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1. The concentration of calcium ions, Ca$^{2+}$, in aqueous solution can be determined by titration with ethylenediaminetetraacetic acid (EDTA), or by atomic absorption spectrophotometry.

(a) EDTA solution must be standardized before use, by titrating it with a standard solution of calcium chloride.

An accurately weighed mass of powdered calcium carbonate was dissolved in concentrated hydrochloric acid and heated to remove all of the carbon dioxide. After cooling, the calcium chloride solution was made up to exactly 250.0 cm$^3$.

(i) State the most appropriate glassware to make 250.0 cm$^3$ of a Ca$^{2+}$ solution. [1]

(ii) Describe the steps involved in preparing the standard solution by dilution of the initial calcium chloride solution. [2]

(b) Calcium chloride is not used to make the standard solution because it readily absorbs water.

Suggest how solid calcium chloride, CaCl$_2$, could be treated in the laboratory to remove all of the absorbed water. [2]

(c) A sample of tap water was titrated against the standardized EDTA solution, using Eriochrome Black T as an indicator.

The following titration data were obtained.

(i) Suggest two reasons why the volume added in trial 1 is much larger than the other values. [2]

(ii) State the name given to a value that differs significantly from others. [1]

(iii) Suggest a reason why trial 1 should not be included when finding the mean volume of EDTA used. [1]

(iv) Calculate the percentage uncertainty in the volume of EDTA used in trial 2. [1]

(v) The end point of the titration is determined when a red complex, X, changes to a blue ion, Y. Explain how this shows that X absorbs radiation with higher energy than Y. [2]

(d) A standard stock solution of calcium chloride was diluted to produce several solutions with accurately known concentrations. The absorbance of these solutions was measured in a spectrophotometer that was calibrated with distilled water.

{{figure_desc:fig-1|scatter plot of Absorbance vs Concentration of Ca$^{2+}$ (ppm)}}

(i) Draw a line of best fit on the graph. [1]

(ii) Determine the concentration, in ppm, of Ca$^{2+}$ ions in a sample of tap water with an absorbance of 0.090. [1]

(iii) $A = \\log_{10} \\frac{I_0}{I}$ where A is absorbance, $I_0$ is intensity of incident radiation and $I$ is intensity of radiation that has passed through the sample.

Determine the percentage of incident radiation that is absorbed by a sample of tap water with an absorbance of 0.090. [2]

(iv) The 4.00 ppm solution was made by diluting 2.00 cm$^3$ of the stock solution to 25.00 cm$^3$. Calculate the concentration, in ppm, of the stock solution. [1]

(v) Calculate the concentration, in mol dm$^{-3}$, of Ca$^{2+}$ in the 4.00 ppm solution. 1 ppm = 1 mg dm$^{-3}$. [1]