A Laboratory Report: Determination of the Enthalpy Change of a Neutralisation Reaction

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> Note: This is a published student-facing sample produced by Assignment Help Center to show how a first-class undergraduate chemistry laboratory report is structured, argued and referenced. The temperature readings in the Results section are illustrative and were generated for teaching purposes only; they do not describe a real experiment or a real cohort. Every quantity has been independently recomputed in Python and checked by hand. Use this as a model for your own writing, not as a source to cite.

Abstract

The enthalpy change of neutralisation for a strong acid reacting with a strong base was determined using simple constant-pressure (“coffee-cup”) calorimetry. Equal volumes (50.0 cm3) of 1.00 mol dm-3 hydrochloric acid and 1.00 mol dm-3 sodium hydroxide, both at approximately room temperature, were mixed in an insulated polystyrene cup and the maximum temperature rise recorded. Across three illustrative trials the mean temperature rise was 6.6 degrees C. Treating the mixed solution as having the mass and specific heat capacity of water, the heat released was calculated as 2.76 kJ for 0.0500 mol of water formed, giving a molar enthalpy of neutralisation of -55.2 kJ mol-1. This lies about 3.4 per cent below the accepted literature value of approximately -57.1 kJ mol-1 (Atkins and de Paula, 2014), a discrepancy consistent with unavoidable heat loss to the surroundings and the simplifying assumptions of the method. The reaction is confirmed to be strongly exothermic, and the near-constant value expected for strong acid-strong base systems is discussed in terms of the underlying ionic reaction.

Introduction

When an acid reacts with a base, a neutralisation reaction occurs in which hydrogen ions and hydroxide ions combine to form water. For a strong acid and a strong base, both of which are fully dissociated in aqueous solution, the spectator ions take no part in the chemical change and the net ionic equation reduces to a single, simple process:

H+(aq) + OH-(aq) -> H2O(l)

Because this net reaction is the same regardless of which particular strong acid and strong base are used, the molar enthalpy change of neutralisation for strong acid-strong base systems is expected to be approximately constant, close to -57 kJ mol-1 (Atkins and de Paula, 2014; Housecroft and Constable, 2010). The standard enthalpy of neutralisation is defined as the enthalpy change when an acid and a base react to form one mole of water under standard conditions. It is a negative quantity because bond formation in the product releases energy: neutralisation is exothermic.

The experimental determination of this quantity relies on calorimetry, the measurement of heat flow accompanying a physical or chemical change. In constant-pressure calorimetry the heat exchanged is equal to the enthalpy change of the system, since at constant pressure q(p) = delta-H (Atkins and de Paula, 2014). A simple insulated vessel such as an expanded-polystyrene cup approximates an adiabatic system, in which no heat is exchanged with the surroundings; any heat released by the reaction is therefore assumed to be absorbed entirely by the solution, raising its temperature. The heat absorbed by the solution is given by the fundamental calorimetric relationship:

q = m c delta-T

where q is the heat transferred (J), m is the mass of solution (g), c is the specific heat capacity of the solution (J g-1 K-1) and delta-T is the temperature change (K, numerically equal to a change in degrees C). Because the solutions are dilute and largely water, it is conventional to assume that the mixed solution has the density (1.00 g cm-3) and specific heat capacity (4.18 J g-1 K-1) of pure water (Monk, 2004). Dividing the heat released by the number of moles of water formed yields the molar enthalpy of neutralisation, which by the sign convention for an exothermic process is reported as a negative value.

The aim of this experiment was to determine the molar enthalpy change of neutralisation for the reaction between hydrochloric acid and sodium hydroxide and to compare the experimental value with the accepted literature value, thereby evaluating the reliability of simple calorimetry. It was hypothesised that the reaction would be exothermic and that the measured value would be slightly less negative than the literature value because a real calorimeter is not perfectly insulated.

Materials and Method

Apparatus and reagents

The apparatus comprised an expanded-polystyrene cup (approximately 200 cm3 capacity) fitted with a loose lid, standing inside a glass beaker for stability; a thermometer graduated to 0.1 degrees C; two 50 cm3 measuring cylinders; and a stirring rod. The reagents were 1.00 mol dm-3 hydrochloric acid (HCl) and 1.00 mol dm-3 sodium hydroxide (NaOH), both standardised solutions.

Procedure

Using a clean measuring cylinder, 50.0 cm3 of 1.00 mol dm-3 hydrochloric acid was transferred to the polystyrene cup. A separate measuring cylinder was used to measure 50.0 cm3 of 1.00 mol dm-3 sodium hydroxide. Both solutions were left to stand so that they reached the same initial temperature; the temperature of the acid was recorded as the initial temperature. The sodium hydroxide was then added rapidly to the acid, the lid was replaced, and the mixture was stirred gently and continuously with the thermometer. The temperature was monitored and the highest (maximum) temperature reached was recorded as the final temperature. The cup was then emptied, rinsed and dried, and the procedure was repeated to obtain three trials in total. The apparatus was allowed to return to room temperature between runs.

Safety and controls

Both reagents are corrosive at higher concentrations, so eye protection and gloves were worn and spillages were rinsed with plenty of water. To improve reliability, the same thermometer and measuring cylinders were used throughout, the volumes and concentrations were held constant, and the reactants were mixed quickly to minimise heat loss before the maximum temperature was reached. In a more rigorous version of the experiment, a temperature-time graph would be plotted and the maximum temperature obtained by extrapolation back to the moment of mixing, correcting for heat lost during the reaction.

Results

The temperature readings below are illustrative and were generated for teaching purposes; they do not represent a real experiment. They are, however, internally consistent, so the descriptive and calculated quantities agree with one another.

The initial and final temperatures for each trial, and the resulting temperature rise, are presented in Table 1. The three trials gave closely similar temperature rises, and no reading was discarded as anomalous.

Table 1

Initial temperature, maximum temperature and temperature rise for three trials (illustrative data)

The mean temperature rise was 6.6 degrees C, with a range of only 0.2 degrees C across the three trials, indicating good repeatability. The fixed quantities used in the subsequent calculations are summarised in Table 2.

Table 2

Fixed experimental quantities and assumed constants

Calculations

Step 1 – Heat released, q. Assuming the mixed solution behaves as 100.0 g of water (50.0 cm3 + 50.0 cm3 at a density of 1.00 g cm-3):

q = m c delta-T = 100.0 g x 4.18 J g-1 K-1 x 6.6 K = 2758.8 J, approximately 2.76 kJ

Step 2 – Moles of water formed, n. The number of moles of acid and of base are equal, so the reaction is stoichiometric and neither reagent is in excess:

n(HCl) = c x V = 1.00 mol dm-3 x 0.0500 dm3 = 0.0500 mol n(NaOH) = 1.00 mol dm-3 x 0.0500 dm3 = 0.0500 mol

From the 1:1 stoichiometry of H+ + OH- -> H2O, the amount of water formed is 0.0500 mol.

Step 3 – Molar enthalpy of neutralisation, delta-H. The heat released corresponds to 0.0500 mol of water; dividing gives the heat per mole, and the negative sign denotes an exothermic change:

delta-H = – q / n = – (2.7588 kJ) / (0.0500 mol) = -55.2 kJ mol-1

Step 4 – Comparison with the literature value. Taking the accepted value as -57.1 kJ mol-1 (Atkins and de Paula, 2014), the percentage difference is:

percentage difference = (-55.2 – (-57.1)) / (-57.1) x 100 per cent = -3.4 per cent

The experimental value is therefore about 1.9 kJ mol-1 less negative than the literature value.

All quantities above were recomputed independently in Python (q = 2758.8 J; delta-H = -55.18 kJ mol-1; difference -3.37 per cent) and agree with the hand calculations to within rounding.

Discussion

The experiment confirmed that the neutralisation of hydrochloric acid by sodium hydroxide is strongly exothermic, releasing 2.76 kJ of heat for 0.0500 mol of water formed and giving a molar enthalpy of neutralisation of -55.2 kJ mol-1. This value is in good agreement with the accepted figure of approximately -57.1 kJ mol-1 for strong acid-strong base neutralisation (Atkins and de Paula, 2014; Housecroft and Constable, 2010), differing by only 3.4 per cent. The negative sign and the magnitude are both consistent with the underlying net ionic reaction, H+(aq) + OH-(aq) -> H2O(l), which is the same for any fully dissociated acid and base and therefore predicts a near-constant molar enthalpy across such systems (Chang and Goldsby, 2016).

That the experimental value is slightly less negative than the literature value is the expected outcome and supports the stated hypothesis. The dominant source of error is heat loss to the surroundings: although the polystyrene cup is a reasonable insulator, it is not perfectly adiabatic, and some heat escapes through the lid, the thermometer and the walls during the time taken to mix, stir and reach the maximum temperature. Any such loss lowers the measured temperature rise and hence produces a numerically smaller enthalpy value. A related systematic error is that the calorimeter itself absorbs some heat; this experiment assumed the heat capacity of the apparatus to be negligible, whereas a calorimeter constant, determined separately and added to the calculation, would raise the magnitude of the result closer to the literature value (Monk, 2004).

Further systematic error arises from the simplifying assumptions about the solution. Treating the mixed solution as pure water – with a density of exactly 1.00 g cm-3 and a specific heat capacity of 4.18 J g-1 K-1 – is only approximate, since dilute salt solutions have slightly higher densities and slightly lower specific heat capacities than water. These approximations introduce a small error into both the assumed mass and the value of c. Random errors are comparatively minor here: the three trials agreed to within 0.2 degrees C, and the thermometer’s resolution of 0.1 degrees C on a rise of 6.6 degrees C corresponds to a reading uncertainty of only about plus or minus 1.5 per cent, smaller than the observed deviation from the literature value. This pattern – small random scatter but a consistent one-directional shift – points to systematic heat loss rather than measurement noise as the principal limitation.

Several improvements would increase accuracy. Plotting a temperature-time graph and extrapolating the cooling portion back to the instant of mixing would correct for heat lost before the maximum was recorded, typically yielding a larger, more accurate temperature rise. Determining and applying a calorimeter constant, using a lid with a smaller aperture, and employing a data-logging temperature probe to capture the true maximum more reliably would each reduce systematic error. Repeating the experiment with a different strong acid-strong base pair, such as nitric acid and potassium hydroxide, would also test the prediction that the molar enthalpy of neutralisation is essentially independent of the identity of the strong acid and base.

Conclusion

Using simple constant-pressure calorimetry, the molar enthalpy change of neutralisation for hydrochloric acid and sodium hydroxide was determined to be -55.2 kJ mol-1 from an illustrative dataset. This is within 3.4 per cent of the accepted literature value of approximately -57.1 kJ mol-1, and the small negative discrepancy is fully consistent with heat loss from an imperfectly insulated calorimeter and with the simplifying assumptions made about the solution. The experiment therefore confirms that strong acid-strong base neutralisation is strongly exothermic and demonstrates that even a basic coffee-cup calorimeter can yield a molar enthalpy value close to the accepted figure, provided its systematic limitations are recognised and, where possible, corrected.

References

Atkins, P.W. and de Paula, J. (2014) Atkins’ Physical Chemistry. 10th edn. Oxford: Oxford University Press.

Chang, R. and Goldsby, K.A. (2016) Chemistry. 12th edn. New York: McGraw-Hill Education.

Hale, J.D., Izatt, R.M. and Christensen, J.J. (1963) ‘A calorimetric study of the heat of ionization of water at 25 degrees’, The Journal of Physical Chemistry, 67(12), pp. 2605-2608. [VERIFY page range and issue]

Housecroft, C.E. and Constable, E.C. (2010) Chemistry: An Introduction to Organic, Inorganic and Physical Chemistry. 4th edn. Harlow: Pearson Education.

Levine, I.N. (2009) Physical Chemistry. 6th edn. New York: McGraw-Hill.

Monk, P. (2004) Physical Chemistry: Understanding our Chemical World. Chichester: John Wiley & Sons.

Rossini, F.D. (1931) ‘The heat of formation of water and the heat of neutralization’, Bureau of Standards Journal of Research, 6(1), pp. 1-35. [VERIFY exact title, volume and pages]