Methanol can be used as a fuel, in a variety of different ways - Scottish Highers Chemistry - Question 7 - 2015

The thermometer should not touch the flame or be placed directly above it because this would cause inaccurate temperature readings. If the thermometer is too close to the flame, the temperature measured would be higher than the actual temperature of the water.

The student should have kept the type of alcohol constant during the experiments, ensuring all trials used the same substance for accurate comparisons.

To calculate the enthalpy of combustion, we use the formula:

$q = mc\Delta T$

Here, the mass of water (m) is 100 g, the specific heat capacity (c) is 4.18 J/g °C, and the temperature rise (ΔT) is 23 °C. Therefore,

$q = 100 imes 4.18 imes 23 = 9614\, J = 9.614\, kJ$

Since 1.07 g of methanol is used, we find moles of methanol as:

$\text{moles of methanol} = \frac{1.07\,g}{32\,g/mol} = 0.0334\, mol\, CH_3OH$

Now, we can find the enthalpy change per mole:

$\text{Enthalpy change} = \frac{9.614\,kJ}{0.0334\,mol} = -288\,kJ/mol$

To calculate the density of methanol, we use the formula:

$\text{Density} = \frac{\text{mass}}{\text{volume}}$

Considering the given mass of alcohol as 19.98 g for a volume of 25.0 cm$^{3}$:

$\text{Density} = \frac{19.98\,g}{25.0\,cm^3} = 0.799\,g\,cm^{-3}$

Knowing whether a reaction is exothermic or endothermic helps chemists in designing processes that are safe and efficient, as exothermic reactions can release energy that needs to be managed to prevent hazards, while endothermic reactions require energy input that must be controlled to optimize production.

Using bond enthalpies to calculate the enthalpy change requires summing the bond enthalpies of bonds broken and subtracting the bond enthalpies of bonds formed. Assuming the necessary values from the booklet, apply the calculation to arrive at the enthalpy change for the reaction: These values should give insight into the energy dynamics of the chemical reaction.