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GCSE Chemistry 2 > Acids - Topic 3 > Flashcards

Acids - Topic 3 Flashcards

1
Q

Solute:

A

the substance that dissolves in the solvent (solutes can be solids, liquids or gases)

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2
Q

Solvent:

A

the liquid in which the solute dissolves in

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3
Q

Solution:

A

consists of a solvent and a dissolved solute

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4
Q

Saturated solution:

A

a solution in which no more solute will dissolve

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5
Q

Suspension:

A

consists of small particles of insoluble, undissolved solid spread throughout a liquid, suspensions are cloudy (like milk is)

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6
Q

Solubility:

A
  • the solubility of a solute is the mass of the solute in grams that will dissolve in 100g of solvent
  • it has the units g/100g of solvent
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7
Q

What happens when two solutions of soluble substances are mixed together?

A

When two solutions of soluble substances are mixed together they can possible react to form an insoluble substance which will not dissolve and be suspended in the liquid and so a precipitate will be formed - one of the two products is insoluble

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8
Q

What happens when an ionic compound dissolves in water?

A

When an ionic compound dissolves in water the ions separate and become spread out among the water molecules - the ions have no longer anything to do with each other

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9
Q

What are most salts often?

A
  • Most salts are soluble in water
  • Often when a solution of two salts is mixed, no such precipitation reaction takes place and the ions will remain dissolved in the water
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10
Q

Examples of precipitates being formed in reactions:

A
  • some salts are insoluble in water. if solutions containing those ions are mixed, the insoluble salt forms as a solid as the solutions are mixed
  • this solid is known as a precipitate, and the reaction as precipitation
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11
Q

Describe how the solubility of a substance changes with temperature:

A
  • The solubility of solutes changes with temperature
  • Solid solutes get more soluble as the temperature gets higher
  • When a hot saturated solution cools, the solvent is not able to dissolve as much solute and so solid crystals start to form
  • Gases become less soluble as the temperature increases
  • When water is heated, before it boils, bubbles of dissolved gas (mainly oxygen) form on the bottom of the container
  • As the temperature increases water is not able to dissolve as much gas
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12
Q

How can the solubility of a solutes changes with temperature be shown?

A
  • By drawing graph of solubility against temperature
  • Drawing a solubility curve
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13
Q

What are salts?

A

Salts are ionic compounds which consist of a positive ion (metal ion or ammonium ion, NH4+) and a negative ion (non-metal ion like chloride ion, Cl-, or non-metal ion radical like sulphate or nitrate)

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14
Q

What salts are soluble in water?

A
  • All potassium, sodium, and ammonium salts
  • All nitrates
  • All common halides (chlorides, iodides, bromides) except silver and lead halides
  • All common sulphates except lead (II) sulphate, calcium sulphate and barium sulphate
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15
Q

Which salts are insoluble in water?

A
  • All common carbonates except sodium, potassium and ammonium carbonates
  • All common hydroxides except sodium, potassium and ammonium hydroxides
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16
Q

Precipitate:

A

An insoluble solid that forms as a product when two solutions are mixed together

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17
Q

Distillate:

A

liquid collected by distillation

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18
Q

Filtrate:

A

liquid that passes through the filter paper

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19
Q

Residue:

A

the insoluble solid left behind on the filter paper during filtration or the solid left behind in the distillation flask

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20
Q

Halides:

A

chlorides, bromides, iodides

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21
Q

What are acids in solutions sources of?

A

acids in solutions are sources of hydrogen ions

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22
Q

What are alkalis in solutions sources of?

A

alkalis in solutions are sources of hydroxide ions

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23
Q

Acid:

A

substance that produces H+ (aq) (hydrogen ions) when dissolved in water

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24
Q

What is the use of the pH scale?

A

pH scale from 0-14 can be used to classify solutions as strongly acidic, weakly acidic, neutral, weakly alkaline, or strogly alkaline

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25
Q

What is a neutral pH?

A

pH 7

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26
Q

What is an acidic pH?

A

below pH 7

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27
Q

What is an alkali pH?

A

above pH 7

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28
Q

What pH is a strong acid and what colour is it?

A
  • pH 0 - 1
  • red
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29
Q

What pH is a weak acid and what colour is it?

A
  • pH 2, 3, 4, 5 & 6
  • pH 2 = orange
  • pH 3 & 4 = yellow
  • pH 5 & 6 = lime green
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30
Q

What pH is a neutral solution and what colour is it?

A
  • pH 7
  • green
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31
Q

What pH is a weak base/alkali and what colour is it?

A
  • pH 8, 9, 10, 11 & 12
  • pH 8 = teal
  • pH 9, 10, 11 & 12 = blue
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32
Q

What pH is a strong alkali and what colour is it?

A
  • pH 13 & 14
  • purple
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33
Q

Indicator:

A

a substance that is used to distinguish between acidic and alkaline solutions

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34
Q

Use of the indicators, litmus, phenolphthalein and methyl orange:

A

used to distinguish between acidic and alkaline solutions

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35
Q

Universal indicator:

A
  • a mixture of several different indicators which is used to measure the approximate pH value of a solution and indicates the strength of the acid or alkali
  • the strength of the acid or alkali is measured on the pH scale
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36
Q

How can indicators be used?

A

can be used either in solution or the solution can be dried onto paper to make test papers

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37
Q

Litmus solution in acid:

A

Red

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38
Q

Litmus solution in neutral solution:

A

Purple

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39
Q

Litmus solution in alkali:

A

Blue

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40
Q

Red litmus paper in acid:

A

Red

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41
Q

Red litmus paper in neutral solution:

A

Red

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42
Q

Red litmus paper in alkali:

A

Blue

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43
Q

Blue litmus paper in acid:

A

Red

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44
Q

Blue litmus paper in neutral solution:

A

Blue

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45
Q

Blue litmus paper in alkali:

A

Blue

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46
Q

Phenolphthalein in acid:

A

Colourless

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47
Q

Phenolphthalein in neutral solution:

A

Colourless

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48
Q

Phenolphthalein in alkali:

A

Pink

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49
Q

Methyl orange in acid:

A

Red

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50
Q

Methyl orange in neutral solution:

A

Orange

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51
Q

Methyl orange in alkali:

A

Yellow

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52
Q

Universal Indicator in acid:

A

Red

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53
Q

Universal Indicator in neutral solution:

A

Green

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54
Q

Universal Indicator in alkali:

A

Purple

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55
Q

What causes a solution to be acidic or alkaline?

A
  • the higher the concentration of hydrogen ions in an acidic solution the lower the pH so the more acidic the solution is as it is the hydrogen ion that causes acidity
    • acids release H+ in a solution
  • the higher the concentration of hydroxide ions in an alkaline solution the higher the pH so the more alkaline the solution is
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56
Q

How does a solvent (e.g. water and methylbenzene) affect acidic properties?

A
  • Water is a polar solvent and methylbenzene is a non-polar solvent
  • A polar solvent will have positive and negative charge (in water, oxygen is slightly more negative and the hydrogen is slightly positive) whereas methylbenzene is netral throughout
  • HCl will dissociate in water into H+ and Cl- ions because HCl is also a polar molecule (Cl is slightly negative and H is slightly positive)
    • The Cl will be attracted to the slightly positive H in water and the H will be attracted to the slightly negative O in water.
    • The HCl molecule will dissociate as a result of these strong attractions
  • With methylbenzene, no such attraction exists, thus HCl will not dissociate and no ions will be made and HCl will remain as a HCl molecule
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57
Q

Dot and cross diagram showing HCl changes when dissolved in water:

A
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58
Q

2 equations to shows what happens when HCl is dissolved in water:

A
  1. HCl (g) ——> H+ (aq) + Cl- (aq)
  2. HCl (g) + H2O (l) ——> H3O+ (aq) + Cl- (aq)
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59
Q

Difference between hydrogen chloride gas and hydrochloric acid in particles:

A
  • HCl gas consists of HCl molecules
  • hydrochloric acid consists of H+ and Cl- ions in solution
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60
Q

What does pH stand for?

A
  • ‘p’ stands for potenz (German word for power)
  • H stands for H+
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61
Q

What type of scale is the pH scale?

A

logarithmic scale (like Richter scale and loudness in decibels)

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62
Q

What is a pH probe?

A
  • used to measure pH of solution to determine pH to 1 or 2 decimal places - more accurate
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63
Q

By what factor does hydrogen ion concentration solution increase/decrease when the pH increases/decreases by 1?

A
  • as hydrogen ion concentration in a solution increases by a factor of 10, the pH of the solution decreases by 1 and vice versa
  • e.g. solution of pH1 contains 10x as many hydrogen ions than a solution that has pH2 (10x as acidic) & pH 1 is 10ox more acidic than pH 3
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64
Q

How are solutions of acids made?

A
  • solutions of acids made by dissolving certain covalent molecules in water
  • the water cause the covalent molecules (as it is a polar substance) to break apart in a way that generates hydrogen ions (H+)
65
Q

What does concentrated and dilute refer to?

A

concentrated and dilute refers to amount of acid molecules dissolved in water

66
Q

Concentrated:

A

larger amount of substance in given volume of solution

67
Q

Dilute:

A

lesser amount of substance in given volume of solution

68
Q

What will a concentrated acid have?

A
  • concentrated solution will have large amount of acid molecules dissolved in a certain volume of water
69
Q

What will a dilute acid have?

A
  • dilute acid will have relatively few acid molecules dissolved in the water
70
Q

If a solution of ethanoic acid contains more hydrogen ions than a solution of hydrochloric acid what can you say about their concentrations?

A
  • ethanoic acid has lower pH than HCl as more of HCl molecules dissociated and so released more hydrogen ions
  • need a high conc of ethanoic acid and a dilute hydrochloric acid for ethanoic acid to contain more hydrogen ions
71
Q

What is the difference between dilute and concentrated & strong and weak?

A
  • Strong and weak is NOT the same as concentrated and dilute
  • concentrated and dilute depends on the amount of acid (substance)
  • strong or weak depends on how many acid molecules dissociate (refers the H+ ion conc. in aq. solutions)
72
Q

What does strong and weak refer to in an acid?

A

strong or weak depends on how many acid molecules dissociate into separate ions (refers the H+ ion conc. in aq. solutions)

73
Q

What does a strong acid do in a solution?

A
  • strong acid fully dissociates in solution (e.g. water) (all acid molecules broken apart and formed separate ions)
  • acids such as HCl, H2SO4 and HNO3 are strong acids
    • HCl (g) ——> H+ (aq) + Cl-(aq)
74
Q

What does a weak acid do in a solution?

A
  • in weak acid only small proportion of molecules dissociates into ions (means only small proportion of solution mainly contains undissociated molecules and just low number of ions)
  • acids like ethanoic acid, CH3COOH are weak acids
75
Q

Why is HCl more dangerous than methanoic acid even if they both have same concentration?

A

HCl more dangerous as it is stronger acid than methanoic acid so more of its molecules dissociate into separate ions and release more H+ ions (so has a lower pH) make it more corrosive

76
Q

Explain the difference between strong and weak acids:

A
  • strong acids are more reactive than weak acids
    • faster reaction
    • more exothermic
    • due to difference in H+ ions as strong acids have more H+ ions
77
Q

Base:

A
  • substance that reacts with an acid to form a salt and water only
  • substances that can accept H+ ions/acids/metal oxides/metal hydroxides
  • all substances whether soluble or not that neutralise acids
78
Q

What are bases normally?

A

bases normally metal oxides or metal hydroxides

79
Q

Alkalis:

A
  • (sub-group of bases)
  • soluble bases
  • also neutralise acids but must be soluble in water
80
Q

What are alkalis normally?

A

soluble hydroxides - potassium, sodium and ammonium hydroxides
Soluble carbonates - potassium, sodium and ammonium carbonates

81
Q

What will all alkalis do?

A
  • all alkalis will either dissolve in water or react with water to produce hydroxide ions (OH-) in a solution
  • it is the OH- that causes alkalinity
82
Q

What do strong alkalis do?

A

strong alkalis like sodium hydroxide (NaOH) and potassium hydroxide (KOH) will completely dissolve in water and dissociate to give high conc of OH- ions

83
Q

What do weak alkalis do?

A
  • weak alkalis like ammonia will only partially dissolve / react with water and dissociate to give a low conc of OH- ions
  • ammonia undergoes a reversible reaction with water which leads to a low conc of OH- ions and so ammonia is a common alkaline solution of approx. pH = 8
84
Q

Salts:

A

ionic substances where hydrogen is replaced by other positive ions (normally metals)

85
Q

Criteria for acids:

A
  • substances that dissolve/react with water to release H+ ions
  • common acids:
    • H2SO4 - sulphuric acid
    • HCl - hydrochloric acid
    • HNO3 - nitric acid
    • H3PO4 - phosphoric acid
    • CH3COOH - ethanoic acid
86
Q

Criteria for bases:

A
  • substances that react with acids to form a salt and water (sometimes CO2 as well)
  • common bases:
    • metal oxides e.g. CaO, Al2O3, Na2O, CuO
    • metal hydroxides e.g. Ca(OH)s, NaOH, KOH
    • metal carbonates e.g. CaCO3, Na2CO3, CuCO3
87
Q

Criteria for alkalis:

A
  • substances that react with water to release OH- ions
  • common alkalis:
    • NH3 - ammonia
    • (water soluble metal hydroxides) e.g. NaOH (sodium hydroxide), KOH (potassium hydroxide), Ca(OH)2 (calcium hydroxide)
88
Q

Criteria for salts:

A
  • ionic substances formed when acids react with bases
  • common bases:
    • sulphates - made from sulphuric acid
    • chlorides - made from hydrochloric acid
    • nitrates - made from nitric acid
    • phosphates - made from phosphuric acid
    • ethanoates - made from ethanoic acid
    • citrates - made fro citric acid
89
Q

Acid + metal

A

metal + acid ——> metal salt + hydrogen

90
Q

Observations when a metal reacts with an acid:

A
  • effervescence
  • exothermic reaction
  • metal no longer visible
91
Q

Acid + base (metal oxides & hydroxides (exc. KOH and NaOH)

A

acid + base (alkali) ——> metal salt + water

92
Q

Observations when acid reacts with base:

A
  • form salt and water
  • if the base starts off as solid it will dissolve
  • exothermic —> temp change
93
Q

metal carbonate + acid

A

metal carbonate + acid ——> metal salt + water + carbon dioxide

94
Q

Observations when metal carbonate reacts with acid:

A
  • fizz and give off carbon dioxide
  • effervescence
95
Q

Ammonia + acid

A

Ammonia + acid ——> ammonium salts

96
Q

Observations when ammonia reacts with acid:

A
  • when ammonia reacts with HCl it produces white fumes of ammonium chloride
97
Q

Neutralisation:

A

a reaction between an acid and a base

98
Q

Acid-alkali neutralisation:

A
  • a reaction of hydrogen ions (H+) from the acid react with hydroxide ions (OH-) from the alkali to form water
  • H+ + OH- ——> H2O (+ salt)
99
Q

What happens when a base is added to an acid solution?

A
  • when a base is added to an acid solution the pH rises
  • eventually, when sufficient bas has been added the pH reacher 7
  • at this point, neither acid nor bases are in excess
  • all acid has been neutralised by the base
  • solution is neutral
  • as more base is added the pH increases above 7 and so solution becomes alkaline
100
Q

How can the neutralisation process be followed?

A
  • neutralisation process can be followed using UI paper or solution and also a pH metre
  • as the pH of solution increases, colour changes from red for acid to green for neutral and finally to purple for alkali
101
Q

Apparatus and materials needed to monitor the pH of a neutralisation reaction:

A
  • (ml = cm3)
  • 100ml beaker
  • 25ml measuring cylinder
  • stirring rod
  • spatula
  • pH chart for (UI) and UI paper or solution
  • pH metre
  • spotting tile
  • lime (calcium hydroxide) solid
  • vinegar (ethanoic acid)
102
Q

Method for monitoring the pH of a neutralisation reaction:

A
  1. Use measuring cylinder to pour 20ml of vinegar into beaker. Add 20ml of water to vinegar.
  2. Stir mixture with glass stirring rod. Place small piece of UI paper on spotting tile. Use glass rod to place drop of vinegar solution on indicator paper. Note the colour.
  3. Place clean tip of pH metre into solution, allow metre to settle for a value and note this down.
  4. Add 0.1g of lime to beaker. Take reading with pH metre to test the solution with fresh piece of UI indicator.
  5. Repeat steps 3 & 4.
  6. Continue with above steps until you have added 1.4g of lime.
103
Q

Graph to show change in pH of vinegar solution as calcium hydroxide added:

104
Q

What is pH of vinegar at start of the neutralisation reaction?

105
Q

Explain the change in pH during the neutralisation reaction:

A
  • the pH gradually increases till 0.8g of lime is added then there is a sudden increase in pH when 0.9 of lime is added and then the pH of the solution curves off as solution becomes saturated
  • there is an increase in pH as lime is added making the solution more alkaline increasing the pH of the solution
106
Q

Does the solution smell of vinegar at end of the neutralisation experiment? Explain answer:

A
  • doesn’t smell as all acid has been reacted
107
Q

How do you know when ethanoic acid has been neutralised in the neutralisation reaction?

A

when pH of solution reaches pH 7.0

108
Q

Why is it important to stir mixture thoroughly in neutralisation reaction?

A
  • to ensure all Ca(OH)2 (alkali) has been fully dissolved in solution so pH is consistent throughout the solution
  • to make sure acid is reacted as much as possible
109
Q

What is the best piece of apparatus to measure the volume of ethanoic acid in the neutralisation experiment and why?

A
  • measuring cylinder
    • has a higher resolution as the increase in graduations on the measuring cylinder are in smaller increments (e.g. to nearest 0.5ml) so level of accuracy increases
110
Q

What are the main errors of neutralisation experiment?

A
  • loss of mass of solid onto balance
  • Ca(OH)2 not fully dissolving in acid
111
Q

How could you improve method of neutralisation experiment?

A

use glass pipette to increase accuracy of measuring liquid

112
Q

3 methods of preparing a salt:

A
  1. excess base method
  2. titration
  3. precipitation
113
Q

When can the excess base method be used?

A
  • can only be used to make a salt if base in insoluble and salt made is soluble
114
Q

What can the excess base method not be used for?

A
  • can’t be used to make sodium, potassium and ammonium salts
  • can’t be used to make soluble salts
115
Q

Excess base method:

A
  1. Choose suitable starting materials - usually metal oxide is used as base and is reacted with appropriate acid (want to make copper sulfate - so use copper oxide and sulphuric acid
  2. Put some acid in beaker and add to it excess base (ensures all acid has been reacted). Gently warm the beaker to speed up the reaction. (if carbonate used reaction much faster and mixture doesn’t need to be warmed) - so add excess copper oxide
  3. After reaction finished, there will be a solution containing the salt and some unreacted solid base, the solution can be separated from unreacted base by filtration - copper oxide is residue and copper sulfate solution is filtrate
  4. The filtrate contains the salt solution. The solution is then heated and some of the water is evaporated away to make a hot saturated solution (this can be tested for by dipping a glass rod into the hot solution and blowing on it - if crystals appear it is saturated)
  5. Take care not to evaporate all the water away - this will result in an anhydrous powder of the salt rather than crystals (which usually contain some water of crystallisation) - if copper sulfate solution has all the water evaporated away then white anhydrous copper sulfate will be formed rather than blue hydrated copper sulfate
  6. Leave the hot saturated solution to cool and crystallise. Crystals appear as dissolved solids are less soluble in cold water than in hot water. The slower the cooling, the larger the crystals.
  7. To separate the crystals from the solution, the mixture can be filtered and then the crystals should be left to dry the warm place
116
Q

What does it mean for copper sulphate crystals made this way through excess base method to be hydrated?

A
  • contain some water molecules ‘locked’ in solid structure
  • CuO(s) + H2SO4(aq) ——> CuSO4 (aq) + H2O(l)
117
Q

How does hydrated copper sulphate differ from anhydrous copper sulphate and how could hydrated be converted to anyhdrous?

A
  • anyhdrous copper sulphate of crystallisation
  • heating the hydrated copper sulphate turns it anyhdrous
  • CuSO4 + 5H2O (blue, hydrated) ——> (heat) CuO4 + 5H2O (white, anhydrous)
118
Q

What is titration?

A

lab technique where 2 solutions mixed together to determine exact volumes of each which react together

119
Q

How does titration work?

A
  • pipette used to measure out one solution into a conical flask
  • other solution added to flask from a burette a little at a time, with mixing until the reaction is complete - end-point of the titration
120
Q

Making salts by titration:

A
  • salts formed when acid and base react together
  • when soluble bases (alkalis) are used all solutions concerned are colourless and it is impossible to tell when 2 bases have exactly neutralised each other - so precise volumes found by titration
  • method applies when making sodium, potassium and ammonium salts since these hydroxides and carbonates are all soluble in water
121
Q

Method for preparing sodium sulphate using titration:

A

Setting up the burette:

  • Rinse out the burette with distilled water, then rinse with little sulphuric acid provided
  • Clamp the burette and make sure tap is closed. Carefully pour more acid from a beaker into the top of the burette using small funnel
  • Place the beaker under the burette tap and then open the tap to fill the jet. Make sure no air bubbles are trapped in the jet.
  • Remove the funnel from top, the read initial volume from burette scale (make sure eyes level with bottom of meniscus). Record volume.

Using the pipette:

  • Collect some sodium carbonate solution in beaker.
  • Fill the pipette then measure out 25.0cm3 of sodium carbonate solution into small conical flask. repeat this 3 times into 3 separate flasks.
  • Add drops of methyl orange indicator to each flask (should be yellow)

The titration:

  • 1st titration is rough. After you have approximate value do 2 accurate titrations.
  • Stand the conical flask with the alkali and indicator in it on a white tile underneath the burette. The jet should be just inside mouth of flask.
  • Use left hand to control burette hand and right hand tot swirl flask gently as you add acid at about 1cm3 at a time. Stop adding acid when indicator changes to permanent red/pink colour. Repeat titration using fresh lot of alkali in a clean flask.
  • Use results from rough titration to anticipate where end point will be.
  • As you get nearer to end point, add acid one drop at a time until indicator orange-pink colour. Remember to record all burette readings.
  • Repeat until you get 2 titrations where volumes of acid used agree to within 0.1cm3

Making sodium sulphate crystals:

  • Now you know accurate volumes of sodium carbonate and sulphuric acid needed to neutralise each other you can mix these to make a solution of the salt
  • Pipette another 20.0cm3 of sodium carbonate solution into evaporating basin. Add it to the exact volume of sulphuric acid required from burette. (don’t use indicator)
  • Heat solution gently with Bunsen burner over tripod and gauze until you get to crystallisation point (test with cold glass rod), then set aside to cool and crystallise slowly at room temperature
122
Q

The methyl orange indicator was used to find when the solution was exactly neutralised. Why was no indicator added when the acid and alkali were mixed in the evaporating basin?

A
  1. volumes needed to make neutral solution are known
  2. an indicator would make the salts impure
123
Q

Why was the sodium sulphate heated gently before leaving it to cool (titration experiment)

A
  • to evaporate away some of water to make a hot saturated solution
124
Q

Describe the appearance of sodium sulphate crystals (titration experiment):

A

white crystals

125
Q

Method for titration to make sodium chloride:

A
  1. Collect some hydrochloric acid. Use acid to fill burette carefully, with tape closed. Fill the nozzle of the burette by running the solution through the tap. Record the start volume reading on the burette.
  2. Collect some sodium hydroxide in a different labelled 100cm3 glass beaker. Measure exactly 25.0cm3 of acid using pipette filler. Run this acid into clean conical flask.
  3. Add 5 drops of phenolphthalein in different labelled 100cm3 beaker. Indicator is pink at start (pH is very low) but when solution is acidic it will change to colourless.
  4. Start running the acid into the conical flask from the burette. After adding 23cm3 add acid one drop at a time, swirling contents of flask to mix them. When one drop of acid turns mixture colourless stop adding acid. Record final volume on burette and use this value and the starting value to work out the amount of acid added.
  5. Refill the burette and clean out the conical flask with water. Repeat whole experiment using same volumes.
  6. When finished pour solution containing salt into large beaker
  7. Salt can be obtained from solution by heating solution evaporating away some of water and then leaving it to cool and crystallise.
126
Q

When is titration used?

A

when soluble salts are prepared from an acid and soluble reactant

127
Q

Why is titration used rather than the excess base method?

A
  • excess base method relies upon the facts the base and salt can be easily separated by filtration
  • while this is largely true for most bases and their corresponding salt, it is not always the case
  • sodium, potassium and ammonium compounds are always soluble
  • means that the sodium, potassium and ammonium bases can’t be separated from corresponding salts by filtration so titration is used instead
128
Q

Titration method to prepare salts:

A

e.g. making sodium chloride

  1. Choose suitable starting chemicals - one must be an alkali (either sodium hydroxide, potassium hydroxide or ammonia). The other must be the appropriate acid. -e .g. to make sodium chloride sodium hydroxide and hydrochloric acid must be used
  2. Use a glass pipette and place exactly 25.0cm3 of the alkali into a conical flask. Add a few drops of indicator. - e.g. 25.0cm3 of sodium hydroxide is measured by pipette into a conical flask and a few drops of phenolphthalein is added, it will turn pink as sodium hydroxide is added
  3. Put the acid in the burette. Add the acid to the alkali in the conical flask a bit at a time. Swirl the conical flask to make sure that the acid and alkali mix and react.
  4. When the alkali is nearly neutralised add the acid a drop at a time (dropwise) until the indicator changes colour. The alkali is now neutralised. Record the vol of acid added from the burette. - e.g. hydrochloric acid is added to the sodium hydroxide until one drop of the acid turns the phenolphthalein colourless
  5. Repeat the whole experiment again using the exact same volume but without the indicator
  6. The solution made contains the salt dissolved in water. Heat the mixture, evaporate away some of the water and make a hot saturated solution
  7. Leave the solution to cool and condense. Filter the crystals off from the solution and leave to dry in a warm place.
129
Q

Why must titration be used?

A
  • both reactants are liquids/soluble so if you have an excess of one you would not be able to easily remove it from the mixture of products - titration allows you to measure the exact amount of volumes that react
  • you can then mix the exact proportions of the two reactants that react
  • the exact amount of acid therefore, has been added to the soluble reactant, meaning that the leftover solution is only salt and water, no acid or alkali, because they have been completely neutralised
130
Q

What method is being used to prepare the copper sulfate crystals?

A
  • excess base and filtration
  • acid + base ——> salt + water
131
Q

Safety for preparation of pure, dry hydrated copper sulfate crystals starting from copper oxide including the use of a water bath:

A
  • H2SO4, particularly when hot, is corrosive
  • compounds of copper are toxic
132
Q

Method for the preparation of pure, dry hydrated copper sulfate crystals starting from copper oxide including the use of a water bath:

A
  1. Measure correct volumes of of dilute H2SO4 into a small beaker, then add correct volume of water
  2. Add excess of copper oxide (insoluble solid) to H2SO4 acid which react to form copper sulfate
  3. Warm the mixture gently without boiling and stir for 2-3mins, until the copper oxide stops dissolving in and reacting with the acid - the acid has now been neutralised
  4. Use filter and filter paper to filter and recover the unreacted copper oxide. Collect the filtrate (copper sulfate solution) in an evaporating dish
  5. Place the filtrate in a hot water bath to evaporate the water until the solution is saturated
  6. Use the ‘glass rod’ test to decide when saturation has occurred
  7. Leave the solution for 3-4 days so that it can cool and crystallise
  8. Remove the crystals with a glass rod and dry them using filter paper
133
Q

Preparation of pure, dry hydrated copper sulfate crystals starting from copper oxide including the use of a water bath: What is copper oxide?

134
Q

Preparation of pure, dry hydrated copper sulfate crystals starting from copper oxide including the use of a water bath: Why would you not use a plastic beaker at the start of the experiment?

A
  • plastic has a low melting point so it would melt
  • it’s insulating
135
Q

Preparation of pure, dry hydrated copper sulfate crystals starting from copper oxide including the use of a water bath: Why was the mixture of copper oxide and dilute acid warmed?

A

To speed up process of copper oxide dissolving and reaction in acid

136
Q

Preparation of pure, dry hydrated copper sulfate crystals starting from copper oxide including the use of a water bath: Why was excess copper oxide used?

A
  • to ensure all the acid had been reacted so that the solution is pure
  • to ensure that the neutralisation reaction had occurred
137
Q

Preparation of pure, dry hydrated copper sulfate crystals starting from copper oxide including the use of a water bath: What colour was the filtrate and why?

A

blue - contains copper which is transition metal

138
Q

Preparation of pure, dry hydrated copper sulfate crystals starting from copper oxide including the use of a water bath: A student doing the experiment got a filtrate with small black grains in it, what were these?

A
  • copper oxide which was not dissolved - ripped filter paper by mistake
139
Q

Preparation of pure, dry hydrated copper sulfate crystals starting from copper oxide including the use of a water bath: Describe the test with the glass rod:

A
  • it tells us when the solution is saturated
  • dip glass rod into mixture, take out and blow on it —> if crystals appear it’s saturated
140
Q

Preparation of pure, dry hydrated copper sulfate crystals starting from copper oxide including the use of a water bath: What happens when the saturated solution was left for 3-4 days?

A

solid blue CuSO crystals appeared

141
Q

Preparation of pure, dry hydrated copper sulfate crystals starting from copper oxide including the use of a water bath: The copper sulphate crystals are hydrated. What does this word mean and what effect does it have on the colour of copper sulphate?

A

contains water of crystallisation - makes the crystals blues

142
Q

Method to carry out acid-alkali titration using burette, pipette and suitable indicator to prepare dry salt:

A
  1. Wash burette using acid then water
  2. Fill burette to 100cm3 with acid with the meniscus’ base on the 100cm3 line
  3. Use 25cm3 glass pipette to add 25cm3 of alkali into conical flask, drawing alkali into pipette using pipette filler
  4. Add few drops of suitable indicator into conical flask
  5. Add acid from burette to alkali until end-point is reached
  6. The titre (volume of alkali needed to exactly neutralise the acid) is the difference between the first (100cm3) and second readings on the burette)
  7. Repeat experiment to gain more precise results
  8. To prepare pure, dry salt - warm solution to evaporate water
  9. Crystals form
143
Q

What happens in precipitation reactions?

A
  • Precipitation reactions are ‘swapping partner’ reactions.
  • In these reactions an insoluble solid is formed which is called a precipitate. e.g. mixing solutions of lead (II) nitrate and potassium iodide produces lead (II) iodide and potassium nitrate
  • Lead (II) iodide is insoluble and so it precipitates out of the mixture as a yellow solid
  • To work out the identity of a precipitate you need to know the solubility rules
144
Q

What is usually used in precipitation reactions?

A
  • State symbols are usually used in precipitation reactions
  • E.g,. sodium chloride(aq) + silver nitrate(aq) —> sodium nitrate(aq) + silver chloride(s)
145
Q

If you had a solution of 148g of sodium nitrate dissolved in 100g of water at 80 degrees C, what do you think might happen if the solution cooled down to room temperature at 20 degrees C?

A
  • wouldn’t stay dissolved
  • crystals formed
  • 62g would come out of solution
146
Q

Explain why crystals form when a saturated solution is cooled down:

A
  • as the solution cools down the solvent is less able to hold the solute in solution
147
Q

Explain why potassium chloride crystallises out of the solution as the temperature falls:

A
  • as the temp falls the solute becomes less soluble
  • therefore it cannot remain all dissolved so some comes out of the solution
148
Q

What is the precipitation method used for?

A
  • used to make insoluble salt from 2 solutions (soluble → insoluble)
  • excess base method won’t work as both base and salt cannot be easily separated
149
Q

Method for precipitation reaction:

A
  1. Mix the 2 solutions (both must be soluble and one solution must contain the cation and the other must contain the anion), they will react and form an insoluble salt which is the precipitate.
  2. Filter the mixture using filter paper. The residue is the precipitate (insoluble salt)
  3. The insoluble salt will contain impurities from the solution. To remove these impurities, wash the solid with distilled water.
  4. Leave the damp solid to dry in a warm place e.g. a drying oven to evaporate the water
150
Q

Colour of group 1/2 metals:

151
Q

Colour of silver compounds:

152
Q

Colour of copper compounds:

153
Q

Colour of lead compounds:

154
Q

Summary of ways of preparing salts:

155
Q

For the excess base method why is excess of the reactant added?

A

to ensure your volume of acid reacts completely and solution is pure

156
Q

For the excess base method why is the excess reactant removed?

A

this is done by filtration of the insoluble reactant and is done so that you are left with just a salt and water

157
Q

For the excess base method why is the remaining solution only salt and water?

A
  • this is because all of your acid has fully reacted and you have filtered off your other reactant, and that the only products of your reaction are a salt and water
  • if you have used a carbonate you would still only have a salt and water remaining as carbon dioxide gas would have been given off into the atmosphere
158
Q

Why should universal indicator not be used in titration experiments?

GCSE Chemistry 2 (29 decks)

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