
Q

How to identify carbonate ions

A
  • Add dilute acid to suspected carbonate
  • collect gas formed through limewater

A positive test will produce effervescence and turn limewater cloudy
CO3^2-(aq) + 2H+(aq) –> H2O(l) + CO2(g)

44
Q

How to identify sulfate ions

A
  • Add dilute HCl and barium chloride
    Positive test will produce a white precipitate
    Ba^2+(aq) + SO4^2-(aq) –> BaSO4(aq)
45
Q

How to identify halide ions

A

Add silver nitrate, add dilute ammonia and then concentrated ammonia.

  • Silver chloride has a white precipitate and is soluble in dilute NH3
  • Silver bromide has a cream precipitate and is soluble in concentrated NH3
  • silver iodide has a yellow precipitate and is insoluble in dilute and concentrated NH3
46
Q

How to identify ammonium ions

A
  • Add NaOH and warm gently
  • Ammonia gas will turn red litmus paper blue and has a distinctive smell
    NH4+(aq) + OH-(aq) –> NH3(g) + H2O(l)

Chemistry (6 decks)

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