common experiments Flashcards
Investigating the Energy in Fuels
what are ?
These are calorimetry experiments that usually test the energy in different alcohols.
The efficiency of fuels can be analysed by comparing the amount of energy they release during combustion.
Investigating the Energy in Fuels
method and drawing
Method:
Measure the mass and initial temperature of the water
Fill the spirit burner with test substance and measure and record its mass
Stir the water constantly with the thermometer and continue heating until the spirit burner burns out
Measure and record the highest temperature of the water
Investigating the Energy in Fuels
calculation
Calculation:
Temperature change of water = final temperature – initial temperature
Number of moles burned = change in mass ÷ molecular mass
Amount of energy = change in temperature x mass of water x specific heat capacity
Amount of energy per mole (J mol-1) = total amount of energy ÷ moles burned
There are several factors that can affect the rate of a reaction. These are:
Surface area of solid reactants
Concentration of the reactants
Temperature at which reaction is carried out
The use of a catalyst
The influence of light on photochemical reactions
Experiment 2a: Surface Area
method and drawing for every passage
Method:
Add dilute hydrochloric acid into a conical flask
Use a capillary tube to connect this flask to a measuring cylinder upside down in a bucket of water (downwards displacement)
Add calcium carbonate chips into the conical flask and close the bung
Measure the volume of gas produced in a fixed time using the measuring cylinder
Repeat with different sizes of calcium carbonate chips (solid, crushed and powdered)
Experiment 2b: Concentration drawing for every passage and method
Method:
Measure 50 cm3 of Sodium Thiosulfate solution into a flask
Measure 5 cm3 of dilute Hydrochloric acid into a measuring cylinder
Draw a cross on a piece of paper and put it underneath the flask
Add the acid into the flask and immediately start the stopwatch
Look down at the cross from above and stop the stopwatch when the cross can no longer be seen
Repeat using different concentrations of Sodium Thiosulfate solution (mix different volumes of sodium thiosulfate solution with water to dilute it)
Experiment 2c:Temperature
Method:
Dilute Hydrochloric acid is heated to a set temperature using a water bath
Add the dilute Hydrochloric acid into a conical flask
Add a strip of Magnesium and start the stopwatch
Stop the time when the Magnesium fully dissolves
Repeat at different temperatures and compare results
Experiment 2d: Catalyst
Method:
Add Hydrogen Peroxide into a conical flask
Use a capillary tube to connect this flask to a measuring cylinder upside down in a bucket of water (downwards displacement)
Add the catalyst Manganese(IV) Oxide into the conical flask and close the bung
Measure the volume of gas produced in a fixed time using the measuring cylinder
Repeat experiment without the catalyst of Manganese(IV) Oxide and compare results
Determining the Formulae of Metal Oxides
Common metals uses in these experiments include
There are two methods to carry out this investigation.
copper, magnesium and calcium.
Combustion of Metal Oxide
Reduction of Metal Oxide
Experiment 3a: Combustion of Metal Oxide
method draw passages calculation
Method:
Measure and record the mass of crucible and lid
Add sample of metal into crucible and measure mass with lid and record
Strongly heat the crucible over a bunsen burner for several minutes
Frequently lift the lid to allow sufficient air into the crucible for the metal to fully oxidise but without letting any of the gaseous metal oxide to escape
Continue heating until the mass of crucible remains constant
Measure the mass of crucible and contents and record
Calculation of empirical formula:
Mass of metal = (mass of crucible + lid + metal) – (mass of crucible + lid)
Mass of metal oxide = (mass of crucible + lid + oxide) – (mass of crucible + lid)
Mass of oxygen = mass of metal oxide – mass of metal
Step 1 – Divide each of the two masses by their Relative Atomic Masses
Step 2 – Simplify the ratio
Metal Oxygen
Mass x y
Mole x / Mr y / Mr
= a = b
Ratio a : b
Step 3 – Write out the formula using the ratio e.g. for magnesium: MgaOb
Experiment 3b: Reduction of Metal Oxide
method draw passages calculation
Method:
Measure and record the mass of the metal oxide
Use a clamp to hold boiling tube horizontally, and place the metal oxide at the end of the tube
Heat using a bunsen burner until all the oxide has completely changed colour, indicating that all oxygen has been reduced
Measure mass and record mass of the remaining powder
Calculation of empirical formula:
Mass of metal = mass of remaining metal powder
Mass of oxygen = (mass of metal oxide) – (mass of metal powder)
Step 1 – Divide each of the two masses by their Relative Atomic Masses
Step 2 – Simplify the ratio
Metal Oxygen
Mass x y
Mole x / Mr y / Mr
= a = b
Ratio a : b
Step 3 – Write out the formula using the ratio e.g. for copper: CuaOb
Investigating Solubility of Salts
The solubility of a solid is the…
The solubility of a substance in water at different temperatures ….and can be plotted on a graph to produce a…
amount of that solid that you can get to dissolve in a fixed amount of water at room temperature.
changes
solubility curve.
Experiment 4: Investigating Solubility of Salts
method, draw passages
Method:
Set water bath to specific temperature
Use water from water bath and add to beaker, making sure to use a thermometer to record the exact temperature of the water in the beaker
Record mass of the salt and add to the beaker
Measure time taken for the salt to dissolve and record
Repeat the experiment for different solids at different temperatures
Experiment 4: Investigating Solubility of Salts
results, graph
Results:
Generally most solids become more soluble as the temperature increases.
This is why sugar dissolves better in hot water than in cold water.
The graph below shows the solubility curves produced from the investigation of the solubility of NaCl, KNO3 and PB2(NO3) at 20ºC, 30ºC and 40ºC.
Solubility curves can be used to compare solubilities of different compounds and to predict the yield produced on crystallisation.
This can be done by extrapolation.
Performing an Acid-Base Titration
Acid base titrations are the most common kind of titration.
You may be asked to calculate the moles present in a given amount, the concentration or volume required to neutralise an acid or base.
Titrations may also be used to
prepare salts or other precipitates and in redox reactions or to prepare complexes.
Experiment 5: Performing an Acid-Base Titration
method draw passages
Method:
Place the conical flask on a white tile so the tip of the burette is inside the flask
Add a few drops of a suitable indicator to the solution in the conical flask
Perform a rough titration by taking the burette reading and running in the solution in 1 – 3 cm3 portions, while swirling the flask vigorously
Quickly close the tap when the end-point is reached (sharp colour change) and record the volume (being sure that you place your eye level with the meniscus)
Now repeat the titration with a fresh batch of alkaline
As the rough end-point volume is approached, add solution from the burette one drop at a time until the indicator changes colour
Record the volume to the correct number of decimal places (to 0.1cm3)
Repeat until you achieve two concordant results (two results that are within 0.1cm3 of each other)
Experiment 5: Performing an Acid-Base Titration
calculations
Make sure you balance the chemical equation before performing the calculations
Use the general formulae below and rearrange it to calculate what the question is asking you to find:
common error when performing titrations is to leave the funnel on the top of the burette after filling.
This is a source of error as
liquid may drip into the burette after it has been calibrated, resulting in inaccuracy.
The common indicators and their colours
A common error is to suggest using Universal Indicator as a suitable indicator for an acid-base titration.
This is incorrect as a
sharp colour change is required to identify the end-point, which cannot be achieved with Universal Indicator.
Preparing soluble salts
There are two methods of preparing soluble salts
Preparation of Soluble Salts by adding Acid to a solid Metal, Base or Carbonate+
Preparation of Soluble Salts by reacting a Dilute Acid and Alkali
Experiment 6a: Preparation of Soluble Salts by adding Acid to a solid Metal, Base or Carbonate
draw passagees and write method
Method:
Add dilute acid into a beaker and heat using a bunsen burner flame
Add the insoluble metal, base or carbonate, a little at a time, to the warm dilute acid and stir until the base is in excess (i.e. until the base stops disappearing and a suspension of the base forms in the acid)
Filter the mixture into an evaporating basin to remove the excess base
Heat the solution to evaporate water and to make the solution saturated. Check the solution is saturated by dipping a cold, glass rod into the solution and seeing if crystals form on the end.
Leave the filtrate in a warm place to dry and crystallize
Decant excess solution and allow crystals to dry
Experiment 6b: Preparation of Soluble Salts by reacting a Dilute Acid and Alkali
method and draw passages
Method:
Use a pipette to measure the alkali into a conical flask and add a few drops of indicator (phenolphthalein or methyl orange)
Add the acid into the burette and note the starting volume
Add the acid very slowly from the burette to the conical flask until the indicator changes to appropriate colour
Note and record the final volume of acid in burette and calculate the volume of acid added (starting volume of acid – final volume of acid)
Add this same volume of acid into the same volume of alkali without the indicator
Heat to partially evaporate, leaving a saturated solution
Leave to crystallise decant excess solution and allow crystals to dry
Preparing insoluble salts
Insoluble salts can be prepared using…
The solid salt obtained is the …thus in order to successfully use this method the solid salt being formed must be….
a precipitation reaction.
precipitate,
insoluble in water.
