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Chemistry > Topic 4 Chemical Change > Flashcards

Topic 4 Chemical Change Flashcards

C4

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1
Q

What is oxidation/reduction?

A

When metals react with oxygen we can have an oxidisation or reduction reaction.
Oxidation - Oxygen is added to an element or a compound after the reaction
Reduction - Oxygen is removed from an element or a compound after the reaction

A common example is the reaction with red-brown copper metal to produce black copper oxide:

2Cu + O2 ⟶ 2CuO

In this reaction copper metal has been oxidised since oxygen has been added to it
Another example is the reaction of zinc oxide with carbon:
ZnO + C ⟶ Zn + CO

In this reaction the zinc oxide has been reduced since it has lost oxygen. The carbon atom has been oxidised since it has gained oxygen.

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2
Q

What is the reactivity series of metals? What are the trends in reactivities of metals in reactions
with acids/water?

A
  • The series shows the metals in order of their reactivity.
  • Metal atoms form positive ions by loss of electrons when they react with other substances
  • The tendency of a metal to lose electrons is a measure of how reactive the metal is
  • A metal that is high up on the series loses electrons easily and is thus more reactive than one which is lower down on the series
  • Note that although carbon and hydrogen are non-metals, they are included in the series as they are useful in extracting metals from their oxides by reduction processes
    (See pg 3 Savemyexams for reactivity Series)
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3
Q

What is a displacement reaction?

A

A reaction where a more reactive metal displaces (PUSHES OUT) a less reactive metal from a compound

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4
Q

How are oxidation and reduction defined in terms of electron transfer?

A

Oxidation – loss of electrons
Reduction – gain of electrons

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5
Q

What is the general equation for a reaction between metals and acids? What type of reaction is this?

A

Metal + acid → salt + hydrogen
Redox reaction, also a displacement reaction

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6
Q

Which metals in the reactivity series will react with acid?

A

Those above hydrogen

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7
Q

What is the general equation for a neutralisation reaction?

A

Base + acid → salt + water

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8
Q

What is the general equation for the reaction between metal carbonate and acid?

A

Metal carbonate + acid → salt + water + carbon dioxide

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9
Q

What is the general equation for the reaction between metal oxides and acids?

A

Metal oxide + acid → a salt + water

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10
Q

What is a redox reaction?

A

A reaction where both oxidation and reduction occurs.
Exam Tip
REDOX are simultaneous reactions as they occur at the same time in the same reaction.

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11
Q

Explain in terms of gain or loss of electrons which species has been oxidised and which species has been
reduced when magnesium reacts with hydrochloric acid

A

Magnesium has lost electrons and thus has been oxidised (Mg to Mg2+)
The hydrogen in HCl has gained electrons and thus has been reduced (H+ to H2)

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12
Q

How is a soluble salt formed?

A

a) React the excess acid with some insoluble chemical (e.g. metal oxide)
b) Filter off the leftovers
c) Crystallise the product

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13
Q

What do acids and alkalis produce in aqueous solutions?

A

Acids produce hydrogen ions, alkalis produce hydroxide ions

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14
Q

What are bases, acids and alkalis?

A

Bases are chemicals that neutralise acids producing salt and water. Bases are usually metal oxides and metal hydroxides. eg are Copper Oxide (insoluble in water), Iron (III) Hydroxide (insoluble in water) and Sodium Hydroxide (soluble in water)
Alkalis are bases which are soluble in water. eg Sodium Hydroxide (soluble in water). In aqueous solutions alkalis produce hydroxide ions (OH-)

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15
Q

What is the pH scale and what does a pH of 7 show?

A

The pH scale is a numerical scale which is used to show how acidic or alkaline a solution is, as it is a measure of the amount of the hydrogen ions present in solution.
It goes from 0 to 14. From 0 - 6 acids; 7 neutral; above 7 alkali

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16
Q

State the general equation for a neutralisation reaction in a short, ionic form.

A

H+(aq) + OH− (aq)→ H2O(l)

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17
Q

What is a strong acid? What is a weak acid?

A

Strong acid is completely ionised (split) in aqueous solution; weak acid is only partially ionised in aqueous solution

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18
Q

What happens to pH as concentration of H+ increases?

A

As the pH scale decreased by 1 unit, the concentration of hydrogen ions increases by 10 times.

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19
Q

What is a concentrated acid and what is a diluted acid? Is this the same as a strong and weak acid?

A
  • Concentrated acid has more moles of acid per unit volume than dilute (dilute refers to solutions of low concentrations)
  • It is not the same - concentration is not the same thing as strength of an acid.
  • Strength refers to whether the acid is completely ionised in water (strong) or only partially (weak).
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20
Q

What does the concentration of acids tell us?

A

Amount of acid molecules in a given volume of solution. A diluted acid has fewer acid molecules in a given volume than a concentrated acid even if the strenght of the acid is the same

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21
Q

As the pH is decreased by one unit, what change is seen in the hydrogen ion concentration?

A

Increases by a factor of 10

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22
Q

Name the following salts: LiNO3, K2CO3, MgBr2,
BaSO4

A

Lithium nitrate
Potassium carbonate
Magnesium bromide
Barium sulfate

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23
Q

What is electrolysis?

A

The passing of an electric current through ionic substances that are molten or in solution to break them down into elements; ions are discharged (they lose/gain electrons) at electrodes to produce these

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24
Q

What is an electrolyte?

A

The liquid/solution which conducts electricity

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25
Q

What is a cathode and what is an anode?

A

Cathode is the negative electrode, anode is the positive electrode

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26
Q

What occurs at the cathode and what occurs at the anode during electrolysis?

A

Reduction occurs at the cathode
Oxidation occurs at the anode

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27
Q

In aqueous electrolysis, which element is discharged at the cathode? Oxygen is produced at the anode unless what?

A

The less reactive element discharges at the cathode. Hydrogen is produced unless there is a less reactive metal, in which case the said metal is produced. Oxygen is produced at the anode unless the solution contains halide ions, in which case halogen molecules are produced.

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28
Q

How is aluminium manufactured? Why is it expensive?

A

Aluminium is made through the electrolysis of aluminium oxide and cryolite.
Lots of energy is needed to produce the current in electrolysis which makes this process expensive.

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29
Q

What are the half equations in the extraction of aluminium?

A

Al3+ + 3 e− → Al (cathode)
2 O2− → O2 + 4 e− (anode)
Oxygen reacts with C of the anode producing CO2.

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30
Q

Why is cryolite used in this process?

A

It lowers the melting point of aluminium oxide, reducing energy costs

31
Q

What are the half equations in electrolysis of the
aqueous Na2SO4?

A

2 H+ + 2 e− → H2 (cathode)
4 OH− → 2 H2O + O2 + 4 e− (anode)

32
Q

What are the half equations in electrolysis of the
molten and aqueous KCl?

A

K+ + e− → K (cathode)
2 Cl− → Cl2 + 2 e− (anode)
2 H+ + 2 e− → H2 (cathode)
2 O2− → O2 + 4 e− (anode), respectively

33
Q

What are the half equations in electrolysis of the aqueous CuBr2?

A

Cu2+ + 2 e− → Cu (cathode)
2 Br− → Br2 + 2 e− (anode)

34
Q

Describe the reaction of Metals with Water

A
  • Some metals react with water
  • Metals above hydrogen in the reactivity series will react with water
  • For some metals such as iron, the reaction may be very slow
  • For other metals such as the alkali metals, the reaction may be quick and potentially hazardous because of their reactivity
  • Metals that react with cold water form a metal hydroxide and hydrogen gas:
    metal + water → metal hydroxide + hydrogen
    For example, calcium:
    Ca + 2H2 O → Ca(OH)2+ H2
    calcium + water → calcium hydroxide + hydrogen
  • Magnesium reacts very slowly with cold water when finely divided
  • Magnesium reacts with gaseous water to form a metal oxide and hydrogen gas:
    Mg + H2O → MgO + H 2
    magnesium + water → magnesium oxide + hydrogen
  • Titanium is another example of a metal that reacts with water to form the metal oxide and hydrogen
35
Q

Describe how metals react with Acids

A
  • Most metals react with dilute acids such as HCl
  • Only the ones below hydrogen in the reactivity series will not react with acids
    -When acids and metals react, the hydrogen atom in the acid is replaced by the metal atom to produce a salt and hydrogen gas:
    metal + acid → metal salt + hydrogen
    For example iron:
    Fe + 2HCI → FeCl2 + H2
    iron + hydrochloric acid → iron(II)chloride + hydrogen
  • In both these types of reactions (water and acids) the metals are becoming positive ions
  • The reactivity of the metals is related to their tendency to become an ion
  • The more reactive the metal the more easily it becomes an ion (by losing electrons)
    (see print out)

Exam Tip
Sometimes metals can fool us with their reactions. Aluminium is high in the reactivity series, but it does not react with water and the reaction with dilute acids can be quite slow. This is because it has a protective oxide layer that prevents reaction with these reagents. It reminds us that these reactions are trends or patterns rather than rules about chemical behaviour.

36
Q

Why are non metals in the reactivity series?

A

A reactivity series will usually contain the elements carbon and hydrogen.
This is because these elements play different roles in our understanding of the reactions of metals and our ability to predict how metals can be extracted from
their ores
From the reactions with water and acids we have seen that whether a reaction takes place depends on the position of the metal in the reactivity series relative to
hydrogen
A reaction takes place if the metal is able to displace hydrogen from water or acids
Carbon is a cheap reducing agent which can be used to remove oxygen from metal oxide ores
Placing carbon in the reactivity series allows us to see whether a metal oxide can be reduced or not by carbon
Metals below carbon can be extracted by heating the oxide with carbon
Metals higher than carbon have to be extracted by other methods, such as electrolysis

37
Q

What are displacement reactions?

A

Thee reactivity of metals decreases going down the reactivity series.
This means that a more reactive metal will displace a less reactive metal from its compounds
Two examples are:
Reacting a metal with a metal oxide (by heating)
Reacting a metal with an aqueous solution of a metal compound
For example, it is possible to reduce copper(II) oxide by heating it with zinc.
The reducing agent in the reaction is zinc:
Zn + CuO → ZnO + Cu
zinc + copper oxide → zinc oxide + copper
Displacement reactions occur when the solid metal is more reactive than the metal that is in the compound
See print out table

38
Q

How can e reactivity between two metals be compared

A

The reactivity between two metals can be compared using displacement reactions in salt solutions of one of the metals
This is easily seen as the more reactive metal slowly disappears from the solution, displacing the less reactive metal
For example, magnesium is a reactive metal and can displace copper from a
copper sulfate solution:
Mg + CuSO4 → MgSO4 + Cu
The blue colour of the CuSO solution fades as colourless magnesium sulfate solution is formed.
Copper coats the surface of the magnesium and also forms solid metal which falls to the bottom of the beaker

39
Q

How are metals extracted from their ores?

A

They have to be extracted from their ores through processes such as electrolysis, using a blast furnace or by reacting with more reactive material
In many cases the ore is an oxide of the metal, therefore the extraction of these
metals is a reduction process since oxygen is being removed
Common examples of oxide ores are iron and aluminium ores which are called haematite and bauxite respectively
Unreactive metals do not have to be extracted chemically as they are often found as the uncombined element
This occurs as they do not easily react with other substances due to their chemical stability
They are known as native metals and examples include gold and platinum which can both be mined directly from the Earth’s crust

Exam Tip
A metal can reduce another metal (remove oxygen) only if it is more reactive
than the metal that is bonded to the oxygen.

40
Q

Explain the link between Extraction of metals and the reactivity series

A

The most reactive metals are at the top of the series
The tendency to become oxidised is thus linked to how reactive a metal is and therefore its position on the reactivity series
Metals higher up are therefore less resistant to oxidation than the metals placed lower down which are more resistant to oxidation
The position of the metal on the reactivity series determines the method of extraction
Higher placed metals (above carbon) have to be extracted from their compounds using electrolysis as they are too reactive and cannot be reduced by carbon
Lower placed metals can be extracted from their compounds by heating with carbon which reduces them
E.g. The oxides of metals which are below carbon can be reduced by heating them with carbon
The carbon removes the oxygen from the metal oxide
Carbon dioxide is formed as well as the metal element:
metal oxide + carbon → metal + carbon dioxide
See print out

41
Q

How to Identify Oxidised & Reduced Species

Zinc displaces copper from a solution of copper(II)sulfate. Using ionic equations, determine which species undergoes oxidation and which species
undergoes reduction. See pg 14 Savemyexams

A

Using the principles of electron loss and gain it is possible to identify which species undergo oxidation and reduction in redox reactions

Answer
1. Write the full equation
2. Write the ionic equation
3. Use the ionic equation to rule out / ignore spectator ions that are present as reactants and products
4. Use the ionic equation to identify the species that is oxidised (OIL)
5. Use the ionic equation to identify the species that is reduced (RIG)

Exam Tip
Remember: OIL RIG - Oxidation Is Loss, Reduction Is Gain of electrons
After writing half equations, you can see if they are correct by checking that the number of electrons on either side is the same, which should combine to give 0 charge.

42
Q

How do metals react with acids?

A
  • Only metals above hydrogen in the reactivity series will react with dilute acids.
  • The more reactive the metal then the more vigorous the reaction will be.
  • Metals that are placed high on the reactivity series such as potassium and sodium are very dangerous and react explosively with acids.
  • When acids react with metals they form a salt and hydrogen gas:
    The general equation is:
    metal + acid ⟶ salt + hydrogen
    Exam Tip
    Sulfuric acid reacts with metals and produces sulfate salts while hydrochloric acid produces chloride salts.
    See pg 15 save my exams
43
Q

What type of reactions are Metal-acid reactions?

A

Metal-acid reactions are redox reactions (Redox means reduction and oxidation at the same time)
Exam Tip
Remember metal atoms tend to lose electrons and in these reactions are usually the species that undergoes oxidation.

44
Q

Are all reactions of acids neutralisations?

A

Not all reactions of acids are neutralisations
For example when a metal reacts with an acid, although a salt is produced there is no water formed so it does not fit the definition of neutralisation

45
Q

Give examples of the importance of neutralisation

A

Neutralisation is very important in the treatment of soils to raise the pH as some crops cannot tolerate pH levels below 7. This is achieved by adding bases to the soil such as limestone and quicklime

46
Q

What happens in a neutralisation reaction?

A

When an acid reacts with an alkali in a neutralisation reaction, the H+ ions react with the OH- ions to produce water.
This is the net ionic equation of all acid-base neutralisations and is what leads to a neutral solution, since water has a pH of 7:

H+ + OH-⟶ H 2O

47
Q

When does a neutralisation reaction occur?

A

When an acid reacts with a base (alkali), a neutralisation reaction occurs

48
Q

What Ph do bases have?

A

Above 7

49
Q

Bases and water

A

Many bases are insoluble in water but the ones that do dissolve in water are called alkalis
They thus form an alkaline solution
Examples of alkalis are soluble metal hydroxides such as NaOH and Ca(OH)

50
Q

What is produced in all acid-base neutralisation reactions?

A

In all acid-base neutralisation reactions, salt and water are produced:

acid + base ⟶ salt + water

If the base is a metal carbonate, carbon dioxide is also produced:
acid + base ⟶ salt + water + carbon dioxide

51
Q

What salt is formed in acid-base neutralisation reactions?

A

The identity of the salt produced depends on the acid used and the positive ions in the base
Hydrochloric acid produces chlorides, sulfuric acid produces sulfate salts and nitric acid produces nitrates

52
Q

What happens when metal oxides and metal hydroxides react with acid?

A

Metal oxides and metal hydroxides act as bases
When they react with acid, a neutralisation reaction occurs

53
Q

What happens when acids react with metal carbonates?

A

Acids will react with metal carbonates to form the corresponding metal salt, carbon dioxide and water

54
Q

Why there could be effervescence produced in an acid-base reaction?

A

Exam Tip
If in an acid-base reaction there is effervescence produced then the base must be a metal carbonate which produces carbon dioxide gas.

55
Q

What is salt?

A

A salt is a compound that is formed when the hydrogen ion in an acid is replaced by a metal ion
For example if we replace the hydrogen ion in HCl with a potassium ion, then the salt potassium chloride is formed, KCl
Salts are an important branch of chemistry due to the varied and important uses of this class of compounds
These uses include fertilisers, batteries, cleaning products, healthcare products and fungicides

56
Q

How are salts named?

A

The name of a salt has two parts
The first part comes from the metal, metal oxide or metal carbonate used in the reaction
The second part comes from the acid
The name of the salt can be determined by looking at the reactants
For example hydrochloric acid always produces salts that end in chloride and contain the chloride ion, Cl
Other examples:
Sodium hydroxide reacts with hydrochloric acid to produce sodium chloride
Zinc oxide reacts with sulfuric acid to produce zinc sulfate
See pg 20 Showmyexams for more examples

57
Q

Do salts have charge?

A

Salts have no overall charge since the sum of the charges on the ions is equal to zero

58
Q

How can a soluble salt be made?

A
  • A soluble salt can be made from the reaction of an acid with an insoluble base
  • During the preparation of soluble salts, the insoluble reactant is added in excess to ensure that all of the acid has reacted
  • If this step is not completed, any unreacted acid would become dangerously concentrated during evaporation and crystallisation
  • The excess reactant is then removed by filtration to ensure that only the salt and water remain
  • Since all of the acid has reacted and the excess solid base has been removed then the solution left can only be salt and water
  • If a carbonate was used as the solid base instead of an oxide or hydroxide, then any carbon dioxide gas produced would have been released into the atmosphere
  • A common example is the preparation of copper(II) sulfate which can be made with copper(II) oxide and dilute sulfuric acid:
    CuO (s) + H2SO4 (aq) ⟶ CuSO4(aq) + H2 2O (l)

The acid could also be reacted with a metal to produce the salt, as long as the metal is above hydrogen in the reactivity series and not too reactive so that a dangerous reaction does not take place

Exam Tip
Exam questions often ask why the solid oxide is added in excess. This is done to avoid leaving any unreacted acid which would become dangerously concentrated during evaporation and crystallisation.

59
Q

What happens when acids are added to water?

A

They form positively charged hydrogen ions (H+)

60
Q

What makes a solution acidic?

A

The presence of H+ ions is what makes a solution acidic

61
Q

What happens when alkalis are added to water?

A

They form negative hydroxide ions (OH-)

62
Q

What makes the aqueous solution an alkali

A

The presence of the OH- ions is what makes the aqueous solution an alkali

63
Q

How can Ph be measured?

A
  • pH can be measured using an indicator or a digital pH meter
64
Q

How does a Ph meter work?

A
  • pH meters contain a special electrode with a thin glass membrane that is dipped in the solution being tested. Hydrogen ions pass through it; the ions alter the voltage detected by the electrode
65
Q

What is a universal indicator to test Ph?

A

It is a special indicator made from a number of dyes. It turns a range of colours as the Ph changes. See picture pg 97 Chemistry book

66
Q

Name one acid you use in labs

A

Hydrocloric acid. This is formed when the gas hydrogen chloride is dissolved in water.
HCl(g) {water} H+(aq) - Cl-(aq)

67
Q

How do you get sodium hydroxide solution?

A

You dissolve solid sodium hydroxide in water.
NAOH(s) {water} Na+(ac) + OH-(aq)

68
Q

What causes acids to b weak or strong?

A

Acids must dissolve in water before they show their acidic properties. When added to water, acids ionise (split up) to produce H+ (aq) ions and negative ions. It is the H+ ions that all acidic solutions have in common. Acids can be either strong or weak, depending on how many ions they produce when they dissolve in water

69
Q

What are strong acids?

A

Strong acids such as HCl and H2SO4 dissociate completely in water, producing solutions with a high concentration of H+ ions and thus a very low pH

70
Q

What are weak acids?

A

Weak acids such as ethanoic acid, CH3COOH, and hydrofluoric acid, HF, only partially ionise in water, producing solutions of pH values between 4 – 6. The reaction is reversible (unlike the ionisation of a strong acid). So the molecules of weak acids split up to form H+ ions and negative ions, the ions recombine to form the original molecules again. At the end, a position of equilibrium is reached in which both whole molecules (the majority) and their ions (the minority) are present. So in ethanoic acid you get:
CH3COOH(aq) – CH3COO-(aq) + H+(aq)
ethanoic acid – ethanoate ion + hydrogen ions

71
Q

Examples of strong acids

A

Hydrocloric acid
nitric acid
sulfuric acid

72
Q

Example of weak acids

A

ethanoic acid (found in vinegar)
citric acid
carbonic acid (found in rain water and fizzy drinks)

73
Q

How are Ph values related to the concentration of H+ ions?

A

As the concentration of H+ (aq) ions increases by a factor of 10, the pH value decreases by 1 unit.
The pH scale is logarithmic, meaning that each change of 1 on the scale represents a change in concentration by a factor of 10.
Two acids of equal concentration, where one
is strong and the other is weak, then the strong acid will have a lower pH due to its capacity to dissociate more and hence put more H+ ions into solution than the weak acid

Exam Tip
Acid strength indicates the proportion of acid molecules that dissociate while concentration is a measure of how much acid there is per unit volume
of water.
See pg 99 Chemistry book

Chemistry (21 decks)

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