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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
what do alkalis ionise to produce and in what conditions
in aqueous solutions, alkalis ionise to produce OH⁻ ions (hydroxide ions)
26
acid + metal oxide OR metal hydroxide →
salt + water e.g. sulfuric acid + copper(II) oxide → copper(II) sulfate + water
27
acid + metal carbonate ->
salt + water + carbon dioxide e.g. nitric acid + copper(II) carbonate → copper(II) nitrate + water + carbon dioxide
28
pH of acids, alkalis and neutral in aqueous solutions
ACID: between 0 and 6 NEUTRAL: 7 ALKALI: between 8 and 14
29
describe pH in terms of ions
- 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
30
how can soluble salts be made from acids
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
31
what happens in neutralisation reactions
in neutralisation reactions between an acid an alkali, hydrogen ions react with hydroxide ions to produce water, whereby the solution becomes pH7.
31
neutralisation general ionic
H⁺ (aq) + OH⁻ (aq) → H₂O (l)
32
what can we measure pH using
a pH probe or universal indicator
33
why are pH probes more advantageous than UI
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
pH colours
pH 0-2 : red pH 3-5: yellow pH 6-8: green pH 9-11: blue pH 12-14: purple
35
what are titrations used to measure
the volumes of acid and alkali solutions that react with each other in order to neutralise eachother
36
define a strong acid
an acid that completely ionises in aqueous solutoon
37
examples of strong acids
hydrochloric, nitric and sulphuric acids
38
define a weak acid
an acid that is only partially ionised in aqueous solution
39
examples of weak acids
ethnic, citric and carbonic acids
40
describe pH 7 in terms of ions
there is an equal concentration of H⁺ and OH⁻ ions
41
why do stronger acids have a lower pH (more acidic) than a weaker acid of the same concentration
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
define pH
a measure of the H⁺ concentration; the higher the concentration, the lower the pH
43
define dilute
a solution that contains a relatively small amount of dissolved solute
44
define concentrated
a solution that contains a relatively large amount of dissolved solute
45
how can you identify a weak acid equation
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
describe the relationship between the change in pH and the concentration of H⁺ ions
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
why do weaker acids have a higher pH (more alkali) than a stronger acid of the same concentration
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
how can an acid be both dilute and strong
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
what does acid strength tell you
what proportion of the acid molecules ionise in water
50
what happens to the pH regardless of strength
it will decrease with increasing acid concentration
51
how can you convert concentration in mol/dm³ to g/dm ³
multiply the concentration in mol/dm³ by the relative formula mass
52
why can solid ionic compounds not conduct electricity
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
define an electrolyte
the ionic liquid or solution broken down by electrolysis
54
what happens when an ionic compound is melted or dissolved in water
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
describe the process of electroysis
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
define electrolysis
the process of breaking down compounds using electricity
57
define electrodes
the rods that conduct electricity which come in pairs; one negative, one positive
58
describe the electrolysis of aluminium oxide
- 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
diatomic molecules
Have No Fear Of Ice Cold Beer H - hydrogen N - nitrogen F - fluorine O - oxygen I - iodine Cl - chlorine Br - bromine
60
what two ways can we extract metals from their compounds
- reduction with carbon (displacement reactions) - electrolysis
61
why do we not reduce metals with other metals
it would be too expensive
62
pros and cons of reducing metals using carbon
PROS: - cheap - requires less energy CONS: - only works for metals less reactive than carbon because carbon needs to be able to displace it
63
why must the anode be replaced regularly
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
when is electrolysis used
when the metal is too reactive to be extracted by reduction with carbon or if the metal reacts with carbon
65
why is extracting aluminium expensive
- 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
why are the electrodes made of graphite
because carbon conducts electricity due to the delocalised electron, and it has a high melting point so can withstand the heat.
66
why are the electrodes made of graphite
because carbon conducts electricity due to the delocalised electron, and it has a high melting point so can withstand the heat
67
why is a mixture used as the electrolyte when electrolysing aluminium oxide
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
what are the electrodes used in electrolysis made of and why
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
what does the cathode produce
it produces hydrogen if the ⁺ ions are more reactive than hydrogen. if they're less reactive than hydrogen, the metal is produced
70
what does the anode produce
it produces oxygen unless the solution contains halide ions; if it does, then the halogen is produced
71
reactivity series + acronym
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
why do the anodes and cathodes discharge ions depending on the relative reactivity of the elements involved
because in the aqueous solution, water molecules break down, producing hydrogen ions and hydroxide ions that are discharged
73
what do the ions travelling to the electrodes create
a flow of charge through the electrolyte
74
describe what happens during a displacement reaction
the more reactive metal gradually disappears as it forms a solution; the less reactive metal coats the surface of the more reactive metal
75
what is always reduced at the cathode and why
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
half equation for oxygen gas produced at the anode
4OH⁻ - 4e⁻ → 2H₂O + 4e⁻
77
half equation for hydrogen gas produced at the cathode
2H⁺
78
half equation for hydrogen gas produced at the cathode
2H⁺ + 2e⁻ → H₂

chem spec (12 decks)

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