Practical skills - Chemistry.

Question 1

Give examples of iritants and the precautions that should be taken.

Answer 1

Dilute acids and alkalis are iritants. Should wear goggles.

Question 2

Give examples of corrosive substances and the precautions that shoulf be taken.

Answer 2

Stronger acids and alkalis are corrosive. Should wear goggles.

Question 3

What safety measures should be taken for flamable substances.

Answer 3

Keep away from naked flames.

Question 4

Give the precautions that shoulf be taken for toxic substances.

Answer 4

Wear gloves and keep away from skin contact - wash hands after use.

Question 5

For oxidising substances what are the precautions that shoulf be taken.

Answer 5

Keep away from flamable and easily oxidised materials.

Question 6

What two things can heating in a crucibe be used for.

Answer 6

Measuring mass loss in thermal decomposition reactions, or mass gain when reacting magnesium with oxygen.

Question 7

Give the method for removing the water of crystalisation as water vapour from calcium sulphate crystals to determin an experimental value for X.

Answer 7

Weigh an empty and dry crucible and lid.
Add 2g of hydrated calcium sulphate to the crucible and weigh again.
Heat strongly with a bunsen.
Allow to cool.
weigh the crucible and contents again.
Heat the crucible again and reweigh untill you reach a constant mass (this will ensure the reaction is complete)

Question 8

Why should large amounts as hydrated calcium sulphate(e.g 50g) not be used when heating it in a cricible.

Answer 8

Since the thermal decomposition is likely to be incomplete.

Question 9

Why does the crucible need to be dry for thermal decomposition reactions.

Answer 9

A wet crucible will give innacurate results since it would cause the mass loss to be too high since the water would evaporate on heating.

Question 10

Why should small amounts of a hyrdayed solid crystal not be used when heating it in a crucible for thermal decomposition. (e.g 0.1 grams)

Answer 10

The percentage uncertainties in weighing the solid will be too high.

Question 11

The the actual value for X in a thermal decomposition reaction shoud be higher than the experimental value what is likely to have happened.

Answer 11

see test.

Question 12

Give two possible things a gas syringe could be used for.

Answer 12

Variety of experiments where the volume of gas is measured.

Probably to work out moles of gas or to follow reaction rates.

Question 13

What should be noted down and why when using a gas syringe.

Answer 13

The Pressure and temperature of the room since the volume of gas wil vary with thease two factors despite the moles staying constant.

Question 14

Give the equation and the units for each variable in the equation related to a gas syringe.

Answer 14

PV=nRT n=8.31 constant , V=M cubed , pressure in pascals Pa.

Question 15

What are potential errors that coud occur when using a gas syringe.

Answer 15

The gas could escpae before the bung is inserted
The syringe can stick
Some gassses like carbon dioxide and sulphur dioxide are not soluble in water and so the true amount of gas is not measured.

Question 16

Giv the method that you would follow to calculate the Mr of propanol.

Answer 16

Extract 0.2cm cubed of propanone into a hyperdermic syringe and then measure the mass of the syringe.
Remove a gas syringe from an oven and note the volume of air already in the barrel.
3. Inject the propanone through the self-seal cap into the barrel. The plunger will move straight away.
4. Put the gas syringe back into the oven.
5. Measure the mass of the empty hypodermic syringe immediately. This change will give you a mass of the volatile liquid in its liquid form as you can’t get this from its gas state.
6. After a few minutes measure the volume of the gas in the gas syringe it will have decreased as the 0.2cm3 of liquid will evaporate due to the heat of the gas syringe as it just came out of the oven, record the temperature of the oven.
shelf and the pressure of the room.
This gives you all the info you need to use ideal gas equation.

Question 17

Giv the method that you would follow to calculate the Mr of propanol.

Answer 17

Extract 0.2cm cubed of propanone into a hyperdermic syringe and then measure the mass of the syringe.
Remove a gas syringe from an oven and note the volume of air already in the barrel.
Inject the prpanone into the self seal cap into the barrel of the gas syringe. The plunger will move streight away.
Measure the mass of the empty hyperdermic syringe.
After a few mins measure the volume of gas in the syringe and note the temp of the oven and the pressure of the room.

Question 18

Draw a gas syringe and give the most important parts of the drawing.

Answer 18

There should be measurment marks of the barrel of the gas syringe and no gaps where gas could escape.

Question 19

Give the method for making up a volumetric solution for a solid.

Answer 19

Weigh the smaple bottle containing the required mass of solid on a 2dp balance.
Transfer this to a beaker.
Reweigh the sample bottle and record the difference in mass.
Add 100cm cubed of distilled water to the beaker.. Use a glass rod to stir to help dissolve the solid.
Sometimes the substance may not dissolve well in cold water so the beaker and its contents could be heated gently until all the solid had dissolved•
Pour solution into a 250cm3 graduated flask via a funnel. Rinse beaker and funnel and add washings from the beaker and glass rod to the volumetric flask.
Shake the volumetric flask thoroughly to ensure a uniform concentration.
Make up to the mark with distilled water using a dropping pipette for last few drops.
• Invert flask several times to ensure uniform solution.
Remember to fill so the bottom of the meniscus sits on the line on the neck of the flask. With dark liquids like potassium manganate it can be difficult to see the meniscus.
Shake the volumetric flask thoroughly to ensure a uniform concentration.

Question 20

Give the method for making up a volumetric solution.

Answer 20

Weigh the smaple bottle containing the required mass of solid on a 2dp balance.
Transfer this to a beaker.
Reweigh the sample bottle and record the difference in mass.
Add 100cm cubed of distilled water to the beaker stiring with a glass rod to help the solid dissolve.
Now transfer the solution into a 250cm cubed graduated flask via a funnel.
Rinse both the beaker and funnel and add the washings from the beaker funnel and glass rod.
Using a dropping pipete make up to the mark with distilled water.
Invert flask several times to ensure a uniform solution.

Question 21

Where should the meniscus sit.

Answer 21

Bottom of the miniscus should sit on the line of the neck of the flask.

Question 22

What can be done if a substance is not disolving well in cold water.

Answer 22

Heat gently untill it all disolves.

Question 23

What is a graduated flask. What is it also known as and why should it not be heated or have hot substances put in it.

Answer 23

A graduated flask has one mark on the neck which is the level to fill to get the acurate volume. It is also known as a volumetric flask. Heat would cause the flask to expad and the volume measured would be incorrect.

Question 24

What problems can occur with measuring out dark luiqids.

Answer 24

With dark luiqids like potasium manganate it can be difficult to see the miniscus.

Question 25

How can mass be measured out acurately in generlal.

Answer 25

In many experiments the best method is the weigh and re-weigh method in which:
On a 2dp balance measure the mass of a weighing bottle with the required mass inside it.
Empty the mass into the reaction vessel or flask.
Re-weigh the now empty weighing bottle
Subtract the mass of the empty weighing bottle fro the first reading to give the exact mass actually added.
If added streight from the bottle washings should be added too.

Question 26

Give the general method for diluting a solution.

Answer 26

Pippete 25 centimeters cubed of the original solution into a 250 cm cubed volumetric flask.
Make up to the mark with distiled water using a dripping pipete for the last few drops.
Invert the flask several times for a uniform solution.

Question 27

Why is a volumetric pippete more acurate than using a measuring cylinder.

Answer 27

It has a smaller uncertainty.The volumetric pipette has one neck Line to fill to whereas the measuring cylinder has loads.

Question 28

What should be used to make a volumetric flask up to the mark.

Answer 28

A teat pippete should be used to make up to the mark so it doesnt go over the line.

Question 29

State what titrations are often used for and the two possible reactions you would see during a titration.

Answer 29

Titrations are done often to find out the concentration of one substance by reacting it with another substance of known concentration. They are often done with neutralisation reactions, but can bedone with redox reactions.

Question 30

Explain generally where the two reactants are placed in a titration.

Answer 30

One substance (generally the one we don’t know the concentration) is put in the conical flask. It is measured using a volumetric pipette.The other substance is placed in the burette

Question 31

Explain what the standard phrase titrate solution A with solution B means.

Answer 31

It means that A should be in the conical flask and B should be in the burette.

Question 32

Why , in a titration is a conical flask used in preferene to a beaker.

Answer 32

A conical flask is used in preference to a beaker because it is easier to swirl the mixture in a conical flask without spilling the contents.

Question 33

Describe in detail how you would use in pipette in the first stage of a titration considering what should and shouldnt be done.

Answer 33

•rinse pipette with substance to go in it (often alkali).•pipette 25 cm3 of solution A into conical flask. The volumetric pipette will have a mark on its neck to show the level to fill to. The bottom of the meniscus should sit on this line.•touch surface of solution with pipette( to ensurecorrect amount is added). A small amount of solution will be left in the pipette at this stage. The calibrationof the pipette will take into account this effect. It should not be forced out.

Question 34

Describe in detail how you would use a beaurete in the second stage of a titration considering things that should and shouldnt be done.

Answer 34

The burette shouldbe rinsed out with substance that will be put in it.If it is not rinsed out the acid or alkali added may be diluted by residual water in the burette or may react with substances left from a previous titration. This would lead tothe concentration of the substance being lowered and a larger titre being delivered.Don’t leave the funnel in theburette because small drops of liquid may fall from the funnel during the titration leading to a false burette reading (would give a lower titre volume) make sure the jet space in the burette is filled with the solution and air bubbles are removed or it will fill during the titration giving a too great a titre.

Question 35

What will happen to the titre reading if the jet space in the beurete is not filled.

Answer 35

If the jet space in the burette is not filled properly prior to commencing the titration it will lead to errors if it then fills during the titration, leading to a larger than expected titre reading.

Question 36

Describe , if given a picture how you would read a beurete and state the accuracy that it should be recorded to in regards to where the miniscus sits.

Answer 36

You should read going down the beurette with the bottom of the miniscus being the value coming downwards that you take.
Even though a burette has marking reading to 0.1cm3, the burette readings should always be given to 2dp either ending in 0.00 or 0.05. 0.05cm3 is the volume of 1 drop of solution delivered from a burette and so this is the smallest difference in readings that can be measured. If the bottom of the meniscus sits on a line it should end with a 0.00 as in the above example 9.00cm3. If the meniscus sits between two lines it should end 0.05. e.g. if the bottom of the meniscus sits between the lines marked 9.1 and 9.2, you should record 9.15

Question 37

Why should only a few drops of indicator be added during a titration.

Answer 37

Indicators are generally weak acids so only add a few drops of them. If too much is added they will affect the titration result.

Question 38

State the two possible ways burette readings should be given and why this is the case dispite a beurette having marking readings of 0.1cm3 , so why do you give your readings to this high degree of acuracy of 0.1cm cubed.

Answer 38

Even though a burette has marking reading to 0.1cm3,the burette readings should always be given to 2dp either ending in 0.00 or 0.05. This is as 0.05cm3 is thevolume of 1 drop of solution delivered from a burette and so this is the smallest difference in readings that can be measured. If the bottom of the meniscus sits on a line it should end with a 0.00 as in the above example 9.00cm3. If the meniscus sits between two lines it should end 0.05. e.g. if the bottom of the meniscussits between the lines marked 9.1 and 9.2, you shouldrecord 9.15

Question 39

State the colour change you would observe with phenolphthalein and the type of titration that it would be suited for.

Answer 39

If acid is added from the burette the colour change would be pink (alkali) to colourless (acid): end point pink colour just disappears [use with titrations using strong alkalis e.g.NaOH

Question 40

Describe the colour change you would see with methyl orange and the type of titration that is would be suitable for.

Answer 40

Methyl orange is a suitable indicator for neutralisation reactions where strong acids are used.It is red in acid and yellow in alkali. It is orange at the endpoint.

Question 41

What should be done during a titrration after the indicator has been added.

Answer 41

Add solution from burette whilst swirling the mixture and add drop-wise at endpoint.

Question 42

Describe a full and general method for a titration.

Answer 42

General Method•
rinse equipment(burette with acid, pipette with alkali, conical flask with distilled water)•
pipette 25 cm3of alkali into conical flask•t
ouch surface of alkali with pipette( to ensure correct amount is added)•
adds acid solution from burette•
make sure the jet spacein the burette is filled with acid•
add a few drops of indicator and refer to colour change at end point•
phenolphthalein[pink (alkali)tocolourless (acid):end point pink colour just disappears][ use if NaOHis used]•methyl orange[yellow (alkali)tored (acid):end point orange] [use ifHClis used]
•use a white tile underneath the flask to help observe the colour change
•add acid to alkali whilsts wirling the mixture and add acid drop wise at end point•
note burette reading before and after addition of acid•
repeats titration until at least 2 concordant results are obtained-two readings within 0.1 of each other.

Question 43

Explain why distiled water can be added to the conical flask during a titration.

Answer 43

Distilled water can be added to the conical flask during a titration to wash the sides of the flask so that all the acid on the side is washed into the reaction mixture to react with the alkali.
It does not affect the titration reading as water does not react with the reagents or change the number of moles of acid added.

Question 44

Why should only distilled water be used to wash out the conical flask between titrations.

Answer 44

Only distilled water should be used towash out conical flasks between titrations because it does not add any extra moles of reagents.

Question 45

Explain how one would be able to get results from a practical titration ( how should titration results be recorded).

Answer 45

Note burette reading before and after addition of solution repeat titration until at least 2 concordant results are obtained-two readings within 0.1 of each other.
Main points are to: record ClEARLY IN A TABLE TOP AND BOTTOM IN FULL - SEE PICTURE.
Include both the initial and final reading of beurete reading.
Under the initial and finlal beurete reading should have a titre columb.
RECORD ABSOLTUTELY EVERYTHING TO 2DECIMAL PLACES. INCLUDE .00 even if its a whole number !!!
You lost a mark for that in E.O.Y
So include results for

Question 46

Why is is neccasary to do multiple titrations.

Answer 46

A single titration could be flawed. Repeating allows for anomalous titres to be spotted and discounted.

Question 47

What are concordant results in a titration.

Answer 47

Two readings within 0.1 of each other.

Question 48

What can we say if 2 or 3 values are within 0.10cm3 and therefore concordant or close in regards to a titration.

Answer 48

We can say results are accurate and repeatable and the titration technique is good and consistent.

Question 49

What should be used to calculate a mean titre.

Answer 49

Only make an average of the concordant titre results.

Question 50

Give some safety precautions for an acid base titration.

Answer 50

Acids and alkalis are corrosive (at low concentrations acids are irritants) Wear eye protection and gloves If spilled immediately wash affected parts after spillage If substance is unknown treat it as potentially toxic and wear gloves.

Question 51

What needs to be done if testing batches during a titration.

Answer 51

In quality control it will be necessary to do titrations/testingon several samples as the amount/concentration of thechemical being tested may vary betweensamples.

Question 52

What needs to be done when titrating mixtures.

Answer 52

If titrating a mixture to work out the concentration of an active ingredient it is necessary to consider if the mixture contains other substances that have acidbase properties. If they don’t have acid base properties we can titrate with confidence.

Question 53

Define a reading. Then define a measurment. Therefore how does a measurment relate to a reading.

Answer 53

Reading- The values found from a single judgement when using a piece of equipment.
Measurment- the values taken as the difference between judgements (readings) of two values ( e.g. using a burette in titration and making a reading of the initial and final readi of the beurete).
The uncertainty of a measurment acounts for two judgments which are the two readings.

Question 54

If uncertainty is not given for a piece of aparatus then what can you asume it to be.

Answer 54

In general, if uncertainty is not indicatedon apparatus, the following assumptionsare made:
For an analogue scale-The uncertainty of a reading(one judgement) is at least±0.5 of the smallest scale reading .The uncertainty of a measurement (two judgements) is at least±1 of the smallest scale reading.-If the apparatus has a digital scale, theuncertainty is +or - the resolution of the apparatus in each measurement.

Question 55

Why would the uncertainty of a beurete be + or - 0.15 cm3.

Answer 55

If The burette used in the titration had an uncertainty for each reading of+/–0.05 cm3 then during a titration two readings would be taken so the uncertainty on the titre volume would be+/–0.10cm3. Then often another 0.05 is added on because of uncertainty identifying the end point colour change.

Question 56

What is the formula that should be used to calculate the percentage uncertainty in each piece of equiptment.

Answer 56

% uncertainty = +or- uncertainty x 100

Measurementmade on apparatus

Question 57

How can you reduce the uncertainty in masuring mass.

Answer 57

Using a balance that measures to more decimal places or usinga larger mass will reduce the % uncertainty in weighing a solid. Weighing sample before and after addition and then calculating difference will ensure a more accurate measurement of themass added (weigh and re-weigh method).

Question 58

If looking at a series of measurments in an investigation which experiments will have the greatest experimental uncertainty.

Answer 58

If looking at a series of measurements in an investigation, the experiments with the smallest readings will have thehighest experimental uncertainties.

Question 59

How would you calculate the total percentage aparatus uncertainty.

Answer 59

To calculate the maximum total percentage apparatus uncertainty in thefinal result add all the individual equipment uncertainties together.

Question 60

How could the percentage uncertainty of a beurette be reduced , and hence the overall titration.

Answer 60

To reduce the % uncertainty in a burette reading it is necessary to make the titre a larger volume. This could be done by: increasing the volume and concentration of the substance in the conical flask or by decreasing the concentration of the substance in the burette.

Question 61

How could the percentage uncertainty of a beurette be reduced during a titration and how is percentage aparatus uncertainty lowered in general for a titration.

Answer 61

To reduce the % uncertainty in a burette reading it is necessary to make the titre a larger volume. This could be done by: increasing the volume and concentration of the substance in the conical flask or by decreasing the concentration of the substance in the burette.
Replacing measuring cylinders with pipettes or burettes which have lower apparatus uncertainty will lower the % uncertainty for the titration.

Question 62

Calculating the percentage difference If we calculated an Mr of 203 and the real value is 214. And then state what it would mean if the percentage uncertainty was greater or less than this value.

Answer 62

By simple actual change over original method you get:
5.41%
Case 1 : If the % uncertainty due to the apparatus < percentage difference between the actual value and the calculated value then there is a discrepancy in the result due to other errors.
Case 2: If the % uncertainty due to the apparatus > percentage difference between the actual value and the calculated value then there is no discrepancy and all errors in the results can be explained by the sensitivity of the equipment.

Question 63

Give the equation for measuring enthalpy of solution and most importantly state the units of each term. State what you need to understand about the value of energy that this equation will give and its limitatios and the common conversions to look out for with the sneeky data they quote in exams.

Answer 63

energy change = mass of solution x heat capacity x temperature change.
Q (J)= m (g)x c(J g-1K-1) x DeltaT ( K)
This equation will only give the energy for the actual quantities used. Normally this value is converted into the ENERGY CHANGE PER MOLE OF LIMITING REAGENT (The enthalpy change of reaction, Delta H).
Energy can only go in the delta H equation where it is in Kj however in this one it is in J.
Aslo this eqution will give temperature CHANGE so do not add 273 like with the ideal gas equation since you will get some very funky nembers , temperature change is the same for both scales numerically.

Question 64

Describe the two different types of basic calorimetric experimental methods that you could be presented with and the really key factor that will be the same in both calculations. What is needed to be worked out.

Answer 64

This could be a solid dissolving or reacting in a solution or it could be two solutions reacting together.
You will need to calculate the moles of both the reactants to determin which is in exess ( you can skip this calculation if you can visually see the moles are the same so just use the data of either reactant). Must take the balanced equation into consideration for exess moles however. Hence you need the moles of the limiting reagent.
For every case ONLY CONSIDER MASS IN TERMS OF SOLUTION.
If one solution and one solid are used DO NOT INCLUDE THE MASS OF THE SOIID in your calculation.
If two solutions are used USE THE COMBINED VOLUME OF THE SOLUTION AS YOUR MASS.
(note your asuming that the solutions have the density of water, which is 1g cm-3.Eg 25cm3 will weigh 25 g.

Question 65

Describe a general method for the calorimetry practical.

Answer 65

General method washest he equipment (cup and pipettes etc) with the solutions to be used
dry the cup after washing
-put polystyrene cup in aglass beaker for insulation and support
-Measure out desired volumes of solutions with volumetric pipettes and transfer toinsulated cup
clamp thermometer into place making sure the thermometer bulb is immersed insolution
-measure the initial temperatures of the solution or both solutions if 2 areused. Do thisevery minute for 2-3 minutes
-At minute 3 transfer second reagent to cup. If a solid reagent is used then add thesolution to the cup first and then add the solid weighed out on a balance.
If using a solid reagent then use‘before and after’ weighing method
stirs mixture( ensures that all of the solution is at the same temperature)
Record temperature every minute after addition for several minutes.

Question 66

Explain what can be done to counteract the cooling effect of a reaction during the calorimetric method, and also what type of reaction would mean exact temperature rise is difficult to gauge. Give an explanation of all the steps that should be taken to do this from start to finish.

Answer 66

If the reaction is SLOW then the exact temperature rise can be difficult to obtain as cooling occurs simultaneously with the reaction. To counteract this we take readings at regular time intervals and extrapolate the temperature curve/line back to the time the reactants were added together.We also take the temperature of the reactants for a few minutes before they are added together to get a better average temperature. If the two reactants are solutions then the temperature of both solutions need to be measured before addition and an average temperature is used.

Question 67

Draw a cooling curve where a second reagent was transfered to a polystyrene cup afeter 3 mins and a temperature rise of 20 Celcius was seen from an initial temp of 18 Celcius.

Answer 67

SEE PHOTO.

Question 68

State the four errors that could arise in a method to determin experimentaly the enthalpy change for a reaction, refer to the asumptions that are made in this method in the first place in your answer.

Answer 68

The four errors here are very different from the errors in falme colorimetry.
• energy transfer fromsurroundings(usually loss), could be gain.
• approximation in specific heat capacity of solution. The method ASSUMES ALL SOLUTIONS HAVE THE HEAT CAPACITY OF WATER. Where The heat capacity of water is 4.18J g-1K-1. In any reaction where the reactants are dissolved in water we assume that the heat capacity is the same as pure water.
• neglecting the specific heat capacity of the calorimeter-we ignore any energy absorbed by the apparatus.
•reaction or dissolving maybe incomplete or slow.•
DENSITY OF THE SOLUTION IS ASUMED TO BE THE SAME AS WATER , where we also assume that the solutions have the density of water, which is 1g cm-3.Eg 25cm 3will weigh 25 g.

Question 69

Example 6. Calculate the enthalpy change of reaction for the reaction where 25.0cm3 of 0.20moldm-3 copper sulfate was reacted with 0.01mol (excess of zinc). The temperature increased 7.0oC.

Answer 69

–146 kJ mol-1

Question 70

What is the difference between delta H and Q .

Answer 70

Q will only give the (negative) energy for the actual quantities used. Normally this value is converted into the energy change per mole of one of the reactants. (The enthalpy changeof reaction,Delta H)

Question 71

How should you deal with signs in a calorimetry equation. And also units.

Answer 71

Probably easiest to leave it out until the end and then state because the temp of solution has increased then the the recation must be exothermic ( decrease in
internal chemical energy) so NEGATIVE SIGN stuck on.
Delta H must always be given as KJ per mol and Q will come out of the equation in J.

Question 72

Example7: 25.0cm 3of 2.00mol dm-3HCl was neutralised by 25.0cm3 of 2.00mol dm-3 NaOH. The temperature increased by 13.5oC. Calculate the enthalpy change per mole ofHCl ?

Answer 72

-56.4kJ mol-1

Question 73

Give an example of a reaction where Hess’s law could be applied where the enthalpy change of the reaction cannot be measured directly by experiments , so alternative reactions are carried out that can be measured experimentally.
Also state why in the example you have given it cant be worked out experimentally.

Answer 73

Hess’s law is used to work out the enthalpy change to form a hydrated salt froman anhydrous salt.

This cannot be done experimentally because itis impossible to add the exact amount of water without the solid dissolving and it is not easy to measure the temperature change of a solid.

Question 74

Describe a detailed method as to how you could calculate the enthalpy change of solution of anhydrous copper(II) sulfate.

Answer 74

1.Weigh out between 3.90 g and 4.10 g of anhydrous copper(II) sulfate in a dry weighing bottle. The precise mass should be recorded.
2.Using a volumetric pipette, place 25 cm3 of deionised water into a polystyrene cup and record its temperature at the beginning (t=0), start the timer and then record the temperature again every minute, stirring the liquid continuously.3
.At the fourth minute, add the powdered anhydrous copper(II) sulfate rapidly to the water in the polystyrene cup and continue to stir, but do not record the temperature.4.
Reweigh the empty weighing bottle5
.At the fifth minute and for every minute up to 15 minutes, stir and record the temperature of the solution in thepolystyrene cup.6.Plot a graph of temperature (on the y-axis) against time. Draw two separate best fit lines; one, which joins the points before the addition, and one, which joins the points after the addition, extrapolating both lines to the fourth minute. STANDARD COOLING CURVE METHOD !
7.Use your graph to determine the temperature change at the fourth minute, which theoretically should have occurred immediately on addition of the solid.
8.Using q= m x cx DeltaT calculate energy change= 20 x 4.18xDelta T
9.Calculate DeltaH solution by dividing q by number of moles of anhydrous copper(II) sulfate in mass added.
10) The above method is then repeated using hydrated copper sulfate. The two DeltaH solution can then be used to calculate the Delta H for the enthalpy change of forming a hydrated salt by using the balanced equation and a Hess’s cycle.

Question 75

What is the method called to meausre the enthalpy change of a reaction experimentaly and what is the method called for measuring the enthalpy of combustion.

Answer 75

Measuring theEnthalpy Changefor a Reaction Experimentally = Calorimetric method.
Measuring Enthalpies of Combustion = Flame Calorimetry.

Question 76

Outline a simple method for measuring the enthalpy of combustion theoretically.
State the 3 measurments you must take during this process.

Answer 76

Enthalpies of combustion can be calculated by using calorimetry. Generally the fuel is burnt and the flame is used to heat up water in a metal cup.
Need to measure:
•mass of spirit burner before and after•
Temperature change of water•
Volume of water in cup (taken as the mass)

Question 77

Give the 4 possible sources of error in the method for flame calorimetry.

Answer 77

1) energy losses from calorimeter
2) •Incomplete combustion of fuel•(Incomplete transfer of energy)
3) •Evaporation of fuel after weighing
4) •Heat capacity of calorimeter not included ,it will absorb heat itself.

Weird one-Measurements not carried out under standard conditions as H2O is gas, not liquid, in this experiment.

Question 78

Example8. Calculate the enthalpy change of combustion for the reaction where 0.650g of propan-1-ol was completely combusted and used to heat up 150g of water from 20.1 to 45.5oC.

Answer 78

–1470 kJ mol-1

Question 79

Give a detailed method as to how you would test for group 2 metals and the results you would get for each metal. Give any equations that support your answer.
State also the principle you can apply to get your answer without really having to memorise anything.

Answer 79

Method: adding DILUTE SODIUM HYDROXIDE
a) Place about 10 drops of 0.1 mol dm–3 metal ion solution in a test tube.
b) Add about 10 drops of 0.6 mol dm–3 sodium hydroxide solution, mixing well.
c) Continue to add sodium hydroxide solution, dropwise with gentle shaking, until in excess
Observations:
TWO MOST INSOLUBLE HYDROXIDES ARE THEN ONLY ONES THAT ACTUALLY DO ANYTHING.
Magnesium hydroxide is classed as insoluble in water and
will appear as a white precipitate.
Simplest Ionic Equation for formation of Mg(OH)2 (s) Is:
Mg2+ (aq) + 2OH-(aq) —> Mg(OH)2 (s).
Calcium hydroxide is classed as SPARINGLY SOLUBLE in water and will appear as a white precipitate (it may need more sodium hydroxide to be added before it appears compared to a magnesium solution.)
Simplest Ionic Equation for formation of Ca(OH)2 (s) Ca2+ (aq) + 2OH-(aq)—>Ca(OH)2 (s).
The other two won’t form anything as are soluble.
Strontium and barium salts will not form a hydroxide precipitate on addition of sodium hydroxide due to their high solubility. The solutions will be highly alkaline

Question 80

What two things can adding a little dilute sodium hydroxide test for .

Answer 80

This test can be used on group 2 metal ions and transition metal ions.
(Transition material is year 2)

Question 81

What is the test for the ammonium ion NH4+ and explain why it does this.

Answer 81

Testing for Ammonium ions (NH4+)
a) Place about 10 drops of 0.1 mol dm–3 ammonium chloride in a test tube.
b) Add about 10 drops of 0.4 mol dm–3 sodium hydroxide solution. Shake the mixture. c) Warm the mixture in the test tube gently using a water bath.
d) Test the fumes released from the mixture by holding a piece of damp red litmus
paper in the mouth of the test tube.

Essentially add sodium hydroxide and warm the mixture
MUST INCLUDE WARM.
Explanation: the ammonia gas that is produced (not ammonium gas) will be alkaline.
Since it is alkaline it will will turn red litmus paper blue.
RED INTO BLUE.

Remember alkalines turn red litmus paper blue and this is true for dipping it in alkaline solutions also.

Question 82

Give the similarities between the two tests for group cations.

Answer 82

Weather using sulphate or hydroxide method. Both group 2 solutions used are of 0.1M concentration. Note the concentration of sulphuric acid used should be 1M and NaOH should be 0.6M.
Both have a 50:50 split of the most insoluble 2 will form a white precipitate and the least insoluble two will not.

Question 83

Give the alternative test for group 2 cations other than the addition of dilute NaOH.
What simple principle can you use.,

Answer 83

Add 1M of sulphuric acid or a soluble sulphate such as Na2SO4 to a solution of 0.1M of the metal and get the reverse two precipitates as you would get for NaOH.

Question 84

Give a detailed method for testing for group 2 cations by using the sulphate ion by way of adding sulphuric acid or soluble sulphate such as Na2SO4. Give all ionic equations and state the overall equation for SrCl2 with Na2SO4.

Answer 84

a) Place about 10 drops of 0.1 mol dm–3 metal ion solution in a test tube.
b) Add about 10 drops of 1.0 mol dm–3 sulphuric acid ( or other soluble sulfate solution. c) Continue to add sulphuric acid solution, drop-wise with gentle shaking, until in excess.
As always the least soluble metals will be the only ones that form a white precipitate.
Strontium and barium solutions will form white precipitates with addition of sulfate ions
Full equation :
SrCl2(aq) + Na2SO4 (aq)—>2NaCl (aq) + SrSO4 (s)
Ionic equation: Sr2+ (aq) + SO42-(aq)—>SrSO4 (s).
Ionic equation: Ba2+ (aq) + SO42-(aq)—>BaSO4 (s)
Magnesium and calcium salts will not form a sulfate precipitate on addition of sulfate ions due to their high solubility.

Question 85

4 marks - Give the test for a sulphate ion. Include all relevant equations to support your answer. Explain the need for each step in terms of why certain reagents have been chosen over others. Give all observations.

Answer 85

BaCl2 solution acidified with Nitric acid is used as a reagent to test for sulphate ions.
If acidified barium chloride is added to a solution that contains sulphate ions a white precipitate of barium sulphate forms: Simplest ionic equation
Ba2+ (aq) + SO42-(aq)—>BaSO4 (s).
The Nitric acid is needed to react with carbonate impurities that are often found in salts which would form a white barium carbonate precipitate and so give a false result : Ba2+ + CO3 2- —> BaCO3 2- (Also white ppt)
You could not use sulfuric acid because it contains sulfate ions and so would give a false positive result by providing source of SO42- for above reaction to happen. Again forming unwanted white ppt.
Carbonate removed as CO2 equation:
2HCl + Na2CO3 —> 2NaCl + H2O + CO2
Observation - Fizzing due to CO2 would be observed if a carbonate was present.

Question 86

6 marks- State how you could positively identify the presence of three test tubes that are known to be KCl , KI , KBr and KF.
State the role of the nitric acid and were you know his reaction from.

Answer 86

This is just the standard three step method.
Step 1:
The test solution is made acidic with nitric acid
The role of nitric acid is to react with any carbonates present to prevent formation of the precipitate Ag2CO3.(s) - This would mask the desired observations.
2 HNO3 + Na2CO3—>2 NaNO3 + H2O + CO2, this is exactly the same reaction as for removing the carbonate ions in the test for sulphate ions. Since BaCO3 is also insoluble.

Step 2:Silver nitrate solution is added drop-wise.
Fluorides produce no precipitate - Silver fluoride is highly soluble in water.
Chlorides produce a white precipitate
Ag+(aq) + Cl- (aq) —> AgCl(s)
Bromides produce a cream precipitate
Ag+(aq) + Br- (aq)—>AgBr(s)
Iodides produce a pale yellow precipitate
Ag+(aq) + I- (aq)—> AgI(s)
The precipitates get increasingly more pale down the group.
Treating with ammonia: If colours look similar. Silver chloride dissolves in dilute ammonia to form a complex ion AgCl(s) + 2NH3(aq)—>[Ag(NH3)2]+ (aq) + Cl- (aq)
Colourless solution
Silver bromide dissolves in concentrated ammonia to form a complex ion
AgBr(s) + 2NH3(aq) —>[Ag(NH3)2]+ (aq) + Br - (aq)
Colourless

Silver iodide does not react with ammonia – it is too insoluble.

Question 87

What is a type of reagent that could be used to test for the carbonate ion. Give a general set of compounds and then a specific few examples. Give the two variations of this method and state when you should use each in an exam.

Answer 87

Any dilute acid will produce effervescence by liberating CO2 from carbonate. This is best choice for a “SIMPLE” test tube reaction messing about with bubbling through stuff isn’t “SIMPLE”.
HCl or nitric acid.
Reaction to Liberate CO2 in the first place is:
2 HNO3 + Na2CO3—>2 NaNO3 + H2O + CO2 or
2HCl + Na2CO3 —> 2NaCl + H2O + CO2

Question 88

What is the incredibly easy test for hydroxide ions that bearfield showed you back during that very first practical endorsement.

Answer 88

Alkaline hydroxide ions will turn red litmus paper blue.

Question 89

A colourless solutions contains a mixture of NaCl and NaBr . Give a detailed method of how you could prepare A PURE SAMPLE OF AgBr.

Answer 89

This answer (and most involving testing halides) are best set out in three stages. This gives a nice structure.
Lights should light up when you see Pure sample that some form of filtering will be involved.
Stage-1 : Add acidified silver Nitrate.
This will produce ppt: Full equation or ionic is fine.
AgNO3 + NaCl —> AgCl(s) + NaNO3
(Exactly the same to form ppt of AgBr(s).)

Stage 2: Recognise only one will dissolve to leave other not dissolved.
Add dilute NH3 in excess so the silver chloride will dissolve but the AgBr will not.
AgCl(s) + 2NH3(aq)—>[Ag(NH3)2]+ (aq) + Cl- (aq)
Bromine not react.
Stage 3: Filter off
Filter to remove silver bromide ppt from solution of crap
WASH IT AND DRY IT to produce pure AgBr.

Question 90

Give there properties of NH3+ gas produced from warming NH4 with NaOH.

Answer 90

It is very soluble. It has a pungent chocking smell.
It is alkaline so turns red litmus paper blue.
*things like to start orange or red In chem tests!

Question 91

Describe how a test tube of MgSO4 , KI , KCl , and KBr can be positively separated.
Give the ionic equation for the removal of carbonate ions as carbon dioxide and the removal of hydroxide ions as water.

Answer 91

Add HNO3 to all test tubes to to remove carbonate impurities and also dont forget hydroxide impurities also. Following this add silver nitrate.
Ionic equations for the removal of both impurities:
NO3 2- + CO3 2- + 2H+ —> CO2 +H2O + 2NO2-
NO32 + OH- +H+ —> H20 +NO3-

After adding the silver nitrate the test tube with MgSO4 will not have changed and so it is required that this is put to the side as all others must have a halide present.
Half each of the remaining three test tubes and and dilute ammonia to half and conc to the other.
The KI tube won’t dissolve in either.
The KBr tube will dissolve in conc
The KCl will dissolve with dilute ammonia.

Question 92

Describe how a test tube of MgSO4 , KI , KCl , and KBr can be positively separated.
Give the ionic equation for the removal of carbonate ions as carbon dioxide and the removal of hydroxide ions as water.
Give any safety precautions.

Answer 92

Add HNO3 to all test tubes to to remove carbonate impurities and also dont forget hydroxide impurities also. Following this add silver nitrate.
Ionic equations for the removal of both impurities:
NO3 2- + CO3 2- + 2H+ —> CO2 +H2O + 2NO2-
NO32 + OH- +H+ —> H20 +NO3-

After adding the silver nitrate the test tube with MgSO4 will not have changed and so it is required that this is put to the side as all others must have a halide present.
Half each of the remaining three test tubes and and dilute ammonia to half and conc to the other.
The KI tube won’t dissolve in either.
The KBr tube will dissolve in conc
The KCl will dissolve with dilute ammonia.
Take a fresh test tube of the one that had no change apon the addition of AgNO3 and add BaCl acidified with nitric acid , a white precipitate will form if MgSO4 is present.
BaCl is toxic so wear nitrile gloves.