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chemical changes Flashcards

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1
Q

What forms when you react 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 happens during oxidation & reduction?

A
  • oxidation is the both the gaining of oxygen and the loss of electrons
  • reduction is 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 do metals form when they react with other substances?

A
  • positive ions
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6
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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7
Q

How can metals be arranged?

A
  • in order of their reactivity in a reactivity series
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8
Q

How can we test metals’ reactivity?

A
  • by reacting the metals with water and with dilute acids
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9
Q

List the reactivity series

A
  • potassium, sodium, calcium, magnesium, aluminium - more reactive than carbon, extracted by electrolysis
  • (carbon)
  • zinc, iron, tin, lead - less reactive than carbon, extracted by reduction
  • (hydrogen)
  • copper, silver, gold - very unreactive
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10
Q

What forms when you react metal + water?

A
  • metal hydroxide + hydrogen
  • e.g. magnesium + water → magnesium hydroxide + hydrogen
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11
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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12
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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13
Q

What forms when you react metal + (dilute) acid?

A
  • salt + hydrogen
  • e.g. magnesium + hydrochloric acid → magnesium chloride + hydrogen
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14
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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15
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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16
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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17
Q

How are unreactive metals found and extracted?

A
  • unreactive (native) metals such as gold are found in the Earth as the pure metal itself; t
  • 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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18
Q

How are reactive metals found and extracted?

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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19
Q

How are metals less reactive that carbon extracted?

A
  • extracted from their oxides by reduction with carbon
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20
Q

Define an ore

A
  • a rock containing enough metal to make it economic to extract the metal
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21
Q

Explain how you would write ionic equations in terms of loss and gain of electrons

A
  • if​ ​sodium​ ​is​ ​oxidised,​ ​it​ ​has​ ​lost​ ​an​ ​electron,​ ​leaving​ ​it​ ​with​ ​a​ ​+1​ ​charge, so​ ​the​ ​ionic​ ​equation​ ​is:​ ​Na​ ​->​ ​Na+​ ​​ ​+​ ​e-​
  • if​ ​sodium​ ​+1​ ​ion​ ​is​ ​reduced,​ ​it​ ​has​ ​gained​ ​an​ ​electron,​ ​leaving​ ​it​ ​with​ ​a charge​ ​of​ ​zero,​ ​so​ ​the​ ​ionic​ ​equation​ ​is:​ ​Na​+​​ ​+​ ​e-​ ​​ ​->​ ​Na
  • remember:​ ​the​ ​charges​ ​on​ ​each​ ​side​ ​of​ ​the​ ​equation​ ​should​ ​add​ ​up​ ​to the​ ​same​ ​number
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22
Q

How can you tell which elements been oxides and which has been reduced in an equation?

A
  • e.g.​ ​2Na​ ​+​ ​2HCl​ ​->​ ​2NaCl​ ​+​ ​H2​
  • HCl​ ​is​ ​made​ ​up​ ​of​ ​H​+​​ ​and​ ​Cl-​ ​​ ​ions​ ​&​ ​NaCl​ ​is​ ​made​ ​up​ ​of​ ​Na+​ ​​ ​and​ ​Cl-​ ​​ ​ions
  • looking​ ​at​ ​just​ ​sodium:​ ​2Na​ ​->​ ​2Na+​ ,​ ​ ​so​ ​the​ ​ionic​ ​equation​ ​must​ ​be: 2Na​ ​->​ ​2Na+​ ​​ ​+​ ​2e-​ ,​ ​ ​meaning​ ​sodium​ ​has​ ​lost​ ​electrons​ ​&​ ​has​ ​been oxidised
  • looking​ ​at​ ​just​ ​chlorine:​ ​2Cl-​ ​​ ​->​ ​2Cl-​ ,​ ​ ​meaning​ ​chlorine​ ​has​ ​not​ ​been oxidised​ ​or​ ​reduced
  • looking​ ​at​ ​just​ ​hydrogen:​ ​2H​+​​ ​->​ ​H2​ ,​ ​ ​so​ ​the​ ​ionic​ ​equation​ ​must​ ​be: 2H+​ ​​ ​+​ ​2e-​ ​​ ​->​ ​H2​ ,​ ​ ​meaning​ ​hydrogen​ ​has​ ​gained​ ​electrons​ ​so​ ​has​ ​been reduced
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23
Q

Why type of reaction is acid + metal?

A
  • redox - as one substance is reduced, another is oxidised
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24
Q

Give an example of an acid + metal reaction and how to identify which substances are which using OIL RIG

A
  • e.g.​ ​2HCl​ ​+​ ​Mg​ ​->​ ​MgCl​2​​ ​+​ ​H​2
  • magnesium:​​ Mg​​ ->​ ​Mg​2+,​​​ so​​ ionic​ ​equation​​is​​ Mg​​->​​Mg2​+​​​ +​​2e-​,​​​Mg​​ has lost​ ​electrons​ ​so​ ​​Mg​ ​has​ ​been​ ​oxidised
  • hydrogen:​ ​2H+​ ​​ ​->​ ​H2​ ,​ ​ ​so​ ​ionic​ ​equation​ ​is​ ​2H+​ ​​ ​+​ ​2e-​ ​​ ​->​ ​H2​ ,​ ​ ​H​ ​has​ ​gained electrons,​ ​so​ ​​H​ ​has​ ​been​ ​reduced
  • because​ ​magnesium​ ​has​ ​been​ ​oxidised​ ​and​ ​hydrogen​ ​has​ ​been​ ​reduced in​ ​the​ ​same​ ​reaction,​ ​this​ ​is​ ​a​ ​​redox​ ​reaction
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25
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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26
Q

What are acids neutralised by?

A
  • alkalis (e.g. soluble metal hydroxides)
  • bases (e.g. insoluble metal hydroxides and oxides)
  • metal carbonates
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27
Q

What forms when you react acid + alkali or base?

A
  • salt + water
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28
Q

What forms when you react acid + metal oxide/ hydroxide?

A
  • salt + water
  • e.g. sulfuric acid + copper(II) oxide → copper(II) sulfate + water
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29
Q

What forms when you react 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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30
Q

What does the salt produced depend on?

A
  • in alkali and base reactions it depends on the acid used
  • the positive ins in the base, alkali or carbonate, i.e. the metal
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31
Q

What does hydrochloric acid produce when reacted with an alkali/base?

A
  • salts called chlorides
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32
Q

What does sulphuric acid produce when reacted with an alkali/base?

A
  • salts called sulfates
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33
Q

What does nitric acid produce when reacted with an alkali/base?

A
  • salts called nitrates
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34
Q

The greater the difference in reactivity between the acid and hydrogen….

A
  • the faster it reacts with acids
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35
Q

Define bases

A
  • any chemical that can neutralise acids to produce a salt and water
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36
Q

Give examples of bases

A
  • insoluble metal hydroxides and metal oxides e
  • e.g., copper oxide, sodium hydroxide metal carbonates
    alkaline
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37
Q

Define an alkali and give an example

A
  • a soluble base 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
  • e.g. sodium hydroxide,
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38
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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39
Q

What is the pH of acids, alkalis and neutral in aqueous solutions?

A
  • ACID: between 0 and 6
  • NEUTRAL: 7
  • ALKALI: between 8 and 14
40
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
41
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
42
Q

Explain how you would make a soluble salt from an acid

A

1) add​ ​the​ ​chosen​ ​solid​ ​insoluble​ ​substance​ ​to​ ​the​ ​acid​ ​then​ ​the​ ​solid​ ​will​ ​dissolve.
2) you​ ​know​ ​the​ ​acid​ ​has​ ​been​ ​neutralised​ ​when​ ​excess​ ​solid​ ​sinks​ ​to​ ​the​ ​bottom,
so​ ​keep​ ​adding​ ​until​ ​this​ ​happens
3) filter​ ​out​ ​excess​ ​solid​ ​leaving​ ​the​ ​salt​ ​solution,​ ​then​ ​evaporate​ ​some​ ​water,​ ​then leave​ ​the​ ​rest​ ​to​ ​evaporate​ ​slowly.

43
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.
44
Q

Name the general iconic equation for neutralisation

A
  • H⁺ (aq) + OH⁻ (aq) → H₂O (l)
45
Q

What can we measure pH using?

A
  • a pH probe or universal indicator
46
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
47
Q

Name the colours of pH scale

A
  • pH 0-2 : red
  • pH 3-5: yellow
  • pH 6-8: green
  • pH 9-11: blue
  • pH 12-14: purple
48
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
49
Q

Define a strong acid

A
  • an acid that completely ionises in aqueous solution
50
Q

Name examples of strong acids

A
  • hydrochloric, nitric and sulphuric acids
51
Q

Define a weak acid

A
  • an acid that is only partially ionised in aqueous solution
52
Q

Name examples of weak acids

A
  • ethnic, citric and carbonic acids
53
Q

Describe pH 7 in terms of ion

A
  • there is an equal concentration of H⁺ and OH⁻ ions
54
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; t
  • his means there are more frequent collisions between reactant particles.
  • hence, higher concentration of H⁺ results in a lower pH
55
Q

Define pH

A
  • a measure of the H⁺ concentration; the higher the concentration, the lower the pH
56
Q

What happens as the pH decrease by one unit?

A
  • The H⁺ conc of solution 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
57
Q

Define a dilute solution

A
  • a solution that contains a relatively small amount of dissolved solute
58
Q

Define a concentrated solution

A
  • a solution that contains a relatively large amount of dissolved solute
59
Q

How can you identify a weak acid in an 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
60
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
61
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
62
Q

What does acid strength tell you?

A
  • what proportion of the acid molecules ionise in water
63
Q

What happens to the pH regardless of strength?

A
  • it will decrease with increasing acid concentration
64
Q

Why is strong and weak not the same as concentrated and dilute?

A
  • the​ ​latter​ ​refers to​ ​the​ ​amount​ ​of​ ​substance​ ​in​ ​a​ ​given​ ​volume​ ​,​ ​whereas​ ​the​ ​former​ ​refers​ ​to the​ ​H+​ ​​ ​ion​ ​conc​ ​in​ ​aq.​ ​solutions
65
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
66
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
67
Q

Define an electrolyte

A
  • the ionic liquid or solution broken down by electrolysis
68
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)
69
Q

How can you break down a solution into elements?

A
  • by passing a current through substances that are molten solution
70
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 (cations) are attracted to the negative electrode (the cathode)
  • and negatively charged ions (anions) are attracted to the positive electrode (the anode).
  • ions are discharged at the electrodes, producing elements
71
Q

What metals are extracted by electrolysis?

A
  • metals that are more reactive than carbon are too reactive to be extracted by reduction so are extracted by electrolysis of molten compounds
72
Q

Define electrolysis

A
  • the process of breaking down compounds using electricity
73
Q

Define electrodes

A
  • the rods that conduct electricity which come in pairs;
  • one negative (cathode), one positive (anode)
74
Q

When a simple ionic compound is electrolysed in the molten sate using inert electrodes, what is produced at each electrode?

A
  • e.g​ ​lead​ ​bromide ​
  • the​ ​metal​ ​(lead)​ ​is​ ​produced​ ​at​ ​the​ ​cathode​
  • ​ ​the non-metal​ ​(bromine)​ ​is​ ​produced​ ​at​ ​the​ ​anode
  • this​ ​is​ ​because​ ​the​ ​metal​ ​is​ ​the​ ​positive​ ​ions​ ​and​ ​the​ ​non-metal​ ​is​ ​the​ ​negative ions
75
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
76
Q

State the overall equation for the electrolysis of aluminium oxide to form aluminium and oxygen

A
  • 2Al₂O₃ → 4Al + 3O₂
77
Q

Name the diatomic molecules

A
  • Have No Fear Of Ice Cold Beer
  • H - hydrogen
  • N - nitrogen
  • F - fluorine
  • O - oxygen
  • I - iodine
  • Cl - chlorine
  • Br - bromine
78
Q

What two ways can we extract metals from their compounds?

A
  • reduction with carbon (displacement reactions)
  • electrolysis
79
Q

Why do we not reduce metals with other metals?

A

it would be too expensive

80
Q

What are the 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

81
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
82
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
83
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
84
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.
85
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
86
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)
  • they conduct electricity due to the delocalised electrons which are free to move and carry charge
87
Q

What does the ions discharged when an aqueous solution is electrolysed using inert electrodes depend on?

A
  • the relative reactivity of the elements involved
88
Q

What does the cathode (negative) produce and when?

A
  • it produces hydrogen if the ⁺ ions are more reactive than hydrogen.
  • if they’re less reactive than hydrogen, the metal is produced
89
Q

Why is hydrogen produced at the cathode when the positive ions are more reactive?

A
  • the more reactive ions want the stay within the solution
90
Q

What does the anode (positive) produce and when?

A
  • it produces oxygen unless the solution contains halide or OH- ions;
  • if it does, then the halogen is produced
91
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
92
Q

What do the ions travelling to the electrodes create?

A
  • a flow of charge through the electrolyte
93
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
94
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)
95
Q

State the half equation for oxygen gas produced at the anode

A
  • 4OH⁻ - 4e⁻ → 2H₂O + 4e⁻
96
Q

State the half equation for hydrogen gas produced at the cathode

A
  • 2H⁺ + 2e⁻ → H₂

ZZ Chem 1 (5 decks)

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