Praticals Flashcards
What are the steps in the experiment that you have to do to find the formula of a hydrated salt
- Weigh an empty crucible
- Add the hydrated salt into the weighed crucible. Weigh the crucible and the hydrated salt
- Using a pipe-clay triangle, support the crucible containing the hydrated salt on a tripod.
- Heat the crucible and contents gently for about one minute. Then heat it strongly for a further three minutes.
- Leave the crucible to cool. Then weigh the crucible and anhydrous salt.
How do you calculate the formula for a hydrated salt from the experimental results
- Calculate the amount, in mol of anhydrous salt
- Calculate the mass and amount, in mol, of water
- Find the smallest whole-number ratio
- Write down the vale of the x and the formula of the hydrated salt
e. g. CuSO4. 5H2O X=5
What assumptions have been made in the calculations of the formula of a hydrated salt
- All of the water has been lost- Difficult to tell if hydrated and anhydrous colours are similar.
Solution = heat to constant mass- the crystals are reheated repeatedly until the mass of the residue no longer changes - No further decomposition- many salts decompose further when heated very strongly e.g copper (II) sulfate decomposes to form black copper (II) oxide. This can be difficult to judge if no colour change.
What are the steps in the experiment that you have to do when finding a relative molecular mass
- You can use the ideal gas equation to find rmm of a volatile liquid. The unknown liquid would need to be liquid at room temperature but have a boiling point below 100 degrees so it vapourises.
- Add a sample of the volatile liquid to a small syringe via a needle. Weigh the small syringe.
- Inject the sample into a gas syringe through the self-sealing rubber cap. Reweigh the small syringe to find the mass of the volatile liquid added to the gas syringe.
- Place the gas syringe in a boiling water bath at 100 degrees.
- The liquid vapourises producing a gas. The pressure is recorded.
What are the calculations needed to determine the relative molecular mass from experimental results
- Convert all the quantities to match the ideal gas equation. e.g m^3, Pa and K
- Use the ideal gas equation to calculate the unknown
- Find the molar mass by dividing the mass by mols
What is the experiment that is carried out to identify an unknown group 2 metal
- Have a conical flask and a gas syringe attached. The unknown metal should be in the bottom of the conical flask
- Weigh a sample of the metal and add to the flask
- Using a measuring cylinder add dilute HCl to the flask and quickly replace the bung
- Measure the maximum volume of gas in the syringe
What calculations need to be done to identify an unknown group 2 metal from experimental results
- Calculate the amount of H2(g) by dividing the volume by 24000
- From the equation and the result in step 1, the amount of metal X can be determined. Molar ratios.
- Work out the relative atomic mass from the mols and mass.
What are 3 things titrations can be used for
- Finding the concentration of a solution
- Identification of unknown chemicals
- Finding the purity of a substance
How do you prepare a standard solution
- The solid is first weighed accurately
- The solid is dissolved in a beaker using less distilled water than will be needed to fill the volumetric flask
- The solution is transferred to a volumetric flask. The last traces of the solution are rinsed into the flask with distilled water.
- The flask is carefully filled to the graduation line by adding distilled water a drop at a time until the bottom of the meniscus lines up exactly with the mark.
- View at eye level for accuracy
- Finally, invert the volumetric flask slowly several times to mix the solution thoroughly
Describe the process of an acid-base titration
- Add a measured volume of solution to a conical flask using a pipette
- Add the other solution to a burette and record the initial burette reading.
- Add a few drops of indicator to the solution in the conical flask.
- Run the solution in the burette into the solution in the conical flask, swirling the flask throughout to mix the two solutions
- Eventually the indicator changes colour at the end point of the titration. This indicates the volume of solution that reacts exactly with the volume of the second solution
- Record the final burette reading. The volume of solution added is the titre.
- A quick, trial titration is carried out first to find the approximate titre.
- The titration is then repeated accurately, adding the solution dropwise as the endpoint is approached.
- Repeat until two titres are concordant.
Describe the calculations carried out for an acid-base titration
- Work out the amount, in mol, of the solute in the solution for which you know both the concentration and volume
- Use the equation to work out the amount of mols of the solute in the other solution
- Work out the unknown information e.g concentration from the titre volume and mols.
Describe how to carry out a titration to identify an unknown carbonate
- Prepare a solution of an unknown carbonate in a volumetric flask.
- Using a pipette transfer a set volume of your solution into a conical flask.
- Using a burette titrate this solution using a known concentration of HCl
Describe the calculations required to identify an unknown carbonate from experimental results
- Calculate the amount of HCl that reacted
- Determine the amount of X2CO3 that reacted
- Scale up to find the amount of X2CO3 in the standard solution you prepared
- Find the molar mass of X2CO3
- Use the molar mass to identify X
What is the sequence of tests used to analyse an unknown inorganic compound for anions
- Carbonate test- add a dilute acid and look for effervescence. Neither sulfate nor halide ions produce bubbles with dilute acid,
- Sulfate test- add Ba2+ ions and look for white precipitate. Carbonates also produce white precipitate with Ba2+ ions so important to carry out carbonate test first.
- Halide test- add Ag+ ions and look for precipitate. Silver carbonate and silver sulfate both produce precipitates and therefore important to carry out the halide test last.
How do you analyse a mixture of chemicals
- Carry out the same sequence on the same solution.
- Carbonate test- if bubbles seen keep adding dilute nitric acid until bubbling stops as this removes all carbonate ions
- Sulfate test- Add excess Ba(NO3)2 to the same solution. Any sulfate ions will precipitate out as barium sulfate
- Filter the solution to remove the barium sulfate
- Halide test- Add AgNO3 followed by NH3 as normal
How do you determine enthalpy change of combustion
- Using a measuring cylinder, measure out a set volume of water. Pour the water into the beaker. And record the initial temperature of the water.
- Add methanol to the spirit burner. Weigh the spirit burner containing methanol
- Place the spirit burner under the beaker and burn the methanol whilst stirring the water with the thermometer.
- After about 3 minutes extinguish the flame. Immediately record the maximum temperature reached by the water.
- Re-weigh the spirit burner containing the methanol. Assume that the wick has not been burnt.
Describe how to calculate the enthalpy change of combustion from experimental results
- Calculate the energy change q of the water in kJ
- Calculate the amount in mols of CH3OH burnt
- Calculate the Enthalpy change in kJmol-1
What are possible reasons for the difference in calculated enthalpy change and the data book value
- Heat loss to the surroundings other than water e.g air and beaker
- Incomplete combustion of methanol- carbon monoxide and carbon produced instead of CO2. See carbon as a black layer of soot on the beaker
- Evaporation of methanol from the wick.
- Non-standard conditions
Use of draught screens and an input of oxygen gas could minimise errors from heat loss and incomplete combustion
Describe the experiment used to determine the enthalpy change of reaction
- Put the two solutions or a solution and a solid into a polystyrene cup.
- Put a thermometer in the cup and record any temperature change.
- Same method can be used to determine enthalpy change of neutralisation.
Describe how to calculate the enthalpy change of reaction from experimental results
- Calculate the energy change in the solution
- Calculate the amount that reacted (use the thing that is not in excess)
- Calculate the enthalpy change
What are cooling curves and describe how to use them
- A method that can be used to correct for heat loss
- Carry out an experiment to determine enthalpy change of reaction.
- Start a stop-clock and take temperatures of the solution every 30s untils the temperature stays constant
- Plot a graph of temperature against time.
- To correct colling, extrapolate the cooling curve section of the graph back to the time when the reactants were added. Draw a vertical line from the time the solutions were mixed to the extrapolated cooling curve
Describe an experiment for monitoring the production of gas using gas collection
- e.g rate of reaction of decomposition of H2O2
- H2O2 is added to the conical flask and the bung is replaced
- The initial volume of gas in the measuring cylinder is recorded.
- Manganese dioxide- catalyst- is then quickly added to the conical flask and the bung is replaced. The stop clock is started
- THe volume of gas produced in the measuring cylinder is recorded at regular intervals until the reaction is complete
- The reaction is complete when no more gas is produced,
- A graph is plotted of total volume of gas produced against time
- Tangents can be drawn to calculate the rate of reactions. Gradient= rate of reaction
Describe an experiment for monitoring the loss of mass of reactants using a balance
e. g. calcium carbonate and HCl
1. The carbonate and the acid are added to a conical flask on a balance.
2. The mass of the flask and contents is recorded initially and at regular time intervals.
3. The reaction is complete when no more gas is produced so no more mass is then lost.
4. A graph of mass lost against time is plotted
5. Draw tangents and gradient = rate of reaction
Describe the experiment for the hydrolysis of haloalkanes
- Reagents- AgNO3 (aq)
- Conditions- ethanol as a co-solvent and warming
- Observations- Measure time taken for a silver halide precipitate to form
- Equation for hydrolysis- CH3CH2X + H2O –> CH3CH2OH + H+ + X-
- Equation from precipitate reaction- Ag+ + X- –> AgX
- Trend in reactivity- Iodo is fastest chloro is slowest
- Why- The strength of the C-X bond affects the reactivity. C-I is a weaker bond strength than C-Cl (strength depends on size of atom not polarity)
Describe an experiment to test the rate of reaction using a colorimeter
e. g. propanone and iodine
1. As the reaction progresses, iodine is used up and its orange/brown colour fades.
2. Prepare standard solutions of known concentration of the coloured chemical iodine in this reaction.
3. Select a filter with the complementary colour of the coloured chemical e.g. green/blue for iodine
4. Zero the colorimeter with water
5. Measure the absorbance readings of the standard solutions of iodine
6. Plot a calibration curve of absorbance against iodine concentration. This is a way to convert an absorbance reading into a concentration of iodine.
7. Carry out the reaction between propanone and iodine. Take absorbance readings of the reacting mixture at measured time intervals.
8. Use the calibration curve to measure the concentration of iodine at each absorbance reading.
9. Finally plot a second graph of concentration of iodine against time. From a concentration-time graph you can determine the order of reaction with respect to the coloured chemical.
Describe how the iodine clock procedure
- Several experiments are carried out using different concentrations of one of the reactants and all other concentrations are kept constant
- The colour change is delayed by including a small amount of another chemical which actually removes iodine as it forms. As soon as this chemical is used up the blue-black appears
- In each experiment the solution is colourless at the start and the time t is measured for the blue-black colour of the starch-iodine to appear
- The initial rate is proportional to 1/t
- A graph of 1/t against concentration is then plotted.
- From the results the order with respect to each reactant is determined and a rate equation written. The rate constant can then be calculated.
How accurate are clock reactions
- You are measuring the average rate during the first part of the reaction.
- Over this time you can assume that the average rate of reaction is constant and is the same as the initial rate.
- In a clock reaction, you are measuring an average rate of a change in reactant over time.
- The shorter the period of time over which an average rate is measure, the less the rate changes over that time period.
- It is relatively accurate despite being an approximation.
Describe an experiment used to determine Kc
e. g CH3COOH + C2H5OH ↔ CH3COOC2H5 + H2O
1. Mix together the solutions in a conical flask and add HCl as an acid catalyst
2. Add HCl into a second flask to act as a control
3. Stopper both flasks and leave for a week to reach equilibrium
4. Carry out a titration on the equilibrium mixture using a standard solution of NaOH
5. Repeat the titration with the control to determine the amount of acid catalyst that had been added
Describe how to calculate Kc from experimental results
- Determine the equilibrium amount of CH3COOH from the results of the two titrations.
- Use the equilibrium equation to determine the equilibrium amounts of each component
- Find the equilibrium concentrations
- Write the expression for Kc and substitute in values and calculate Kc
Describe an experiment using a pH meter
- A pH meter consists of an electrode that is dipped into a solution and connected to a meter that displays the pH readings.
- Using a pipette, add a measured volume of acid to a conical flask
- Place the electrode of the pH meter in the flask.
- Add the aqueous base to the burette and add to the acid in the concial flask
- After each addition swirl the contents. Record the pH and the total volume of the aqueous base added.
- Repeat steps 4 and 5 until pH starts to change more rapidly. Then add the aqueous base dropwise for each reading until the pH changes less rapidly.
- Then add the base 1cm3 at a time again until an excess has been added and the pH has been basic, with little change for several additions.
- A graph of pH against total volume of aqueous base added is then plotted.
- You could attached the pH meter to a data logger instead and use a magnetic stirrer in the flask.
Describe how to carry out a manganate (VII) titration
- MnO4- ions are reduced and so the other chemical used must be a reducing agent that is oxidised
- A standard solution of potassium manganate (VII) is added to the burette.
- Using a pipette add a measured volume of the solution being analysed to the conical flask. An excess of dilute H2SO4 is also added to provide the H+ ions required for the reduction of MnO4- ions.
- During the titration, the manganate (VII) solution reacts and is decolourised as it is being added. The end-point of the titration is judged by the first permanent pink colour, indicating when there is an excess of MnO4- ions present
- Repeat the titration until you obtain concordant titre
What are two examples of manganate (VII) titrations
- Iron Fe2+
2. Ethanedioic acid (COOH)2
Describe how to analyse the percentage purity of an iron (II) compound
- Prepare a 250cm3 solution of the impure FeSO4.7H2O in a volumetric flask
- Using a pipette measure out 25cm3 of this solution into a conical flask. Then add 10cm3 of 1 mol dm-3 H2SO4 (an excess)
- Using a burette, titrate this solution using a standard solution of potassium manganate
- Finally analyse your results to determine the percentage purity
Describe how to calculate the percentage purity of an iron (II) compound from experimental results
- Calculate amount of MnO4- that reacted
- Determine the amount of Fe2+ that reacted. molar ratios
- Scale up to find the amount of Fe2+ in 250 cm3 e.g multiply by 10
- Find the mass of FeSO4.7H2O in the impure sample
- Find the percentage purity of FeSO4.7H2O in the impure sample.
Percentage purity= mass of FeSO4.7H2O/ mass of impure sample *100
Describe the procedure of an iodine/thiosulfate titration
- Add a standard solution of Na2S2O3 to the burette.
- Prepare a solution of the oxidising agent to be analysed. Using a pipette, add this solution to a conical flask.
- Then add an excess of potassium iodide
- The oxidising agent reacts with iodide ions to produce iodine, which turns the solution a yellow-brown colour
- Titrate this solution with Na2S2O3. During the titration, the iodine is reduced back to I- ions and brown colour fades quite gradually, making it difficult to decide on an end point.
- When the end point is being approached, the iodine colour has faded enough to become a pale straw colour. When this happens a small amount of starch indicator is added
- A deep blue-black colour forms to assist with the identification of the end point.
- As more sodium thiosulfate is added the blue-black colour fades. At the end-point, all the iodine will have just reacted and the blue-black colour disappears.
State 3 oxidising agents that can be determined by an iodine/thiosulfate titration
- ClO- content in household bleach
- Cu2+ content in copper (II) compounds
- Cu content in copper alloys
Describe how to calculate the ClO- concentration in bleach from experimental results
- Calculate the amount of S2O32- that reacted
- Determine the amount of I2 and ClO- that reacted.
- Determine the amount of ClO- in 250cm3
- Determine the concentration of ClO- in bleach
Describe how to measure standard cell potentials
- Prepare two standard half-cells
- for a metal/metal ion half-cell, the metal ion must have a concentration of 1 mol dm-3
- For an ion/ ions half-cell, both metal ions present in the solution must have the same concentration. There must be an inert electrode, usually platinum
- For a half-cell containing gases, the gas must be at 100kPa pressure, in contact with a solution with an ionic concentration of 1 mol dm-3. There must be an inert electrode usually platinum.
- For all half cells the temperature must be 298K - Connect the metal electrodes of the half-cells to a voltmeter using wires.
- Prepare a salt bridge by soaking a strip of filter paper in a saturated aqueous solution of KNO3
- Connect the two solutions of the half-cells with a salt bridge
- Record the standard cell potential from the voltmeter
