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chemistry > EL 7&8: Ionic bonding, precipitation reactions, ionisation enthalpy > Flashcards

EL 7&8: Ionic bonding, precipitation reactions, ionisation enthalpy Flashcards

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1
Q

What is a valence electron?

A

An outer shell electron that can participate in the formation of a chemical bond (if the shell is not full).

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2
Q
A

A

Elements so no ionic bonding involved, so skip s-block

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3
Q

What is ionic bonding?

A

Overall attraction in a lattice, consisting of attraction between ions of opposite charge and repulsion between ions of like charge.

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4
Q

Name 3 ways of making ionic salts.

A
  • Acid + alkali/base → salt + water
  • Acid + carbonate → salt + water + carbon dioxide
  • Acid + metal → salt + hydrogen
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5
Q

Draw a dot-and-cross diagram of sodium hydroxide.

A
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6
Q

Draw a dot-and-cross diagram of magnesium fluoride.

A
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7
Q

Give an example of a non-polar solvent.

A

Hexane.

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8
Q

Many ionic substances dissolve readily in water. List the ones which are insoluble.

A
  • Silver/lead/barium/calcium sulphates
  • Silver/lead halides
  • Metal carbonates
  • Metal hydroxides, except group 1 metals
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9
Q

Name 3 types of ionic substance which are always soluble in water.

A
  • Group 1 compounds
  • Ammonium compounds
  • Nitrates
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10
Q

What is a precipitate?

A

A suspension of solid particles produced by a chemical reaction in solution.

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11
Q

How would you test for Cu2+ ions? Name the change observed and the product responsible.

A

Precipitation reaction

  • Add NaOH solution to Cu2+ solution
  • Blue copper(II) hydroxide, Cu(OH)2, precipitate forms

Blue-green flame test

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12
Q

How would you test for Fe2+ ions? Name the change observed and the product responsible.

A

Precipitation reaction

  • Add NaOH solution to Fe2+ solution
  • Dirty green iron(II) hydroxide, Fe(OH)2, precipitate forms

​(No flame test)

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13
Q

How would you test for Fe3+ ions? Name the change observed and the product responsible.

A

Precipitation reaction

  • Add NaOH solution to Fe3+ solution
  • Orange-brown iron(III) hydroxide, Fe(OH)3, precipitate formed
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14
Q

How would you test for Pb2+ ions? Name the change observed and the product responsible.

A

Precipitation reaction

  • Add KI solution to Pb2+ solution
  • Yellow lead iodide, PbI2, precipitate forms

(No flame test)

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15
Q

How would you test for Cl- ions? Name the change observed and the product responsible.

A

Precipitation reaction

  • Add AgNO3 solution to Cl- solution
  • White silver chloride, AgCl, precipitate formed
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16
Q

How would you test for Br- ions? Name the change observed and the product responsible.

A

Precipitation reaction

  • Add AgNO3 solution to Br- solution
  • Cream silver bromide, AgBr, precipitate forms
17
Q

How would you test for I- ions? Name the change observed and the product responsible.

A

Precipitation reaction

  • Add AgNO3 solution to I- solution
  • Yellow silver iodide, AgI, precipitate forms
18
Q

What are the colours of the respective precipitates formed when solutions of chloride, bromide and iodide are reacted with silver nitrate solution?

A
  • Chloride = white
  • Bromide = cream
  • Iodide = yellow

More saturated colour down group

19
Q

How would you test for SO42- ions? Name the change observed and the product responsible.

A

Precipitation reaction

  • Add BaCl2 solution to SO42- solution
  • White barium sulfate, BaSO4, precipitate forms
20
Q
A

C

A: Ba(OH)2 is insoluble so (s) not (aq). Barium too electropositive for its chloride to be hydrolysed

B: Ba(OH)2 is formed

D: BaSO4 is insoluble so (s) not (aq)

21
Q
A

C

Silver chloride, silver sulfate and silver carbonate are all insoluble

but c uses silver carbonate - what’s special about this one?

22
Q
A

B

23
Q

Some students place a small piece of calcium in water with a few drops of universal indicator solution. The indicator turns an alkaline colour. After a time, the solution goes cloudy.
The equation for the reaction is Ca + 2H2O → Ca(OH)2 + H2

  1. Use the equation to suggest two observations that the students made, other than the indicator turning an alkaline colour.
  2. A student says that the cloudiness is caused by insoluble calcium hydroxide being formed slowly. Comment on this statement.
  3. The students filter off the solid from the solution, dry it and find it weighs 4.05 g. They calculate that the percentage yield of their reaction is 55.0%. Calculate the mass of calcium they started with.
  4. State an assumption you made in order to calculate (3).
A
  1. Calcium dissolves / reacts away, + bubbling / fizzing of hydrogen gas
  2. Ca(OH)2 is sparingly soluble, so dissolves at first, then reaches its maximum solubility and forms a precipitate, causing cloudiness
  3. 100 x 4.05 / theoretical = 55 → theoretical = 100 x 4.05 / 55 = 7.364 g → mol Ca(OH)2 = mol Ca = 0.09937 mol → mass Ca = 0.09937 x 40.1 = 3.98 g
  4. No Ca(OH)2 was in solution
24
Q

Some students attempt to make pure copper sulphate crystals from copper carbonate and sulphuric acid. They plan to use the following methods:

  • A: react excess copper carbonate with sulphuric acid. Evaporate. Allow to crystallise.
  • B: React copper carbonate with excess sulphuric acid. Filter, wash + dry.
  • C: react excess copper carbonate with sulphuric acid. Filter. Evaporate filtrate until dry.

None of these students would end up with pure crystals of copper sulphate. State what each one would end up with.

A
  • A: anhydrous CuSO4 and CuCO3
    • CuCO3 was in excess so some left over. Anhydrous since both are insoluble
  • ​B: nothing would collect on filter paper
    • ​Wouldn’t get anhydrous CuSO4 bc hasn’t been evaporated yet
  • C: anhydrous CuSO4
    • Get completely dry solid rather than crystals bc they evaporate fully, rather than until half volume then crystallising
25
Q

Define “first ionisation enthalpy”.

A

The energy needed to remove one electron from each of one mole of isolated gaseous atoms of an element, forming one mole of gaseous 1+ ions.

26
Q

Write the equation representing the first ionisation enthalpy of hydrogen.

A

H(g) → H+(g) + e-

27
Q

Define “second ionisation enthalpy”.

A

The energy needed to remove one electron from each of one mole of isolated gaseous 1+ ions of an element, forming one mole of gaseous 2+ ions.

28
Q

State and explain the trend in ionisation enthalpies across periods.

A

Ionisation enthalpy increases across periods:

  • All outer electrons in a period are in same shell
  • Proton number increases so greater electrostatic attraction between outer electrons + nucleus
  • So more energy required to remove an electron
29
Q

State and explain the trend in ionisation enthalpies down groups.

A

Ionisation energies decrease down groups:

  • More electron shells
  • Distance between nucleus + outer shell increases
  • Weaker electrostatic attraction
  • Easier to remove an electron
30
Q

How does observed data of 1st ionisation enthalpies provide evidence for:

  • electron shells?
  • electron sub-shells?
A

Electron shells

  • Ionisation enthalpy decreases down groups
  • Supports idea that inner shells shield outer ones, causing less attraction to nucleus

Electron sub-shells

  • Ionisation enthalpy increases across periods, with variations
  • So energy needed to remove electrons depends on their location
  • Supports idea of sub-shells with varied electron energies
31
Q

Why is there a decrease in 1st ionisation enthalpy between nitrogen and oxygen, despite the general trend of an increase across periods?

A
  • N has electronic configuration 1s22s22p3 and O 1s22s22p4
  • Extra repulsion from paired electron subshell in oxygen
  • Less energy needed to remove one of these electrons, despite increased nuclear charge

chemistry (50 decks)

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