The Periodic Table Flashcards

1
Q

What relationships are there between elements in the same:

  • group?
  • period?
A
  • Groups: similar chemical properties
  • Periods: trends in physical and chemical properties
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2
Q

Define first ionisation energy.

A

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

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3
Q

What factors affect ionisation energy and why?

A
  1. Nuclear charge: the more protons there are in the nucleus, the more positively charged the nucleus is and the stronger the attraction for the electrons.
  2. Atomic radius: The closer the electron is to the nucleus, the stronger the nuclear attraction.
  3. Shielding: As the no. shells of electrons between the nucleus and valence electrons increases, the nuclear attraction on the valence electrons decreases.
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4
Q

What is the trend in 1st ionisation energy down a group and why?

A
  • More shells of electrons so increased shielding effects.
  • Atomic radius is larger so nuclear attraction is lower.
  • These effects outweight the increase in nuclear charge, meaning that nuclear attraction on the valence electrons decreases overall.
  • Less energy required to remove them therefore the 1st ionisation energies become less endothermic.
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5
Q

What is the general trend in first ionisation energies across a period?

A
  • Proton number increases, increasing nuclear attraction.
  • Atomic radius becomes smalller.
  • Electrons are added to same shell so shielding doesn’t increase.
  • Nuclear attraction therefore increases overall therefore the 1st ionisation energies become from endothermic across the period.
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6
Q

When going across a period, why is there a drop in first ionisation energies:

  1. from group 2 to 3?
  2. from group 5 to 6?
A
  1. Outer electron in Grp. 3 elements is in a p-orbital rather than an s-orbital, which has a slightly higher energy so the electron is slightly further from the nucleus. There is also additional shielding from the s electrons. These factors outweight the increased nuclear charge so the ionisation energy drops slightly.
  2. Outer electron in Grp. 6 is removed from a full orbital instead of a singly-occupied orbital, therefore it experiences spin-pair repulsion so the ionisation energy drops slightly.
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7
Q

Explain the trend in successive ionisation energies.

A

The energy per ionisation increases because:

  • e- is removed from a shell closer to the nucleus/proton:electron ration is greater
  • nuclear attraction is greater
  • More E required to remove e-
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8
Q

How can you tell what group an element is in from its successive ionisation energies?

A

The number of IEs before a big jump (which signals the new electron is being removed from a shell closer to the nucleus) gives the number of valence electrons and therefore the group that the element is in.

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9
Q

What is a giant covalent lattice?

A

A network of atoms all bonded by strong covalent bonds.

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10
Q

What is are allotropes?

A

Different forms of the same element in the same state.

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11
Q

What is metallic bonding?

A

A strong electrostatic attraction between the metal cations and delocalised electrons.

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12
Q

What structure do solid metals form?

A

A giant lattice of closely packed cations surrounded by a sea of delocalised electrons.

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13
Q

What factors affect the:

  1. melting/boiling point
  2. malleability
  3. solubility
  4. thermal conductivity
  5. electrical conductivity

of metals?

A
  1. The no. of delocalised electrons per cation and the ionic radius.
  2. No bonds between cations therefore they are free to slide over each other, making metals malleable and ductile.
  3. No solid metals are soluble as the metallic bonds are too strong.
  4. The delocalised electrons can pass kinetic energy to each other, so metals are good thermal conductors.
  5. The delocalised electrons are free to move and carry charge, so metals are good electrical conductors.
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14
Q

What factors affect the:

  1. melting/boiling points
  2. thermal conductivity
  3. electrical conductivity
  4. solubility

of giant covalent substances?

A
  1. high melting point as lots of strong covalent bonds which all need to be broken
  2. vibrations travel easily through the stiff lattice so they are good thermal conductors
  3. all the electrons are held in localised bonds so they are not free to move and carry charge, so they are bad electrical conductors
  4. they are insoluble due to the strength of the covalent bonds
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15
Q

Explain the variation in melting and boiling points across a period.

A
  • For the metals: the ionic radius decreases and ionic charge increases across the period so the metallic bonds get stronger so the m.p. and b.p. increase.
  • For the giant covalent structures: high m.p. and b.p.
  • For the simple molecular elements: m.p. and b.p. much lower as only weak IM forces need to be overcome. The more atoms per molecule that an element exists as (e.g. S8 vs. F2), the more dipersion forces there are and therefore the higher the m.p. and b.p.
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16
Q

Why do Group 2 metals become more reactive as you go down the group?

A
  • Atomic radius increases
  • More shielding
  • These effects outweight the increase in nuclear charge
  • Nuclear attraction decreases
  • Ionisation energy decreases (v. important to make a statement about energy e.g. less energy required to remove electrons)
17
Q

What do group 2 metals produce when they react with:

  • water?
  • oxygen?
  • dilute acid?
A
  • M(OH)2 + H2 (g)
  • MO(s)
  • salt + H2 (g)
18
Q

What is the rate of reaction of each of the Group 2 metals with water/oxygen/dilute acid?

A
  • Be - doesn’t react
  • Mg - v. slowly
  • Ca - steadily
  • Sr - quickly
  • Ba - v. quickly
19
Q

What is formed when group 2 oxides react with water and what trend do these products show? What is the exception to this pattern?

A
  • Metal hydroxides which dissolve to form alkaline solutions. The oxides form more strongly alkanline solutions as you go down the group, as the hydroxides get more soluble.
  • MgO reacts v. slowly with water and Mg(OH)2 isn’t very soluble.
20
Q

Give two uses for group 2 compounds.

A
  • Ca(OH)2 is used in agriculture to neutralise acidic soils.
  • Mg(OH)2 and CaCO3 are antacids, used to neutralise stomach acid.
21
Q

Explain the trend in reactivity of the halogens.

A
  • Greater atomic radius down group.
  • More shells of electrons.
  • Increased shielding effects.
  • Nuclear attraction on incoming electron lower therefore more difficult to gain an electron.
  • Reactivity decreases down the group.
22
Q

Give the colour and state of fluorine, chlorine, bromine and iodine in their standard states.

A

Fluorine: pale yellow gas

Chlorine: pale green gas

Bromine: red-brown liquid

Iodine: grey solid

23
Q

Explain the trend in b.p. and m.p. down the halogens.

A
  • Volatility decreases down the group
  • More electrons per molecule as you go down the group
  • More ID-ID interactions
  • More energy required to overcome IM attractions
  • Higher m.p. & b.p.
24
Q

How do the reactions of halogens with metal halides show the trend in reactivity of the halogens?

A

Each halogen (not halide) gives the solution a particular colour, so if the solution changes colour you can see that a halogen has been displaced by a more reactive one.

25
Q

What colour are chlorine (Cl2), bromine (Br2) & iodine (I2) in:

  • aqueous solution?
  • organic solution?
A
  • Cl2: colourless, v. pale green
  • Br2: orange, orange
  • I2: brown, purple
26
Q

What is a disproportionation reaction?

A

A reaction in which the same element is both oxidised and reduced.

27
Q

How to test for the presence of specific halide ions?

A
  • Cl-: white ppt., dissolves in dil. NH3(aq)
  • Br-: cream ppt., dissolves in conc. NH3(aq)
  • I-: yellow ppt., insoluble in conc. NH3(aq)
28
Q

Write the equation for the formation of bleach.

A

2NaOH(aq) + Cl2(g) → NaClO(aq) + NaCl(aq) + H2O(l)

29
Q

Write the equation for the sterilisation of water.

A

Cl2(g) + H2O(l) ⇌ HCl(aq) + HClO(aq)

30
Q

What are the benefits and disadvantages of using chlorine to sterilise water?

A
  • Advantages:
    • Kills bacteria
    • Some chlorine remains in water and prevents reinfection further down supply.
    • Prevents growth of algae, eliminating bad tastes/smells & removes discolouration caused by organic compounds.
  • Disadvantages:
    • Toxic - irritates the respiritory system if inhaled
    • Can react with organic compounds from decomposing plants to form chlorinated hydrocarbons which are carcinogenic.
31
Q

What is the procedure to determine the identity of an anion?

Which false positive does this order avoid?

A
  • Add dilute nitric acid to test for carbonates - bubbles of CO2 will be produced if positive.
  • Add barium nitrate solution to test for sulphate ions - dense, white barium sulphate ppt. will form.
  • Add silver nitrate solution - white, cream or yellow ppt. will form depending on which halide is present.
  • Avoids:
    • Barium carbonate forming if barium nitrate to a solution before nitirc acid is used to remove all carbonate ions (BaCO3 is alos a white ppt. so couldn’t distinguish btwn. carbonate and sulphate ions)
    • Silver sulphate forming if silver nitrate is added before barium nitrate is used to remove all sulphate ions (Ag2SO4 is also a white ppt.)
32
Q

What is the procedure to determine the identity of a cation?

A
  • Add sodium hydroxide solution:
    • green ppt. - Fe(OH)2
    • orange-brown ppt. - Fe(OH)3
    • blue ppt. - Cu(OH)2
    • grey-green ppt. - Cr(OH)3
    • pale brown ppt. - Mn(OH)2
  • If no ppt., warm solution and waft damp litmus paper above mouth of test tube:
    • turns litmus paper blue (Warming causes the ammonium to combine with the hydroxide ions from NaOH added earlier to form ammonia gas and water. The ammonia gas redissolves on the litmus paper, reforming ammonium ions and producing hydroxide ions which turn the litmus paper blue.)