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Chemistry > C456 > Flashcards

C456 Flashcards

0
Q

In an atom, the quantities of what two particles are equal?

A

Protons and electrons.

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1
Q

Describe the three main subatomic particles.

A

Protons are heavy and positively charged.
Neutrons are heavy and neutral.
Electrons are tiny and negatively charged.

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2
Q

What do the different state symbols represent?

A

(S) - solid
(L) - liquid
(G) - gas
(Aq) - aqueous

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3
Q

Some elements produce distinctively coloured flames when burnt. What are these colours for lithium, sodium and potassium?

A

Red, yellow and lilac.

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4
Q

Why do different elements have different line spectra?

A

When heated, electrons release energy as light. Different elements emit different wavelengths as each element has a different electron arrangement. This can be used to identify elements (this is called spectroscopy).

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5
Q

Identify elements which have been identified by their line spectra.

A

Caesium, rubidium and helium.

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6
Q

How did Dobereiner organise his elements?

A

Triads - groups of three chemically similar elements. The middle element in the triad (organised by RAM) was the average of the other two.

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7
Q

How did Newlands organise his elements?

A

He listed the elements in rows of seven. The columns had similar properties but this broke down on the third row. It was denounced because he grouped dissimilar elements, he grouped metals and nonmetals and he didn’t leave any gaps.

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8
Q

How did Mendeleev organise his elements?

A

He put them in a table but he had to leave gaps. These gaps predicted undiscovered elements - when they were discovered, this was convincing elements for it.

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9
Q

Where are nonmetals found on the periodic table?

A

The right.

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10
Q

What are columns called in the periodic table?

A

Groups. Roman numerals are used for them.

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11
Q

What does each new row, or period, represent?

A

Another full electron shell.

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12
Q

How many electrons are there in the first three energy levels?

A

2, 8, 8. (Allegedly)

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13
Q

How are ions made?

A

When an atom loses or gains an electron.

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14
Q

How do ionic bonds occur?

A

When oppositely charged ions are attracted to each other. Ionic compounds form regular lattices. When these dissolve or melt, they conduct electricity as the ions can move.

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15
Q

In compounds, the total charge adds up to 0. What is the formula for iron (III) sulphate?

A

Fe2(SO4)3

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16
Q

What is group 1 known as? What changes as you go down the group?

A

Alkali metals.

They become more reactive, the density increases, the melting and boiling points decrease.

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17
Q

Reaction of potassium with water produced what? Give the reaction.

A

Potassium and hydroxide and hydrogen gas.

2K + 2H2O -> 2KOH + H2

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18
Q

What is group 7 known as? What changes as you go down the group?

A

The halogens.

They become less reactive, and the boiling and melting points increase.

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19
Q

What colours are the halogens?

A

F - Yellow
Cl - Green
Br - Orange
I - Purple

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20
Q

Describe the reactions of halogens with alkali metals, iron and displacement reactions.

A

With alkali metals, metal halides are formed, such as sodium chloride.
With iron, iron halides are formed.
Displacement reactions occur when a more reactive element displaces a less reactive one.

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21
Q

Give examples of oxidising, harmful, flammable, explosive, toxic, and corrosive materials.

A

Liquid oxygen, copper sulphate, petrol, some peroxides, hydrogen cyanide, concentrated sulphuric acid.

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22
Q

What safety precautions should be taken when using alkali metals or halogens.

A

Alkali metals can combust, and should be stored in oil to stop reactions with water vapour. They should never be touched.
Halogens are harmful and toxic. Liquid halogens should not touch the skin. They have poisonous vapours.

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23
Q

What gases, in what proportions, are there in the atmosphere?

A

78% nitrogen, 21% oxygen, 1% argon, 0.04% carbon dioxide.

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24
Q

Contrast the intramolecular forces and the intermolecular forces in the atmosphere.

A

The atoms within the molecules are held together by strong covalent bonds. However, the forces between the molecules are very weak. Only a little energy is needed to separate the molecules, so there are low melting and boiling points. These don’t conduct electricity.

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25
Q

Describe covalent bonding.

A

They involve sharing electrons. This is to ensure that the atoms have a full outer shell. They bond due to the electrostatic attraction between the electrons and the positive nuclei.

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26
Q

The earth’s hydrosphere contains many dissolved compounds. Give some examples.

A

Sodium chloride, magnesium chloride, potassium bromide…. The list goes on.

27
Q

Describe ionic compounds.

A

They are made of ions - oppositely charged ions attract each other. These build up into massive lattices. They have high melting and boiling points and do not normally conduct electricity, as there are no free electrons. However, when they are dissolved the ions are free to move and so can conduct electricity.

28
Q

Describe some common flame tests.

A 
When burnt:
Sodium gives an orange flame.
Potassium gives a lilac flame.
Calcium gives a red flame.
Copper gives a blue-green flame.
29
Q

Describe the coloured precipitates we can look for if we try to identify positive ions.

A

Add sodium hydroxide.
Calcium produces a white hydroxide precipitate.
Copper produces a blue precipitate.
Iron(II) produces a sludgy green precipitate.
Iron(III) produces a brownish precipitate.
Zinc produces a white precipitate which then dissolves to form a colourless solution.

30
Q

What is special about ionic equations?

A

They only show the useful parts of a reaction - e.g. The formation of a single product.

31
Q

How can we identify carbonates?

A

With dilute hydrochloric acid, carbonates fizz as they give off carbon dioxide. We can test for CO2 by bubbling it through limewater. It will go cloudy in the presence of CO2.

32
Q

How can we test for sulphates?

A

Add dilute HCl followed by barium chloride solution.
A white precipitate of barium sulphate means the original compound was a a sulphate. The HCl removes traces of carbonates.

33
Q

How can we test for halides?

A

Add dilute nitric acid (to remove carbonate ions) followed by silver nitrate solution.
A chloride gives a white precipitate of silver chloride.
A bromide gives a cream precipitate.
An iodide gives a yellow precipitate.

34
Q

The minerals in the lithosphere often contain which three elements?

A

Silicon, oxygen and aluminium.

35
Q

Carbon forms two giant covalent structures. Compare and contrast them.

A

Carbon atoms in diamond form four covalent bonds, making a very rigid structure. It is the hardest natural substance. It has a high melting point and is insoluble in water. It cannot conduct electricity.
Carbon atoms in graphite form three covalent bonds, creating sheet of carbon atoms that can slide easily. They can be rubbed of (viz. pencils). It has a high melting point. The free electrons conduct electricity.

36
Q

Name another giant covalent structure.

A

Silicon dioxide - similar properties to diamond.

37
Q

What are ores? Give two examples.

A

Ores are rocks that contain minerals from which metals can be extracted. These include haematite (a source of iron) and chalcopyrite (a source of copper).

38
Q

How can reactive metals be extracted? Give the symbol equation.

A

Via reduction with carbon or carbon monoxide.

2Fe2O3 + 3C -> 4Fe + 3CO2

39
Q

When the metal is more reactive than carbon, how do we extract the metal? Why does it work?

A

Electrolysis - the decomposition of a substance with electricity. A liquid - the electrolyte - is needed, normally free ions dissolved in water or a molten ionic compound. The free ions conduct electricity. In electrolysis electrons are taken from ions at the positive anode and given to ions at the negative cathode. They then become atoms or molecules.

40
Q

Electrolysis removes aluminium from its ore bauxite. Explain how.

A

The aluminium oxide is melted. The positive aluminium ions are attracted to the cathode where they pick up three electrons. They then sink to the bottom. The negative oxide ions are attracted to the anode where they lose two electrons. The atoms combine to form O2 molecules.
2Al2O3 -> 4Al + 3O2

41
Q

How do we find the relative atomic mass of an element?

A

We look at the top, bigger number next to an element in the periodic table.

42
Q

Describe metallic bonds.

A

Metallic bonds involve free electrons which come from the outer shell of each atom in the structure. The metal ions are held together by free electrons which can move.

43
Q

Explain how metallic bonds give metals their properties.

A

They conduct heat and electricity because of the free electrons which carry them.
They are strong and malleable - the layers can slide over each other. The bonds are strong.
They have high melting and boiling points due to the strong metallic bonds.

44
Q

What are the advantages and disadvantages of mining ores?

A

Advantages - useful products, local people get jobs and services can be improved.
Disadvantages - ores are a finite resource, it uses energy, and it scars the landscape.

45
Q

Recycling metals is important. Why? Give some stats?

A

Recycling metals can save 95% of the energy, a good deal of the ore, and a massive amount of waste.

46
Q

What is chemical synthesis?

A

Making complex chemical compounds from simpler ones.

47
Q

What products does the chemical industry make?

A

Food additives, cleaning products, drugs and fertilisers.

48
Q

Give some stats about the chemical industry.

A

200000 people in the UK are employed in the industry. Pharmaceuticals, toiletries and cleaning products are among the most important products.

49
Q

What are some common acids?

A

Solids, like citric acid and tartaric acid;
Liquids, like sulphuric, nitric and ethanoic acid;
Gases, like hydrogen chloride.

50
Q

Name some common alkalis.

A 
Sodium hydroxide (in cleaning products);
Potassium hydroxide (in batteries);
Calcium hydroxide (soil neutraliser).
51
Q

How can pH be determined?

A

Indicators like litmus paper, or universal indicator.

pH meters are more accurate however.

52
Q

Neutralisation reactions make what?

A

A salt and water.

An H+ ion from the acid reacts with the OH- from the alkali to form H2O.

53
Q

Acids react with metals to form…

A

A salt and hydrogen. The more reactive the metal, the more explosive the reaction. Copper doesn’t react with dilute acids. This can be ascertained by the squeaky pop test.

54
Q

An acid and a metal oxide or hydroxide form what?

A

Salt and water.

55
Q

Acid + Metal carbonate makes…

A

… Salt, water and carbon dioxide.

56
Q

What are the seven stages of chemical synthesis?

A 
Choosing the reaction,
Risk assessment,
Calculate the quantities of reactants,
Choose the apparatus and conditions,
Isolate the product,
Purification,
Measuring the yield and purity. 
(I could make a tune perhaps)
57
Q

Discuss three purification techniques.

A

Filtration - separate an insoluble solid from a liquid.
Evaporation / Crystallisation - heat the solution! leaving behind solid crystals. Repeat it a few times.
Drying - either dry it in a drying oven, or use desiccators.

58
Q

What is the formula for percentage yield?

A

Actual yield * 100 divided by theoretical yield.

59
Q

Describe a titration.

A

Add a known volume of alkali with indicator to a titration flask. Fill a burette with acid. Add a little acid to the flask, swirling regularly. Go slow when the endpoint is to be reached. When the indicator changes colour, note the volume of acid used.
(NB. Solids need to be put into a solution.)

60
Q

Some product need to be very pure. Give two examples.

A

Pharmaceuticals or petrochemicals.

61
Q

The purity of a compound can be worked out using a titration. How?

A

The concentration of the desired solution = n * (conc. of neutraliser * vol. of neutraliser)/ vol. of solution
Mass = concentration * volume
% purity = calculated mass/ original mass

62
Q

Why is energy management important in the chemical industry.

A

If a chemical synthesis reaction is exothermic, energy must be removed. Vice versa for endothermic.

63
Q

Contrast the rates of different reactions.

A

One of the slowest is iron rusting.
A medium one is metal reacting with acid.
A fast one is an explosion.

64
Q

Why does the chemical industry need to control rates of reaction?

A

Safety (avoiding explosions)

And economic reasons (not paying for lots of energy use)

65
Q

Rate of reaction depends on what four factors?

A

Temperature (due to more kinetic energy)
Concentration (due to more likely collisions)
Surface area (more chance of a collision)
Catalysts (provision of another chemical pathway)

66
Q

How can we measure the rates of reaction?

A

Precipitation, and the resulting colour change (quite subjective, though)
Change in mass (gas released into room, though)
Volume of gas given off (plunger can be blown out, though)

Chemistry (3 decks)

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