Making up a standard solution
Stage 1: transferring a [blank_start]known[blank_end] mass of the solid
• Weigh the solid into a weighing boat recording the mass (to appropriate precision).
• Transfer the solid to a [blank_start]beaker[blank_end].
• Re-weigh the weighing boat and record the [blank_start]difference[blank_end] in mass.
Stage 2: dissolving the solid in distilled water
• Add the [blank_start]minimum[blank_end] distilled water to the beaker to dissolve.
• Stir the beaker with the [blank_start]glass rod[blank_end] until the solid has dissolved.
Stage 3: rinsing all glassware and making up the solution
• Transfer the solution to a [blank_start]volumetric flask[blank_end] using a funnel.
• Rinse all used glassware into the volumetric flask (glass rod, beaker and [blank_start]funnel[blank_end]).
• Add [blank_start]distilled[blank_end] water up to the [blank_start]graduation line[blank_end] (on the volumetric flask).
• [blank_start]Invert[blank_end] the flask several times.
Which of the following should you NOT do when making up a standard solution?
Weigh the weighing boat with the solid in it, and then again after the solid has been transferred to a beaker.
Invert the volumetric flask before using the solution.
Add distilled water to just above the graduation line, then use a dropping pipette to remove the extra.
Any glassware that has been in contact with the solution should be rinsed into the volumetric flask when making up a standard solution.
Carrying out an acid-base titration
Stage 1: Preparing the conical flask
• Using a [blank_start]glass pipette[blank_end], transfer the desired amount of the acidic solution into the conical flask.
• Add several drops of an appropriate [blank_start]indicator[blank_end] to the conical flask.
• Rinse any solution on the sides of the flask into the flask with [blank_start]distilled water[blank_end].
Stage 2: preparing the burette
• Fill the burette up with the basic solution using a [blank_start]funnel[blank_end].
• Run the solution through the jet, ensuring that there are no [blank_start]air bubbles[blank_end] present.
• Remove the funnel from the burette.
Stage 3: performing the titration
• Place the conical flask on a [blank_start]white tile[blank_end] beneath the burette.
• Record the starting volume of the burette.
• Titrate the solution, swirling the [blank_start]conical flask[blank_end] until a [blank_start]permanent[blank_end] colour change is observed.
• Record the end volume in the [blank_start]burette[blank_end] and calculate the titre.
• Repeat until [blank_start]concordant[blank_end] results (titre values within 0.1 cm³ of each other) are achieved and calculate a mean titre (never include [blank_start]rough[blank_end] values).
Using distilled water to rinse the sides of the conical flask in an acid-base titration will negatively affect your results.
Why do you remove the funnel from the burette during a titration?
So that any remaining droplets of solution do not fall into the burette and so change the volume/titre value.
To stop you knocking it out/it falling out when adjusting the clamp to allow you to read the meniscus.
So that it can be used by another group.
Which of the following is NOT a suitable indicator for use in ANY titration?
Potassium permanganate solution
Methyl orange solution
Universal indicator solution
Experimentally finding the enthalpy change of combustion
Stage 1: preparing the chemicals
• Add the fuel to the [blank_start]spirit burner[blank_end] and record the initial mass.
• Add a known quantity of water to the [blank_start]calorimeter[blank_end] and clamp in position above the spirit burner.
• Use thermometer to measure the initial temperature of the [blank_start]water[blank_end].
Stage 2: carrying out the experiment
• Use the spirit burner to heat water in the calorimeter to roughly (40ᵒC) and record the [blank_start]temperature change[blank_end].
• Reweigh the spirit burner and calculate the [blank_start]change in mass[blank_end] of the fuel.
Stage 3: interpreting the results
• Calculate the energy released ([blank_start]J[blank_end]) using the equation q = mcΔT where q = [blank_start]energy change[blank_end]; m = mass of water heated; c = [blank_start]specific heat capacity of water[blank_end] (4.18); ΔT = change in temperature.
• Calculate the enthalpy change ([blank_start]Jmol⁻1[blank_end]) using the equation ΔH = q / (moles of fuel).
• Convert into [blank_start]kJmol-1[blank_end] by dividing your answer by 1000.
change in mass
specific heat capacity of water
Why are data book values often higher than those calculated experimentally?
Heat is lost to the surroundings, meaning the temperature change is smaller than it should be and so the enthalpy change value is smaller.
People make mistakes when doing experiments - the data book values are the values we would get if we always added the correct amount of water etc.
Often, experimental values are not calculated under standard conditions.
The technique has changed leading to more accurate values, but the data books have not been updated yet.
The fuel will sometimes undergo incomplete combustion, meaning that the full amount of energy is not released.
Alcohols are very volatile and so some will evaporate from the wick. The value calculated for the number of moles of fuel burnt will be higher than the true value, meaning the value for enthalpy change will be lower than the true value.
It is important to use a very precise thermometer when calculating the enthalpy change of a reaction.
Which of the following are ways in which the calorimetry method can be improved to minimise heat loss? CHOOSE 3
Reduce the distance between the flame and the calorimeter
Use a glass rod to stir the water/solution in the calorimeter
Put a draught screen around the flame
Use a larger spirit burner
Put a lid on the beaker/calorimeter
Experimentally finding the enthalpy change of a reaction, accounting for heat loss
Add a (5.00g) sample of solid to (50g) of [blank_start]water[blank_end] and start the stop watch.
Record the [blank_start]temperature[blank_end] of the solution every subsequent minute for approximately 5 minutes.
Plot a graph of temperature (y-axis) versus time (x-axis).
[blank_start]Extrapolate[blank_end] back to the time of [blank_start]mixing[blank_end] (T=0s) to establish the [blank_start]maximum[blank_end] temperature change.
Calculate ΔH in kJ mol-1 using the equation: ΔH = (m × c × ∆T)/(n ×1000 )
Why is the line of best fit extrapolated back to the time of mixing when calculating enthalpy change?
To account for the heat lost to the surroundings
To eliminate the possibility of anomalous results
To find the best temperature at which to carry out the reaction
Investigate how the rate of reaction changes with temperature
1. Set a [blank_start]water bath[blank_end] to 20ᵒC.
2. Accurately measure out 20cm³ of 1.0 mol dm⁻³ hydrochloric acid and pour into a [blank_start]conical flask[blank_end].
3. Allow the solution of hydrochloric acid to reach the [blank_start]temperature[blank_end] of the water bath.
4. [blank_start]Accurately[blank_end] weigh out 1.0g of magnesium ribbon.
5. Place the magnesium into the conical flask, immediately attaching the [blank_start]gas syringe[blank_end] and starting a timer.
6. Record the [blank_start]volume[blank_end] of gas evolved in the first 20 seconds.
7. Repeat steps 1 – 6 for a [blank_start]range[blank_end] of temperatures (30ᵒC – 60ᵒC).
8. Plot a graph of temperature ([blank_start]x-axis[blank_end]) versus initial rate (cm³/20s) ([blank_start]y-axis[blank_end]).
It is not possible to accurately measure the effect of temperature on rate of reaction at temperatures above 60ᵒC.