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(GCSE) Chemistry > C5. Chemical Changes (Y10 - Summer 1) > Flashcards

C5. Chemical Changes (Y10 - Summer 1) Flashcards

1
Q

🟢 What sides of the pH scale are alkaline and acidic

A

The pH scale goes from 0 - 14.

The numbers from 6 - 0 are more acidic, while numbers from 8 - 14 are more alkaline

The pH of 7 is neutral.

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2
Q

🟢 What are Bases?

A

Bases are insoluable substances that are neutralised by acids, and they form a salt + water, and sometimes carbon dioxide.

Examples of bases are:

  • Metal Oxides
  • Metal Hydroxides
  • Metal Carbonates

(ANYTHING THAT IS A CARBONATE IS DEFINITELY A BASE)

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3
Q

🟢 What are Alkalis?

A

Alkalis are substances on the upper end of the pH scale and release hydroxide ions in water OH- (aq).(Alkalis = Hydroxides that dissolve in water)

(ANYTHING THAT IS A HYDROXIDE IS DEFINITELY AN ALAKALI)

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4
Q

🟢 What are Acids?

A

Acids are substances on the lower side of the pH scale that are on that release hydrogen ions when added to water.

There are 2 types of acids, Strong Acids and Weak Acids

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5
Q

🟢 How are Strong and Weak Acids diffferent?

A

Strong acids release all hydrogen ions when added to water.

Weak acids only release some of the acid molecules release H+ ions in water.

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6
Q

🟢 Reactions of acids with bases:

  • Acid + Metal Oxide —>
  • Acid + Metal Hydroxide —>
  • Acid + Metal Carbonate —>
A

Acid + Metal Oxide —> Salt + Water

Acid + Metal Hydroxide —> Salt + Water

Acid + Metal Carbonate —> Salt + Water + Carbon Dioxide

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7
Q

🟢 What is type of salt is made out of:

  • Sulphuric Acid (H2SO4)
  • Nitric Acid (HNO3)
  • Hydrochloric Acid (HCl)
A 
Sulphuric Acid (H2SO4):
 Sulphate
Nitric Acid (HNO3): 
 Nitrate
Hydrochloric Acid (HCl):
 Chloride
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8
Q

🟢 Reactions of acids with alkalis:

  • Metal + Acid —>
  • Alkali (Metal Hydroxides) + Acid —>
  • Carbonate + Acid —>
  • Base (Metal Oxides) + Acid —>
A
  • Metal + Acid —> Salt + Hydrogen
  • Alkali + Acid —> Water + Salt
  • Carbonate + Acid —> Water + Carbon Dioxide + Salt
  • Base + Acid —> Water + Salt
  • Remember*:
  • MASH
  • AAWS
  • CAWCS
  • BAWS
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9
Q

🟢 Strong Concentrated Acid:

  • How much is dissolved
  • Proportion of molecules that react with water to form H+ (aq)
  • Example of Acid
  • Typical pH
  • Reaction with Magnesium
  • Colour with Universal Indicator
A

How much is dissolved:
Lots

Proportion of molecules that react with water to form H+ (aq):
All

Example of Acid:
Hydrochloric Acid

Typical pH:
0 - 1

Reaction with Magnesium:
Vigourous Fizzing

Colour with Universal Indicator:
Red

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10
Q

🟢 Strong Dilute Acid:

  • How much is dissolved
  • Proportion of molecules that react with water to form H+ (aq)
  • Example of Acid
  • Typical pH
  • Reaction with Magnesium
  • Colour with Universal Indicator
A

How much is dissolved:
Not Much

Proportion of molecules that react with water to form H+ (aq):
All

Example of Acid:
Hydrochloric Acid

Typical pH:
2 - 3

Reaction with Magnesium:
Fizzing

Colour with Universal Indicator:
Red

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11
Q

🟢 Weak Concentrated Acid:

  • How much is dissolved
  • Proportion of molecules that react with water to form H+ (aq)
  • Example of Acid
  • Typical pH
  • Reaction with Magnesium
  • Colour with Universal Indicator
A

How much is dissolved:
Lots

Proportion of molecules that react with water to form H+ (aq):
Few

Example of Acid:
Ethanoic Acid

Typical pH:
4 - 5

Reaction with Magnesium:
Gentle Fizzing

Colour with Universal Indicator:
Orange

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12
Q

🟢 Weak Dilute Acid:

  • How much is dissolved
  • Proportion of molecules that react with water to form H+ (aq)
  • Example of Acid
  • Typical pH
  • Reaction with Magnesium
  • Colour with Universal Indicator
A

How much is dissolved:
Not Much

Proportion of molecules that react with water to form H+ (aq):
Few

Example of Acid:
Ethanoic Acid

Typical pH:
5 - 6

Reaction with Magnesium:
Hardly Any Fizzing

Colour with Universal Indicator:
Yellow

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13
Q

🟢 What salt is fromed from:

Iron + Hydrochloric Acid –>
Hydrochloric Acid + Copper Carbonate–>
Iron (II) Hydroxide + Sulphuric Acid –>
Nitric Acid + Calcium Oxide–>

(Give word equations and also chemical equations)

A

Iron + Hydrochloric Acid –> Iron Chloride + Hydrogen
(Fe + 2HCl –> FeCl2 + H2)

Hydrochloric Acid + Copper Carbonate–> Water + Carbon Dioxide + Copper Chloride
(2HCl + CuCO3 –> CuCl2 + H2O + CO2)

Iron (II) Hydroxide + Sulphuric Acid –> Water + Iron (II) Sulphate
(Fe(OH)2 + H2SO4 –> FeSO4 + 2H2O)

Nitric Acid + Calcium Oxide–> Water + Calcium Nitrate
(2NHO3 + CaO –> Ca(NO3)2 + H2O

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14
Q

🟢 What reactants form these salts:

___ + ___ –> Zinc Sulfate + Water + Carbon Dioxide

___ + ___ –> Magnesium Nitrate + Hydrogen

___ + Potassium Oxide –> Potassium Chloride + ___

Calcium Hydroxide + ___ –> Calcium Citrate + ___

A

Sulfuric Acid + Zinc Carbonate –> Zinc Sulphate + Water + Carbon Dioxide

Magnesium + Nitric Acid –> Magnesium Nitrate + Hydrogen

Hydrochloric Acid + Potassium Oxide –> Potassium Chloride + Water

Calcium Hydroxide + Citric Acid –> Calcium Citrate + Water

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15
Q

🟢 What is the Difference between Strong and Weak Acids

A

Strong acids completely ionise in aqueous solutions, while weak acids partially ionise in aqueous solution

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16
Q

🟢 What is the Ionic Equation for any acid with any alkali

A

H+ (aq) + OH- (aq) –> H2O (l)

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17
Q

🟢 What are Examples and Formulas of 3 Strong Acids

A

Examples of Strong Acids are:

Hydrochloric Acid —> Hydrogen + Chloride ions
(HCl(aq) –> H^+(aq) + Cl^-(aq))

-Sulfuric Acid —> Hydrogen + Sulphate Ions
(H2SO4(aq) –> 2H^+(aq) + SO4 ^2- )

-Nitric Acid —> Hydrogen + Nitrate Ions
(HNO3(aq) —> H+(aq) + NO3-)

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18
Q

🟢 What are 3 Examples of Weak Acids

A

Examples of Weak Acids are:

-Ethanoic Acid ⇌ Hydrogen + Acetate
(CH3COOH (aq) ⇌ H+ + CH3COO-

  • Citric Acid
  • Carbonic Acid
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19
Q

🟢 What colour does Methyl Orange make:

  • Strong Acids (lots of H+)
  • Weak Acids (a few H+)
  • Water (H+ = OH-)
  • Weak Alkali (a few OH-)
  • Strong Alkali (lots of OH-)
A 
Strong Acids (lots of H+):
Red
Weak Acids (a few H+):
Pink

Water (H+ = OH-):
Orange

Weak Alkali (a few OH-):
Yellow
Strong Alkali (lots of OH-):
Yellow
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20
Q

🟢 What colour does Phenolphtalein Solution make:

  • Strong Acids (lots of H+)
  • Weak Acids (a few H+)
  • Water (H+ = OH-)
  • Weak Alkali (a few OH-)
  • Strong Alkali (lots of OH-)
A 
Strong Acids (lots of H+):
Colourless
Weak Acids (a few H+):
Colourless

Water (H+ = OH-):
Colourless

Weak Alkali (a few OH-):
Pink
Strong Alkali (lots of OH-):
Pink
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21
Q

🟢 What colour does Litmus Paper make:

  • Strong Acids (lots of H+)
  • Weak Acids (a few H+)
  • Water (H+ = OH-)
  • Weak Alkali (a few OH-)
  • Strong Alkali (lots of OH-)
A 
Strong Acids (lots of H+):
Red
Weak Acids (a few H+):
Red

Water (H+ = OH-):
Blue

Weak Alkali (a few OH-):
Blue
Strong Alkali (lots of OH-):
Blue
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22
Q

🟢 What colour does Universal Indicator Solution make:

  • Strong Acids (lots of H+)
  • Weak Acids (a few H+)
  • Water (H+ = OH-)
  • Weak Alkali (a few OH-)
  • Strong Alkali (lots of OH-)
A 
Strong Acids (lots of H+):
Red - pH of 1
Weak Acids (a few H+):
Orange - pH of 5

Water (H+ = OH-):
Green - pH of 7

Weak Alkali (a few OH-):
Green - pH of 10
Strong Alkali (lots of OH-):
Blue - pH of 14
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23
Q

🟢 What colour does Universal Indicator Paper make:

  • Strong Acids (lots of H+)
  • Weak Acids (a few H+)
  • Water (H+ = OH-)
  • Weak Alkali (a few OH-)
  • Strong Alkali (lots of OH-)
A 
Strong Acids (lots of H+):
Red - pH of 1
Weak Acids (a few H+):
Orange - pH of 5

Water (H+ = OH-):
Green - pH of 7

Weak Alkali (a few OH-):
Green - pH of 10
Strong Alkali (lots of OH-):
Blue - pH of 14
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24
Q

🟢 What is the pH scale, and How does it show Relative Acidity?

A

pH is a measure if the concentration if H+ ions in solution.

The pH scale is logarthmic, meaning that each change of 1 on the scale represented a change in concentration by a factor of 10.

-For 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.

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25
Q

🟢 What is a Redox Reaction

A

Redox reactions are reactions where both oxidation and reduction are taking place.

Displacement reactions are examples of redox reactions as one species is being oxidised (losing electrons) while the other is being reduced (gaining electrons).

For example, if magnesium was added to copper sulphate solution, the magnesium metal would be oxidised, while the copper ions were being reduced.

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26
Q

🟢 What is Oxidisation?

A

Oxidisation:

Oxidation is the loss of electrons by a reactant.

When a metal element is reacting to form a compound then it is being oxidised.

For example:
Mg(s) + O2(g) → MgO(s)

The metal atoms are losing electrons to form an ion. They are being oxidised.

Mg(s) → Mg2+(aq) + 2e−

This is known as an ion-electron equation.

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27
Q

🟢 What is Reduction?

A

Reduction:

Reduction is the opposite of oxidation. It is the gain of electrons.

Compounds reacting that result in metal elements being formed are examples of reduction reactions.

For example:
Cu2+(aq) + 2e− → Cu(s)

The metal ions are gaining electrons to form atoms of the element. They are being reduced.

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28
Q

🟢 What happens when metals react with acids (use Mnemonic)

A

O xidisation
I s
L oss of electrons

R eduction
I s
G ain of electrons

29
Q

🟢 What is a half equation (+ examples)

A

A half equation shows what happens when ions gain or lose electrons:

Mg (s) - 2e^- –> Mg^2+ (aq)

Fe (s) - 2e^- –> Fe^2+ (aq)

2n (s) -2e^- –> 2n^2+ (aq)

30
Q

🟢 What factors will affect how readily a metal atom will lose it’s outer shell electrons

A
  • The number of electrons
    i. e, easier to lose one lectron (group 1 elements) than 2 electrons (group 2 elements).
  • The distance between the outer shell electrons from the nucleus
    i. e, the bigger the atom, the easier it is to lose electrons.
  • Electron Shielding
    i. e, the greater the number of shielding electrons, the weaker the nuclear attraction on the outer shell electrons, the easier it is to lose.
31
Q

🟢 Explain the trend in reactivity down group 1 elements

A

The trend in reactivity in group 1 is that the further down you go in the group, the more reactive the metals get. When they react, they will always lose their one outer shell electron. This is because atoms further down the group will be larger, resulting in the outer shell electrons to experience a weaker nuclear force of attraction. Also, electron shielding does not help the cause, as it so sometimes blocks some of the force, making it even easier for the outer shell electrons to transfer to another atom.

32
Q

🟢 What does the reactivity series of a metal show (+ what is it)

A

Tne reactivity of a metal is related to its tendancy to form positive ions. Metals can be arranged in order of their reactivity from their reactions with water and dilute acids in a reactivity series.

33
Q

🟢 Order of Metals in the Reactivity Series (Most Reactive to Least Reactive)

A
  • Potassium, K
  • Sodium, Na
  • Lithium, Li
  • Calcium, Ca
  • Magnesium, Mg
  • Zinc, Zn
  • Iron, Fe
  • Copper, Cu
  • Gold, Au
34
Q

🟢 Potassium: (Does it Displace Hydrogen)

  • Reaction with O2
  • Reaction with H2O
  • Reaction with Acid
  • Half-Equation
A

(More Reactive Than Hydrogen)

Reaction with O2:
-Burn in oxygen to form a metal oxide

Reaction with H2O:
-Fizz vigorously as H2(g) is produced. Other product is the metal hydroxide

Rection with Acid:
-Too Dangerous

Half Equation:
- K(s) –> K+ (aq) + 1e-

35
Q

🟢 Sodium: (Does it Displace Hydrogen)

  • Reaction with O2
  • Reaction with H2O
  • Reaction with Acid
  • Half-Equation
A

(More Reactive Than Hydrogen)

Reaction with O2:
-Burn in oxygen to form a metal oxide

Reaction with H2O:
-Fizz vigorously as H2(g) is produced. Other product is the metal hydroxide

Rection with Acid:
-Too Dangerous

Half Equation:
- Na(s) –> Na+ (aq) + 1e-

36
Q

🟢 Lithium: (Does it Displace Hydrogen)

  • Reaction with O2
  • Reaction with H2O
  • Reaction with Acid
  • Half-Equation
A

(More Reactive Than Hydrogen)

Reaction with O2:
-Burn in oxygen to form a metal oxide

Reaction with H2O:
-Fizz vigorously as H2(g) is produced. Other product is the metal hydroxide

Rection with Acid:
-Too Dangerous

Half Equation:
- Li(s) –> Li+ (aq) + 1e-

37
Q

🟢 Calcium: (Does it Displace Hydrogen)

  • Reaction with O2
  • Reaction with H2O
  • Reaction with Acid
  • Half-Equation
A

(More Reactive Than Hydrogen)

Reaction with O2:
-Burn in oxygen to form a metal oxide

Reaction with H2O:
-Fizz vigorously as H2(g) is produced. Other product is the metal hydroxide

Rection with Acid:
-Fizz as H2(g) is produced. Other product is a salt

Half Equation:
- Ca(s) –> Ca^2+ (aq) + 2e-

38
Q

🟢 Magnesium: (Does it Displace Hydrogen)

  • Reaction with O2
  • Reaction with H2O
  • Reaction with Acid
  • Half-Equation
A

(More Reactive Than Hydrogen)

Reaction with O2:
-Burn in oxygen to form a metal oxide

Reaction with H2O:
-Fizz as H2(g) is produced. Other product is the metal hydroxide (takes longer)

Rection with Acid:
-Fizz as H2(g) is produced. Other product is a salt

Half Equation:
- Mg(s) –> Mg^2+ (aq) + 2e-

39
Q

🟢 Zinc: (Does it Displace Hydrogen)

  • Reaction with O2
  • Reaction with H2O
  • Reaction with Acid
  • Half-Equation
A

(More Reactive Than Hydrogen)

Reaction with O2:
-Burn in oxygen to form a metal oxide

Reaction with H2O:
-Fizz as H2(g) is produced. Other product is the metal hydroxide (takes longer)

Rection with Acid:
-Fizz as H2(g) is produced. Other product is a salt

Half Equation:
- Zn(s) –> Zn^2+ (aq) + 2e-

40
Q

🟢 Iron: (Does it Displace Hydrogen)

  • Reaction with O2
  • Reaction with H2O
  • Reaction with Acid
  • Half-Equation
A

(More Reactive Than Hydrogen)

Reaction with O2:
-Burn in oxygen to form a metal oxide

Reaction with H2O:
-Fizz as H2(g) is produced. Other product is the metal hydroxide (takes longer)

Rection with Acid:
-Fizz as H2(g) is produced. Other product is a salt

Half Equation:
- Fe(s) –> Fe^2+ (aq) + 2e-

41
Q

🟢 Copper: (Does it Displace Hydrogen)

  • Reaction with O2
  • Reaction with H2O
  • Reaction with Acid
  • Half-Equation
A

(Less Reactive Than Hydrogen)

Reaction with O2:
-Tarnishes

Reaction with H2O:
-No Reaction

Rection with Acid:
-No Reaction

Half Equation:
- Cu(s) –> Cu^2+ (aq) + 2e-

42
Q

🟢 Gold: (Does it Displace Hydrogen)

  • Reaction with O2
  • Reaction with H2O
  • Reaction with Acid
  • Half-Equation
A

(Less Reactive Than Hydrogen)

Reaction with O2:
-No Reaction

Reaction with H2O:
-No Reaction

Rection with Acid:
-No Reaction

43
Q

🟢 Half Equation of Hydrogen

A

H2 (g) –> 2H+ (aq) + 2e-

44
Q

🟢 Explain, in terms of electrons, what happens when metals react

A

Metal atoms when they react, lose electrons to form positive ions. They are oxidised.

45
Q

🟢 Explain, in terms of electrons, why magnesium is more reactive than copper

A

Magnesium loses electrons more easily than copper, as it is more reactive than copper. This means magnesium would displace copper in a reaction.

46
Q

🟢 Zinc reacts with a solution of copper (Ii) sulphate to form zinc(II) sulphate in a displacement reaction.

A

Zinc displaces copper because zinc is more reactive than a solution of copper, meaning zinc loses it’s electrons more easily.

47
Q

🟢 What is Reduction and Oxidisation?

A

Reduction and Oxidisation usually happen together. Reactions where both reduction and oxidisation take place are called Redox reactions. There are two definitions of reduction and oxidisation.

48
Q

🟢 Oxidisation and Reduction in terms of Oxygen and Electrons

A

Oxidisation in terms of Oxygen:
-Gain of Oxygen

Oxidisation in terms of Electrons:
-Loss of Electrons

Reduction in terms of Oxygen:
-Loss of Oxygen

Reduction in terms of Electrons:
-Gain of Electrons

49
Q

🟢 How to Write Half Equations and Ionic Equations

A

We can write two half equations for each Redox reaction (one for the Reduction process and one for the Oxidisation process) and combine them to give an overall Ionic eauation to show Redox.

This overall ionic equation leaves out any ions that do not change. It just shows what happens to anything that changes.

50
Q

🟢 How To Work out an Ionic Equation

A
  1. Work out the half equations to show the metal being ozidised and the metal ions being Reduced
  2. Ensure the number of electrons balance am combine the two half equations to give an ionic equation to show Redox.
51
Q

🟢 2Al + Fe2O3 –> Al2O3 + 2Fe

  • Oxidisation Half Equation
  • Reduction Half Equation
  • Ionic Equation
A

Oxidisation Half Equation:
Al - 3e- –> Al3+
(atom - electrons –> ions)

Reduction Half Equation:
Fe3+ + 3e- –> Fe
(ion + electrons –> atoms)

Ionic Equation:
Al + Fe3+ –> Al3+ + Fe
(atom + ion –> ion + atom)

52
Q

🟢 Mg + 2AgNO3 –> Mg(NO3)2 + 2Ag

  • Oxidisation Half Equation
  • Reduction Half Equation
  • Ionic Equation
A

Oxidisation Half Equation:
Mg - 2e- –> Mg2+
(atom - electrons –> ions)

Reduction Half Equation:
Ag+ + e- –> Ag
(ion + electrons –> atoms)

Ionic Equation:
Mg + 2Ag –> Mg 2+ + 2Ag
(atom + ion –> ion + atom)

53
Q

🟢 Cu +2AgNO3 –> Cu(NO3)2 + 2Ag

  • Oxidisation Half Equation
  • Reduction Half Equation
  • Ionic Equation
A

Oxidisation Half Equation:
Cu - 2e- –> Cu2+
(atom - electrons –> ions)

Reduction Half Equation:
Ag + e- –> Ag
(ion + electrons –> atoms)

Ionic Equation:
Cu + Ag+ –> Cu2+ + 2Ag
(atom + ion –> ion + atom)

54
Q

🟢 Mg + ZnO –> MgO+Zn

  • Oxidisation Half Equation
  • Reduction Half Equation
  • Ionic Equation
A

Oxidisation Half Equation:
Mg - 2e- –> Mg2+
(atom - electrons –> ions)

Reduction Half Equation:
Zn2+ + 2e- –> Zn
(ion + electrons –> atoms)

Ionic Equation:
Mg + Zn2+ –> Mg2+ + Zn
(atom + ion –> ion + atom)

55
Q

🟢 How does Magnesium Sulphate - MgSO4(aq) react with:

  • Magnesium
  • Zinc
  • Iron
  • Copper
A

Magnesium - No reaction

Zinc - No reaction

Iron - No reaction

Copper - No reaction

56
Q

🟢 How does Zinc Sulphate - ZnSO4(aq) react with:

  • Magnesium
  • Zinc
  • Iron
  • Copper
A

Magnesium - Slight fizzing + black powder

Zinc - No reaction

Iron - No reaction

Copper - No reaction

57
Q

🟢 How does Iron Sulphate - FeSO4(aq) react with:

  • Magnesium
  • Zinc
  • Iron
  • Copper
A

Magnesium - Slight fizzing

Zinc - Small bubbling

Iron - No reaction

Copper - No reaction

58
Q

🟢 How does Copper Sulphate - FeSO4(aq) react with:

  • Magnesium
  • Zinc
  • Iron
  • Copper
A

Magnesium - Forms Black Powder

Zinc - Turns Black

Iron - Slightly rusts

Copper - No reaction

59
Q

🟢 What is a Metal Ore?

A

If a metal can be extracted for profit from the compounds in a rock, then the rock is called on ore.

60
Q

🟢 How are Metals Removed From Ores

A

Most ores contain METAL OXIDES. To extract the metal from the metal oxide, the oxygen is removed. Reactions that remove oxygen are called reduction reactions.

e.g. Al2O3 → Al

However, when all metals are extracted, metal ions in the compounds gain electrons to form metal atoms. This means that all extraction reactions involve reduction.

e.g. NaCl→ Na (Na+ + e– →Na)

61
Q

🟢 Where Do Metals Come From?

A

Only a few metals are found as elements on Earth – these are the least reactive metals (e.g. gold, platinum)

Most metals are produced by chemical reactions (“extracted”) from compounds found in rocks (e.g. aluminium is produced from aluminium oxide found in bauxite).

62
Q

🟢 What Metals are already found as elements

A
  • Platinum

- Gold

63
Q

🟢 What Metals need Thermal Decomposition to sperate them from ores?

A
  • Silver

- Mercury

64
Q

🟢 What Metals need displacement with Carbon to sperate them from ores?

A
  • Copper
  • Lead
  • Tin
  • Nickel
  • Iron
  • Chromium
  • Zinc
65
Q

🟢 What Metals need electrolysis (electrical decomposition) to sperate them from ores?

A
  • Aluminium
  • Magnesium
  • Calcium
  • Sodium
  • Potassium
66
Q

🟢 What Metals are easier to extract

A

Etals that are less reactive and have weak bonds in compounds are easier to extract.

67
Q

🟢 Explain why is the extraction of calcium from calcium chloride by electrolysis involves a redox reaction (include half-equations)

A

This extract involves a redox recation because the the calcium ion will be being reduced, and will gain 2 electrons to make a calcium atom.

This is shown by the half equation: Ca2+ + 2e– → Ca

68
Q

🟢 One method of extracting zinc involves the reaction of zinc oxide with carbon. Explain, both in terms of
oxygen and electrons, why this extraction is involves reduction. (include half-equations)

A

This extraction will result in carbon displacing the zinc, to form carbon dioxide. This will involve reduction as the zinc ions gain 2 electrons (Zn2+ + 2e- –>Zn) and the zinc oxide loses oxygen. This will result in the products of zinc + carbon dioxide.

69
Q

🟢 Why is the electrolysis of calcium from molten calcium chloride a redox process

A

This is a redox reaction because you can see that Ca2+ gains two electrons, therefore, meaning it is reduced, while Cl- loses electrons as well, meaning it is being oxidised. This shows that both oxidisation and reduction are taking place in this reaction, therefore making it a redox reaction.

(GCSE) Chemistry (24 decks)

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