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chemical changes Flashcards

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1
Q

metal + oxygen ->

A

metal oxides
e.g., magnesium + oxygen -> magnesium oxide

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2
Q

what are oxidation reactions

A

reactions where metals react with oxygen; this is because the metals gain oxygen

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3
Q

what is oxidation and what is reduction

A

oxidation -> both the gaining of oxygen and the loss of electrons
reduction -> both the losing of oxygen and the gaining of electrons)

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4
Q

what are redox reactions

A

reactions in which both reduction and oxidation happen at the same time

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5
Q

what is the reactivity of a metal related to

A

its tendency to form positive ions; the easier they lose electrons to form positive ions, the more reactive they are.
i.e. a more reactive metal will more easily lose its outer electrons to form a positive ion than a less reactive metal

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6
Q

how can we test metals’ reactivity

A

by reacting the metals with water and with dilute acids

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7
Q

metal + water →

A

metal hydroxide + hydrogen
e.g. magnesium + water → magnesium hydroxide + hydrogen

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8
Q

how can you tell the reactivity of metals above calcium

A

by testing a range of different metals reacting with water, we can work out a reactivity series from most reactive to least reactive; we can tell the relative reactivity by comparing how vigorous the reaction with water is. to quantify this, we can test the temperature change and the rate of production of hydrogen gas. these values allow us to compare the reactivity of different metals via their reactions with water

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9
Q

what is the problem with comparing metal reactivities by reactions with water and how can you resolve this

A

some metals that are less reactive than calcium do not actually react with water at all. to compare their relative reactivities, they’re reacted with dilute acids.

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10
Q

metal + dilute acid →

A

salt + hydrogen
e.g. magnesium + hydrochloric acid → magnesium chloride + hydrogen

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11
Q

how can you tell the reactivity of metals below calcium

A

as long as the metal is more reactive than hydrogen, it will have a reaction with dilute acids - this means that we can compare the reactivities of less reactive metals by comparing how vigorous the reactions are, the temperature change and the rate of production of hydrogen

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12
Q

how can the rate of production of hydrogen be detected

A

using a splint test (squeaky pop test) and comparing how loud the squeaky pops are
OR
using a gas syringe and comparing volume of hydrogen produced per second

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13
Q

what occurs in a displacement reaction

A

a more reactive element will displace a less reactive element from its compound

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14
Q

how are unreactive metals found

A

unreactive (native) metals such as gold are found in the Earth as the pure metal itself; this means that these metals don’t need to be chemically extracted because they do not easily react with other elements in the ground e.g. oxygen

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15
Q

how are reactive metals found

A

reactive metals like iron and copper are found as compounds (e.g. iron oxide) that require chemical reactions to extract the metal

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16
Q

define an ore

A

a rock containing enough metal to make it economic to extract the metal

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17
Q

what do acids ionise to produce and in what conditions

A

in aqueous solutions, acids ionise to produce H⁺ ions (hydrogen ions)

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18
Q

what does hydrochloric acid produce

A

salts called chlorides

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19
Q

what does sulphuric acid produce

A

salts called sulfates

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20
Q

what does nitric acid produce

A

salts called nitrates

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21
Q

the greater the difference in reactivity between the acid and hydrogen

A

the faster it reacts with acids

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22
Q

define bases

A

any chemical that can neutralise acids to produce a salt and water

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23
Q

examples of bases

A
  • insoluble metal hydroxides and metal oxides e.g., copper oxide, sodium hydroxide
  • metal carbonates
  • alkalines
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24
Q

define an alkali and give an example

A

a soluble base e.g. sodium hydroxide, which can dissolve in water and can neutralise acids to produce a salt and water. this makes sodium hydroxide an alkali and therefore also a base

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25
Q

what do alkalis ionise to produce and in what conditions

A

in aqueous solutions, alkalis ionise to produce OH⁻ ions (hydroxide ions)

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26
Q

acid + metal oxide OR metal hydroxide →

A

salt + water
e.g. sulfuric acid + copper(II) oxide → copper(II) sulfate + water

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27
Q

acid + metal carbonate ->

A

salt + water + carbon dioxide
e.g. nitric acid + copper(II) carbonate → copper(II) nitrate + water + carbon dioxide

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28
Q

pH of acids, alkalis and neutral in aqueous solutions

A

ACID: between 0 and 6
NEUTRAL: 7
ALKALI: between 8 and 14

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29
Q

describe pH in terms of ions

A
  • the lower the pH, the more acidic, meaning the higher the concentration of H⁺ ions
  • the higher the pH, the more alkaline, meaning the higher the concentration of OH⁻ ions
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30
Q

how can soluble salts be made from acids

A

by reacting them with solid insoluble substances, such as metals, metal oxides, hydroxides or carbonates. the solid is added to the acid until no more reacts and the excess solid is filtered off to produce a solution of the salt

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31
Q

what happens in neutralisation reactions

A

in neutralisation reactions between an acid an alkali, hydrogen ions react with hydroxide ions to produce water, whereby the solution becomes pH7.

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31
Q

neutralisation general ionic

A

H⁺ (aq) + OH⁻ (aq) → H₂O (l)

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32
Q

what can we measure pH using

A

a pH probe or universal indicator

33
Q

why are pH probes more advantageous than UI

A

pH probes determine pH electronically, which is much more accurate and has a higher resolution than using universal indicator, which is subjective to the observer, is qualitative and has a very low resolution only used to !estimate! the pH

34
Q

pH colours

A

pH 0-2 : red
pH 3-5: yellow
pH 6-8: green
pH 9-11: blue
pH 12-14: purple

35
Q

what are titrations used to measure

A

the volumes of acid and alkali solutions that react with each other in order to neutralise eachother

36
Q

define a strong acid

A

an acid that completely ionises in aqueous solutoon

37
Q

examples of strong acids

A

hydrochloric, nitric and sulphuric acids

38
Q

define a weak acid

A

an acid that is only partially ionised in aqueous solution

39
Q

examples of weak acids

A

ethnic, citric and carbonic acids

40
Q

describe pH 7 in terms of ions

A

there is an equal concentration of H⁺ and OH⁻ ions

41
Q

why do stronger acids have a lower pH (more acidic) than a weaker acid of the same concentration

A

because strong acids fully ionise in aqueous solutions (all of the molecules ionise to release H⁺ ions) so will have a higher concentration of H⁺ ions than a weak acid of the same concentration; this means there are more frequent collisions between reactant particles. hence, higher concentration of H⁺ results in a lower pH

42
Q

define pH

A

a measure of the H⁺ concentration; the higher the concentration, the lower the pH

43
Q

define dilute

A

a solution that contains a relatively small amount of dissolved solute

44
Q

define concentrated

A

a solution that contains a relatively large amount of dissolved solute

45
Q

how can you identify a weak acid equation

A

unlike strong acid equations, the ⇌ symbol is used in the equation to show that the reaction is a reversible reaction and does not go to completion

46
Q

describe the relationship between the change in pH and the concentration of H⁺ ions

A

as the pH decreases by one unit, the H⁺ concentration increases by a factor of 10
e.g. at pH 0 → concentration of H⁺ ions = 1
e.g. at pH 1 → concentration of H⁺ ions = 0.1
etc.

47
Q

why do weaker acids have a higher pH (more alkali) than a stronger acid of the same concentration

A

because weak acids only partially ionise in aqueous solutions, meaning that only some of the molecules ionise to release H+ ions. this means that weak acids have a lower concentration of H+ ions than strong acids of the same concentration; this means there are less frequent collisions between reactant particles. hence, lower concentration of H⁺ results in a higher pH

48
Q

how can an acid be both dilute and strong

A

dilute because there are not many acid molecules present, but strong because a very high proportion of the acid molecules that are present ionise to release H+ ion

49
Q

what does acid strength tell you

A

what proportion of the acid molecules ionise in water

50
Q

what happens to the pH regardless of strength

A

it will decrease with increasing acid concentration

51
Q

how can you convert concentration in mol/dm³ to g/dm ³

A

multiply the concentration in mol/dm³ by the relative formula mass

52
Q

why can solid ionic compounds not conduct electricity

A

because the ions are fixed in place by strong electrostatic forces of attraction, so the ions aren’t free to move and carry charge

53
Q

define an electrolyte

A

the ionic liquid or solution broken down by electrolysis

54
Q

what happens when an ionic compound is melted or dissolved in water

A

the ions are free to move about within the liquid or solution, and these liquids or solutions are able to conduct electricity (these are electrolytes)

55
Q

describe the process of electroysis

A

an electric current is passed through electrolytes, which causes the ions to move to the electrodes. positively charged ions are attracted to the negative electrode (the cathode) and negatively charged ions are attracted to the positive electrode (the anode). ions are discharged at the electrodes, producing elements

56
Q

define electrolysis

A

the process of breaking down compounds using electricity

57
Q

define electrodes

A

the rods that conduct electricity which come in pairs; one negative, one positive

58
Q

describe the electrolysis of aluminium oxide

A
  • aluminium oxide is melted and dissolved in cryolite to lower melting point
  • at cathode, positively charged aluminium ions are attracted. they’re reduced here and gain three electrons to become aluminium metal
  • Al³⁺ + 3e⁻ → Al
  • at anode, negatively charged oxide ions are attracted. they’re oxidised here and lose two electrons each to form oxygen gas
  • 2O²⁻ - 4e⁻ → O₂
  • the carbon anode reacts with the oxygen produced to make carbon dioxide and is used up overtime
  • C⁻ + O₂ → CO₂
  • he overall equation for the electrolysis of aluminium oxide to form aluminium and oxygen is:
    2Al₂O₃ → 4Al + 3O₂
59
Q

diatomic molecules

A

Have No Fear Of Ice Cold Beer
H - hydrogen
N - nitrogen
F - fluorine
O - oxygen
I - iodine
Cl - chlorine
Br - bromine

60
Q

what two ways can we extract metals from their compounds

A
  • reduction with carbon (displacement reactions)
  • electrolysis
61
Q

why do we not reduce metals with other metals

A

it would be too expensive

62
Q

pros and cons of reducing metals using carbon

A

PROS:
- cheap
- requires less energy

CONS:
- only works for metals less reactive than carbon because carbon needs to be able to displace it

63
Q

why must the anode be replaced regularly

A

the oxygen gas produced by the anode reacts with the carbon anode under the high temperatures to form carbon dioxide gas which is released, meaning the anode will eventually burn away

64
Q

when is electrolysis used

A

when the metal is too reactive to be extracted by reduction with carbon or if the metal reacts with carbon

65
Q

why is extracting aluminium expensive

A
  • melting compounds like aluminium oxide requires lots of energy (even with the cryolite)
  • a lot of energy is needed to produce the electric current
  • you need to replace the anodes frequently
    all of this makes it quite costly
66
Q

why are the electrodes made of graphite

A

because carbon conducts electricity due to the delocalised electron, and it has a high melting point so can withstand the heat.

66
Q

why are the electrodes made of graphite

A

because carbon conducts electricity due to the delocalised electron, and it has a high melting point so can withstand the heat

67
Q

why is a mixture used as the electrolyte when electrolysing aluminium oxide

A

aluminium oxide has a very high melting point, so to lower, we mix it with cryolite so that less energy is needed to extract aluminium, saving money

68
Q

what are the electrodes used in electrolysis made of and why

A

graphite or platinum because they don’t react with any other materials (they’re inert) and they conduct electricity due to the delocalised electrons which are free to move and carry charge

69
Q

what does the cathode produce

A

it produces hydrogen if the ⁺ ions are more reactive than hydrogen. if they’re less reactive than hydrogen, the metal is produced

70
Q

what does the anode produce

A

it produces oxygen unless the solution contains halide ions; if it does, then the halogen is produced

71
Q

reactivity series + acronym

A

Please Stop Calling Me A Careless Zebra Instead Try Learning How Copper Saves Gold

P - potassium
S - sodium
C - calcium
M - magnesium
A - aluminium
C - (carbon)
Z - zinc
I - iron
T - tin
L - lead
H - (hydrogen)
C - copper
S - silver
G - gold

72
Q

why do the anodes and cathodes discharge ions depending on the relative reactivity of the elements involved

A

because in the aqueous solution, water molecules break down, producing hydrogen ions and hydroxide ions that are discharged

73
Q

what do the ions travelling to the electrodes create

A

a flow of charge through the electrolyte

74
Q

describe what happens during a displacement reaction

A

the more reactive metal gradually disappears as it forms a solution; the less reactive metal coats the surface of the more reactive metal

75
Q

what is always reduced at the cathode and why

A

the least reactive ion, because they have lower tendencies to form an ion, meaning they have a lower tendency to remain an ion (higher reactivity means more easily becomes an ion)

76
Q

half equation for oxygen gas produced at the anode

A

4OH⁻ - 4e⁻ → 2H₂O + 4e⁻

77
Q

half equation for hydrogen gas produced at the cathode

A

2H⁺

78
Q

half equation for hydrogen gas produced at the cathode

A

2H⁺ + 2e⁻ → H₂

chem spec (12 decks)

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