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Chemistry > 2 Inorganic Chemistry > Flashcards

2 Inorganic Chemistry Flashcards

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1
Q

what happens when lithium reacts with water

A

relatively slow reaction

lithium doesn’t melt

effervescence can be observed

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2
Q

what happens when sodium reacts with water

A

sodium melts into a ball as large amounts of heat energy is released

hydrogen released catches fire and causes the Na to dash around the surface of water

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3
Q

what happens when potassium reacts with water

A

creates a lilac flame

melts into a shiny ball which dashes across the surface of the water

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4
Q

what are the group 1 metals called

A

alkali metals

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5
Q

what happens to the reactivity as you go down group one

A

increases because as we go down the group each element has one more outer shell meaning the distance between the outer most shell and the nucleus increases. this decreases the electrostatic force of attraction meaning its easier to lose the one electron as less energy is required

group one elements only need to lose one electron to form a complete outer shell

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6
Q

trends in the group one / alkali metals

A

increase in reactivity as we go down the group because the number of electron shells increase

all soft and easy to cut and get softer as we move down the group

low melting points which decrease as we go down the group

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7
Q

how do the alkali metals react with oxygen

A

when the alkali metals react with oxygen they form metal oxides

this causes a dull coating on the metals

this happens quicker as we go down the group

alkali metal + oxygen -> metal oxide

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8
Q

how do the alkali metals react with water

A

reactivity increases as you go down the group

general reaction formula:

alkali metal + water -> metal hydroxide + hydrogen

The hydroxides formed all have the same general formula and are colourless, aqueous solutions

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9
Q

predicting properties in group 1

A

reactivity increases as we go down the group

melting point decreases as we go down the group

get softer as we go down the group

the density increases as we go down the group

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10
Q

colour and physical state (at room temp) of fluorine

A

yellow gas

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11
Q

colour and physical state (at room temp) of chlorine

A

pale yellow/green gas

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12
Q

colour and physical state (at room temp) of bromine

A

red/brown liquid

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13
Q

colour and physical state (at room temp) of iodine

A

grey solid

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14
Q

colour and physical state (at room temp) of astatine

A

black solid

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15
Q

trend in properties of the halogens

A

melting and boiling points increase as you go down the group because the atoms become larger and therefore the intermolecular forces get stronger

the colours of the halogens get darker in colour as you go down the group

halogens decrease in reactivity as you go down the group

the rate of reaction is slower as you go down the group

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16
Q

what is a metal halide

A

when a halogen reacts with a metal they form ionic compounds which are metal halide salts

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17
Q

what is a halide

A

a halogen ion

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18
Q

what is a displacement reaction

A

where a more reactive halogon is added to a salt (another halogen thats reacted to a metal) and the more reactive halogen displaces the less reactive halogen from the salt

Cl2(aq) + 2KI(aq) → 2KCl(aq) + I2(aq)

here chlorine displaces iodine showing how chlorine is more reactive

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19
Q

why does the reactivity decrease when you go down the halogen group

A

as you go down the group the amount mass of the halogens increase and so does the number of electron shells

halogens form negative ions (they gain an electron)

the distance from the nuclei (positive) to the outer shell increases which means there is a smaller force of attraction so it is harder to gain an electron

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20
Q

what % of nitrogen is there in the air

A

78%

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21
Q

what % of oxygen is there in the air

A

21%

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22
Q

what % of argon is there in the air

A

0.9%

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23
Q

what % of CO2 is there in the air

A

0.04%

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24
Q

how to determine the percentage by volume of oxygen in air using experiments involving the reactions of metals (e.g. iron) with air

A
  1. place an excess of wet iron filings into a conical flask (100cm3) full of air
  2. place a bung on the top of the conical flask connected to a syringe

over time the iron will rust (react with oxygen) until all of the oxygen is used up. this will move the syringe reading from 100 to 79ish showing oxygen makes up 21% of the air

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25
Q

how to determine the percentage by volume of oxygen in air using experiments involving the reactions of non-metals (e.g. phosphorus) with air

A
  1. place a small amount of phosphorus onto a evaporating dish in a bell jar which is sitting in a trough of water
  2. set the phosphorus on fire

As the phosphorus burns it uses up the oxygen inside the bell jar and the water level rises

By making careful measurements of water levels before and after the experiment you can determine the percentage of oxygen in the air

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26
Q

combustion of magnesium in oxygen

A

intense white flame
white powder produced

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27
Q

combustion of hydrogen in oxygen

A

exothermic
water is produced

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28
Q

combustion of sulpher in oxygen

A

blue flame
colourless, poisonous gas produced

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29
Q

what is thermal decomposition

A

a reaction where a substance breaks down due to the action of heat

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30
Q

what is the formation of carbon dioxide from the thermal decomposition of metal carbonates

A

metal carbonate → metal oxide + carbon dioxide

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31
Q

what is the happens when copper(II) carbonate is thermally decomposed

A

copper(II) carbonate → copper(II) oxide + carbon dioxide

green -> black

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32
Q

what happens if we increase the amount of CO2 in the atmosphere

A

contribute to the enhanced greenhouse effect and climate change

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33
Q

what type of gas is CO2

A

a greenhouse gas

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34
Q

practical: determine the approximate percentage by volume of oxygen in air using a metal or a non-metal

A
  1. place an excess of wet iron filings into a conical flask (100cm3) full of air
  2. place a bung on the top of the conical flask connected to a syringe

over time the iron will rust (react with oxygen) until all of the oxygen is used up. this will move the syringe reading from 100 to 79ish showing oxygen makes up 21% of the air

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35
Q

what is the order of the reactivity series

A

potassium, sodium, lithium, calcium, magnesium, aluminium, carbon, zinc, iron, lead, hydrogen, copper, silver, gold

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36
Q

how can metals be arranged into a reactivity series based on their reaction with water

A

whichever reacts most violently with water is the most reactive

for example potassium is the most reactive with water because it is the top of the reactivity series

metal + water ⟶ metal hydroxide + hydrogen

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37
Q

how can metals be arranged into a reactivity series based on their reaction with dilute acids

A

whichever reacts most violently with dilute acid is the most reactive

for example potassium is the most reactive with dilute acid because it is the top of the reactivity series

only elements more reactive than copper will react with acid

metal + acid ⟶ salt + hydrogen

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38
Q

how can metals be arranged in a reactivity series based on their displacement reactions between metal and metal oxides

A

a more reactive metal will displace a less reactive metal from a compound

to do this in metal oxides you can heat it

for example more reactive zinc will displace less reactive copper
zinc + copper(II) oxide → zinc oxide + copper

we can repeat this reaction with different metals to determine which is more / less reactive
so potassium will displace everything as it as the most reactive metal

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39
Q

how does displacement work with the reactivity series

A

a more reactive metal will displace a less reactive metal from a compound

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40
Q

how can metals be arranged in a reactivity series based on their displacement reactions between metal and aqueous solutions of metal salts

A

a more reactive metal will displace a less reactive metal from a compound so

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 copper(II)sulfate solution:
Mg + CuSO4→ MgSO4 + Cu

The blue colour of the CuSO4 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

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41
Q

what conditions does iron need to rust

A

water and oxygen

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42
Q

what is rust

A

hydrated iron oxide

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43
Q

how to stop iron rusting

A

barrier methods

galvanising

sacrificial protection.

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44
Q

what is the barrier method to stop iron rusting

A

the barrier such as paint or oil stop the oxygen reaching the iron and therefore stopping rusting

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45
Q

what is galvanising to stop iron rusting

A

by coating the iron in a more reactive metal (such as zinc) means that more reactive metal will react with oxygen and water instead of iron stopping it from rusting

If the coating is damaged or scratched, the iron is still protected from rusting because more reactive preferentially corrodes as it is higher up the reactivity series than iron

The iron stays protected as it accepts the electrons released by more reactive metal, remaining in the reduced state and thus it does not undergo oxidation

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46
Q

what is OILRIG

A

oxidation is loss of electrons

reduction is gain of electrons

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47
Q

what is oxidation

A

Oxidation is any reaction in which a substance gains oxygen

oxidation is loss of electrons

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48
Q

what is reduction

A

Reduction is a reaction in which a substance loses oxygen

reduction is gain of electrons

49
Q

what is a redox reaction

A

a reaction where reduction and oxidation happen (loss and gain of electrons)

Oxidation cannot occur without reduction happening simultaneously, hence these are called redox reactions

50
Q

what is a oxidising agent

A

makes oxidation happen

so it is the thing that is being reduced as it will take the electrons from the thing being oxidised allowing oxidation to happen

the thing that is reduced is the oxidising agent

51
Q

what is a reducing agent

A

makes reduction happen

so it is the thing that is being oxidised as it will gain the electrons from the thing being reduced allowing reduction to happen

the thing that is oxidated is the reducing agent

52
Q

practical: investigate reactions between dilute hydrochloric and sulfuric acids and metals

A
  1. place 3 test tubes in a rack
  2. Using a small measuring cylinder, add 5 cm3 of dilute hydrochloric acid to each of three test tubes
  3. Add about 1 cm length of magnesium ribbon to the first tube, observe and note down what you see
  4. Use a lit splint to test for any gases given off
  5. To the second test tube add a few pieces of iron filings and to the third some zinc turnings
  6. Observe what happens, test for any gases and note down your observations
    7.Repeat the experiment with dilute sulfuric acid

The metals can be ranked in reactivity order Mg > Zn > Fe

53
Q

where are metals extracted from

A

from ores found in the Earth’s crust

54
Q

what are unreactive metals found as

A

often found as the uncombined element as they do not react easily so they don’t combine with other things

55
Q

what is an ore

A

a naturally occurring rock which contains metals or metal compounds (eg iron oxide) in sufficient amounts to make them worthwhile extracting them

56
Q

how is the method of extraction of a metal from a ore related to its reactivity

A

a more reactive metal can displace a less reactive metal

the metals (aluminium) more reactive than carbon cannot be displaced by carbon so they need to be displaced by electrolysis

metals (iron) less reactive then carbon can be displaced by carbon so they can be extracted by carbon extraction

silver and gold are unreactive so they don’t react with oxygen or anything else to form compounds which means they exist by themselves and therefore don’t need to be extracted

57
Q

what are the uses of aluminium

A

aeroplane bodies - very strong and low density (so it doesn’t weigh a lot)

power cables - good electrical conductor

food cans - non-toxic, resistant to corrosion and acidic food

window frames - resistant to corrosion

58
Q

what are the uses of copper

A

wires - great electrical conductor

cooking pans - great heat conductor

water pipes - unreactive and malleable

59
Q

what are the uses of iron

A

making steel - steel is more useful

60
Q

what are the uses of low-carbon steel

A

ships, cars, bridges - strong and malleable

61
Q

what are the uses of high-carbon steel

A

tools (knives + screwdrivers) - strong and less malleable so won’t change shape

62
Q

what are the uses of stainless steel

A

cutlery, kitchen sinks - Cr forms oxide layers that is resistant to corrosion and stays shiny and clean

63
Q

what does malleable mean

A

can be hammered into different shapes without cracking

64
Q

what does ductile mean

A

can be drawn into wires

65
Q

what is low-carbon steel made of

A

Fe + 0.1% C

66
Q

what is high-carbon steel made of

A

Fe + 1% C

67
Q

what is stainless steel made of

A

Fe + 1% C + 10% Cr

68
Q

what is an alloy

A

a mixture of a metal and one or more elements, usually other metals or carbon

69
Q

why are alloys harder than pure metals

A

the different-sized atoms/ions which distorts the regular arrangement therefore preventing the layers of metal ions from sliding over each other making it harder

70
Q

things to comment on for electrolysis

A

large amounts of electricity are required making it very expensive

71
Q

things to comment on for carbon extraction

A

much cheaper than electrolysis and is a source of heat as well

72
Q

test for hydrogen

A

place a lit splint into hydrogen and you will hear a squeaky pop

73
Q

test for oxygen

A

place a glowing splint into oxygen and it will relight

74
Q

test for carbon dioxide

A

bubble through limewater and it will turn cloudy

75
Q

test for chlorine

A

place a damp blue litmus paper into chlorine and it will bleach (turn white)

76
Q

whats an anion

A

a negative ion

77
Q

test for Cl-, Br-, I-

A

silver nitrate solution

dissolve the sample (which you are testing for the anion) in water
add nitric acid (prevents false positives)
add silver nitrate solution

if Cl, Br, I are present then a precipitate will form

Cl - white precipitate
Br - cream precipitate
I - yellow precipitate

78
Q

test for carbonates

A

hydrochloric acid

add the solution into a test tube with a delivery tube connected to another test tube
add hydrochloric acid to the solution
if the solution effervesces then the gas produced will be captured
test the captured gas for CO2 in limewater
if carbonate was present then the limewater will turn cloudy

79
Q

test for sulphates

A

barium chloride

add hydrochloric acid to the solution and look for effervescence
if there’s not effervescence then add barium chloride (stops false positives)
if theres sulphates present then a white precipitate will form

80
Q

test for water

A

anhydrous copper(II) sulfate

add the solution to anhydrous copper(II) sulfate
if water is present then the solution will turn from white to blue

81
Q

physical test to show water is pure

A

check boiling point

gently heat the solution to 100’c
if the water is pure then it will all boil at 100’c

82
Q

litmus indicator

A

red in acid
blue in alkali

83
Q

phenolphthalein indicator

A

colourless in acid
pink in alkali

84
Q

methyl orange

A

red in acid
yellow in alkali

85
Q

universal indicator colours

A

0-3 strongly acidic - dark red/red/orange
4-6 weakly acidic - orange/yellow/light green
7 neutral - green
8-10 weakly alkaline - dark green/light blue
11-14 strongly alkaline - dark blue/purple

add a few drops to a solution and match the colour it turns to the universal indicator chart to find out the pH

86
Q

what happens if an acid is added to a alkali

A

a neutralisation reaction

the H+ in the acid reacts with the OH- in a base forming water

87
Q

solubility of common sodium, potassium, ammonium compounds

A

they are soluble with everything

88
Q

solubility of nitrates

A

they are soluble with everything

89
Q

solubility of common chlorides

A

soluble with everything except silver and lead(II)

90
Q

solubility of common sulfates

A

soluble with everything except barium, calcium, lead(II)

91
Q

solubility of common carbonates

A

insoluble except with sodium, potassium, and ammonium

92
Q

solubility of common hydroxides

A

insoluble except with sodium, potassium, and calcium (calcium hydroxide is slightly soluble)

93
Q

what is a acid

A

a proton (H+ ion) donor

94
Q

what is a base

A

a proton (H+ ion) acceptor

95
Q

why do metals less reactive the hydrogen not react with acids

A

they can’t displace the H so can’t react

96
Q

reaction between metals and acids

A

Acid + Metal → Salt + Hydrogen

a salt is formed because the H is displaced from the acid

observations:
the metal will disappear as it is used up in the reaction
there is fizzing because a gas (H) is produced

97
Q

reaction between bases and acids

A

Acid + Base → Salt + Water

observations:
the base disappears as it is used up in the reaction

98
Q

reaction between metal carbonates and acids

A

Acid + Carbonate → Salt + Water + Carbon dioxide

observations:
the carbonate disappears because it is used up in the reaction
there is fizzing because a gas (CO2) is being produced

99
Q

whats an alkali

A

a OH- donor
bases that are soluble in water

100
Q

examples of bases

A

a metal oxide, metal hydroxide and ammonia

101
Q

describe an experiment to prepare a pure, dry sample of a soluble salt, starting from an insoluble reactant

A

solid metal, insoluble base or insoluble carbonate

metal/metal oxide/metal carbonate + acid –> salt + water

  1. gently heat acid in a beaker until warm
  2. add excess insoluble metal/oxide/carbonate
  3. to ensure all acid is reacted stir
  4. stop heating and filter out any excess undissolved metal
  5. filter it into an evaporating basin
  6. gently heat again until crystals first appear
  7. as soon as they appear stop heating
  8. leave evaporating basin in a warm place in the sun for a few days to allow water to evaporate
102
Q

describe an experiment to prepare a pure, dry sample of a soluble salt, starting from an acid and alkali

A

acid + alkali (soluble base) –> salt + water

titration
1. the alkali and a few drops of indicator to a conical flask
2. add acid to a burette and note the starting volume
3. slowly add the acid to the alkali until the indicator changes colour
4. record the final volume of acid and calculate the total volume of acid added
5. get the same amount of alkali without the indicator and add the total volume of acid added to the other one (this creates pure crystals without the indicator in them)
5. heat the salt solution until crystals just start to appear
6. as soon as they appear, remove from heat
7. leave evaporating basin in a warm place in the sun for a few days to allow water to evaporate

103
Q

describe an experiment to prepare a pure, dry sample of an insoluble salt, starting from two soluble reactants

A

soluble salt 1(aq) + soluble salt 2(aq) –> insoluble salt(s) + soluble salt 3(aq)

  1. mix the 2 soluble salts together in water using a stirring rod
  2. filter the solution using a funnel and filter paper
  3. pour distilled water through the filter paper to wash the insoluble salt
  4. place the insoluble salt in a warm dry sunny place to dry the salt
104
Q

practical: prepare a sample of pure, dry hydrated copper(II) sulfate crystals starting from copper(II) oxide

A

soluble salt from insoluble reactant method

  1. Add 50 cm3 dilute acid into a beaker and warm gently using a Bunsen burner
  2. add the copper(II) oxide slowly to the hot dilute acid and stir until the base is in excess (i.e. until the base stops
    dissolving and a suspension of the base forms in the acid)
  3. Filter the mixture into an evaporating basin to remove the excess base
  4. Gently heat the solution in a water bath or with an electric heater to evaporate the water and to make the
    solution saturated
  5. Check the solution is saturated by dipping a cold glass rod into the solution and seeing if crystals form on the
    end
  6. Leave the filtrate in a warm sunny place to dry and crystallise
105
Q

practical: prepare a sample of pure, dry lead(II) sulfate

A

insoluble salt from 2 soluble reactants

  1. Measure out 25 cm3 of 0.5 mol dm3 lead(II)nitrate solution and add it to a small beaker
  2. Measure out 25 cm3 of 0.5 mol dm3 of potassium sulfate add it to the beaker and mix together using a stirring
    rod
  3. Filter to remove precipitate from mixture
  4. Wash filtrate with distilled water to remove traces of other solutions
  5. leave in sunny warm place to dry
106
Q

what is a precipitate

A

a solid formed when 2 solutions are mixed

107
Q

test for ammonia

A

hold damp red litmus paper in the NH3
will turn blue

108
Q

colour of Li+ flame

A

(bright) red

109
Q

colour of Na+ flame

A

yellow

110
Q

colour of K+ flame

A

lilac

111
Q

colour of Ca2+ flame

A

orange/red (brick red)

112
Q

colour of Cu2+ flame

A

blue/green

113
Q

method to carry out the flame test

A
  1. turn the bunsen burner onto non-luminous
  2. get a nichrome wire loop and dip it into HCl
  3. Dip the nichrome wire into the substance
  4. hold the wire into the blue cone in the bunsen
  5. observe the colour change and repeat
114
Q

test for ammonium cation

A
  1. place 1cm of ammonia solution into a test tube
  2. add 2cm of sodium hydroxide solution into the test tube
  3. place the test tube into a beaker containing warm water
  4. hold damp red litmus over the test tube

turns from red to blue

115
Q

test for Cu2+ cation

A
  1. place 1cm of the salt (Cu) solution in a test tube
  2. slowly add drops of 0.5mol sodium hydroxide into test tube until precipitate is visible

light blue gelantinous precipitate is formed
floats on surface

116
Q

test for Fe2+ cation

A
  1. place 1cm of the salt (Fe) solution in a test tube
  2. slowly add drops of 0.5mol sodium hydroxide into test tube until precipitate is visible

green gelantinous precipitate is formed
floats on the surface

117
Q

test for Fe3+ cation

A
  1. place 1cm of the salt (Fe) solution in a test tube
  2. slowly add drops of 0.5mol sodium hydroxide into test tube until precipitate is visible

red/brown gelantinous precipitate is formed
floats on the surface

118
Q

how to carry out an acid-alkali titration

A

Titrations are a method of analysing the concentration of solutions
They can determine exactly how much alkali is needed to neutralise a quantity of acid – and vice versa

Use the pipette and pipette filler and place exactly 25 cm3 known acid solution into the conical flask
fill the burette with an alkali
Place the conical flask on a white tile so the tip of the burette is inside the flask
Add a few drops of phenolphthalein indicator to the solution in the conical flask
Perform a rough titration by taking the burette reading and running in the solution in 1 – 3 cm3 portions, while swirling the flask vigorously
Quickly close the tap when the end-point is reached (sharp colour change) and record the volume, placing your eye level with the meniscus
Now repeat the titration with a fresh batch of acid
As the rough end-point volume is approached, add the solution from the burette one drop at a time until the indicator just changes colour
Record the volume to the nearest 0.05 cm3
Repeat until you achieve two concordant results (two results that are within 0.1 cm3 of each other) to increase accuracy

Chemistry (5 decks)

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