solubility and titration experiment

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• Creating a Saturated Solution
• Titration of Borax Solution

Source: Smaa Koraym at Johns Hopkins University, MD, USA

In this experiment, you will create a saturated solution of sodium tetraborate decahydrate, also called borax. In water, borax dissociates into two sodium cations and one tetraborate anion. When an aqueous borax solution is saturated, it means that it contains the maximum amount of dissolved solute, which is borax, for that volume of solvent, which is water. Any additional solid won't appear to dissolve.

Since solubility is temperature-dependent, each group will create a saturated solution at a different temperature. More borax will dissolve at higher temperatures, resulting in a solution with a higher borax concentration.

• Before starting the experiment, put on the necessary personal protective equipment, including a lab coat, chemical splash goggles, and gloves.

Table 1: Temperature and Volume Data

• Label a 150-mL beaker ‘borax solution’, and transfer the borax into it. Measure 50 mL of deionized water and pour it into the beaker. Add a magnetic stir bar to the solution, set the beaker on the hotplate, and turn on the stir setting.
• Use a digital thermometer clamp to hold a digital thermometer inside the beaker and above the stir bar. Heat the beaker to about 10 to 20 °C higher than the assigned temperature. Note: Students assigned the 10 °C temperature will let the borax dissolve at room temperature.
• For all assigned temperatures, add 200 mL of deionized water to a 400-mL beaker. Then, place a glass thermometer in the beaker of water. Set the beaker on the hot plate and heat the water to 20 °C higher than your assigned temperature. Note: Students assigned the 10 °C temperature should use room temperature water.
• Label the five 250-mL Erlenmeyer flasks as ‘trials 1 through 5’. Once the borax has fully dissolved in the solution and it appears homogeneous, remove the beaker from the stir plate and set it on the benchtop to cool to the assigned temperature.
• Once the temperature has been reached, place the beaker on a stack of paper towels to act as an insulator and maintain a constant temperature. Note: Students assigned to the 10 °C temperature should place their beaker containing borax solution into a larger beaker filled with ice and water.
• Label another 400-mL beaker as ‘waste’, and pipette 5 mL of the hot water into your waste beaker. Repeat this several times to warm up the pipette and prevent it from getting clogged as the borax cools.
• Record the exact temperature of the borax solution in your notebook. Then, pipette 5 mL of the saturated borax solution from the top of the beaker and dispense it into the first 250-mL Erlenmeyer flask. Do not allow the pipette to come into contact with the crystallizing borax at the bottom of the beaker.
• Pipette 5 mL of the hot water into the Erlenmeyer flask containing the borax solution. Repeat this for each of the four remaining flasks, adding 5 mL of borax solution and 5 mL of water. Remember to keep an eye on the temperature of your borax solution, as it will cool on the benchtop. Note: The temperature of the saturated solution transferred into each of the five flasks should be within 2 – 3 °C of each other. Students working with the higher temperatures may need to reheat their solution before proceeding.

Now you will determine how much borax has dissolved in your saturated solution. Recall the chemical reaction showing how borax dissociates in water, forming the tetraborate ion. Since the tetraborate ion is a base, it will react with acid following a neutralization reaction.

When the amount of acid is twice the amount of tetraborate, the solution is neutralized. To do this, we will slowly dispense HCl into the borax solution until it is neutralized, meaning that the acid and base react to form water and salt and a neutral pH. We'll use the pH indicator bromocresol green to let us know when the solution is neutralized, as it turns from blue to pale greenish-yellow when the pH is neutral.

• To begin the titration, first, measure 40 mL of deionized water and add it to one of the Erlenmeyer flasks. Swirl the flask to make sure all of the borax is in solution. Repeat this for the other four flasks.
• Add 2 - 3 drops of 0.1% bromocresol green to each flask.
• Clip a burette clamp on a ring stand and secure a burette in the clamp while making sure that the burette is vertical and as straight as possible. Make sure that the stopcock on the burette is closed.
• Place a funnel in the open end of the burette and fill the burette with deionized water. Set your waste beaker under the burette and open the stopcock to let all of the water rinse out of the burette. Then, close the stopcock.
• Obtain 0.5 M HCl from your instructor and pour 10 mL into the burette. Use the markings on the side of the burette as a guide. Open the stopcock on the burette to allow all of the HCl to drain into the waste beaker
• Close the stopcock again, and then fill the burette with 50 mL of the HCl. Open the stopcock slightly to allow the liquid to fill the tip of the burette and remove any bubbles.
• Record the initial volume of HCl in your lab notebook.
• For the first trial, titrate the HCl into the borax solution in increments of 1 mL. Gently swirl the flask after each addition to make sure that the solution is well mixed. Note: At the endpoint of the titration, the indicator will turn the solution from light blue to a pale greenish-yellow color. If your solution turns a dark yellow color, this indicates that you have passed the endpoint.
• When you have reached the endpoint, record the volume of HCl remaining in the burette.
• Refill the burette to the 50-mL mark with more HCl and repeat the titration for all of the other flasks of borax solution. Make sure to record the final volumes of HCl used for each titration.
• To clean up from the experiment, place the 400-mL waste beaker under the burette spout and open the burette to drain the remaining HCl into it.
• Fill the burette with deionized water and allow it to rinse through the burette.
• Use your remaining borax solution to neutralize the acid in your waste beaker. Swirl the beaker until it stops bubbling.
• Add baking soda to the waste beaker and swirl the solution. Continue adding baking soda and swirling the solution until it stops bubbling. Note: You may not observe bubbling because the borax does most of the neutralization.
• Wash the contents of the waste beaker down the sink with copious amounts of water. The flasks containing borax and HCl from your titrations are neutralized, so they can also be poured down the sink.

Table 2: Determining Δ G, Δ H, and Δ S

• Obtain data from other groups so that you have an average HCl volume for each temperature.
• The solubility constant is determined from the stoichiometry of the dissociation reaction of borax and water. Since the reactant is solid, it is not included in the expression. Because two sodium cations are formed for every tetraborate, it is assumed that the concentration of sodium is equal to twice the concentration of tetraborate. Thus, the expression can be simplified as shown: K sp = 4[[B 4 O 5 (OH) 4 ] 2- ] 3
• The concentration of tetraborate can be calculated from the volume of HCl used to reach the endpoint of the titration and neutralize the base. The moles of HCl equal the volume of HCl times the molarity of the solution. Since two moles of HCl are needed to neutralize one mole of tetraborate, the moles of tetraborate can easily be calculated.
• Calculate the concentration of tetraborate (based on 5 mL volume of saturated solution) and determine the reaction constant. Repeat the calculation for each temperature and compare your results.
• Use the titration data to calculate the Gibbs free energy, ΔG, at each temperature value to determine whether the dissociation of borax is a spontaneous reaction. Recall the ΔG expression, where R is the gas constant, T is the temperature of the solution in Kelvin, and K sp is our reaction constant. Using our reaction constants and the corresponding temperature, calculate ΔG for each temperature.
• Remember that if ΔG is positive, the reaction is not spontaneous, meaning that energy needs to be put into the reaction for it to proceed. However, if ΔG is negative, the reaction proceeds spontaneously. In general, this reaction is not spontaneous at low temperatures, but it is spontaneous at higher temperatures. This supports the theory that borax prefers the salt crystal structure form at room temperature and lower, but it prefers to go into solution after a certain temperature is reached. Use K sp to calculate ΔH and ΔS.
• Plot the natural log of K sp as a function of 1/T reported in Kelvin. The slope of this line equals -ΔH/R, so calculate ΔH. The intercept of the line equals ΔS/R. Determine ΔS.
• Since the change in enthalpy of this reaction is about 90 kJ/mol and is positive, the reaction is endothermic, meaning that it absorbs energy. The change in entropy is positive and is about 290 J/mol·K, which indicates the favorable production of disorder. This is expected as the crystal structure of the salt breaks down.

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Ca(OH) 2 ⇌ Ca 2+ + 2OH¯

K sp = [Ca 2+ ] [OH¯] 2

molarity = moles ÷ volume (in liters) 0.0250 mol/L = x ÷ 0.02250 L x = 0.0005625 mol

0.0005625 mol / 0.02500 L = 0.0225 mol/L

it is exactly half the [OH¯]. This is because of the 1:2 molar ratio from the balanced equation.

K sp = (0.01125) (0.0225) 2 = 5.70 x 10¯ 6

moles HCl = (0.102 mol / L) (0.00813 L) = 0.00082926 mol

0.00082926 mol / 2 = 0.00041463 mol Remember, every one Ca(OH) 2 titrated requires 2 H +

0.00041463 mol times 74.0918 g/mol = 0.030721 g This is grams per 25.0 mL

0.030721 g / 0.0250 L = 1.23 g/L (to three sig fig)

[Ca 2+ ] = 0.00041463 mol / 0.0250 L = 0.0165852 M [OH¯] = 0.0165852 M times 2 = 0.0331704 M

K sp = (0.0165852) (0.0331704) 2 = 1.82 x 10¯ 5

moles HCl = (0.000050 mol / L) (0.00670 L) = 0.000000335 mol

0.000000335 mol / 2 = 0.0000001675 mol Remember, every one Pb(OH) 2 titrated requires 2 H +

[Pb 2+ ] = 0.0000001675 mol / 0.0250 L = 0.0000067 M [OH¯] = 0.0000067 M times 2 = 0.0000134 M

K sp = (0.0000067) (0.0000134) 2 = 1.20 x 10¯ 15

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AP®︎/College Chemistry

Course: ap®︎/college chemistry   >   unit 4.

• Acid–base titrations

Worked example: Determining solute concentration by acid–base titration

• Redox titrations
• Introduction to titration

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Skills to Develop

• Perform and interpret titration calculations

A titration is a laboratory technique used to precisely measure molar concentration of an unknown solution using a known solution. The basic process involves adding a standard solution of one reagent to a known amount of the unknown solution of a different reagent. For instance, you might add a standard base solution to an mystery acid solution. As the addition takes place, the two reagents in the solutions, in this the acid and base, react. You also add an indicator , which is a molecule that changes color when the original reagent (the acid in the mystery solution, say) is completely consumed by reaction with the standard solution reagent. If you know exactly how much standard was added before the color change, you can calculate how many moles of the unknown were present at the beginning, and thus the concentration of the unknown.

Many of the standard reagents you might use in the lab, especially HCl and NaOH, which are very common and important, are hard to prepare at precise concentration without titration. The reason is that HCl is purchased as a concentrated solution, which can vary a little in concentration because both the HCl and the water can evaporate. NaOH can be purchased as a solid, but it is hygroscopic which means that it absorbs water from the air. It can absorb so much water that it actually dissolves. For this reason, even if you buy it dry, once you open the bottle, it might start to absorb water, and it would be difficult to know when you measure it what % water it is. Thus, if you work in a biochemistry lab, for instance, you might want to control the pH of your solutions by adding a little bit of dilute HCl or NaOH, because chloride and sodium ions are very common and probably are already included in the solution, but you might want to know how concentrated your solutions are. To determine this, you would use a standard solution made of some easier-to-mass acid or base to titrate the solution you actually want to use. Once titrated, you could dilute it precisely to the concentration you want. Some other reagents you might want standard solutions of react with air; these you might also titrate if they have been waiting a long time so you know what the current concentration is.

Titrations might seem a little old-fashioned. Actually, the number of automated titration machines available (try a google search!) suggest that titrations are still very important in industry. One reason might be that titrations can be good for studying newly discovered molecules, for instance to measure the molecular weight and other properties that we will study more later.

Traditionally, you take a known mass or volume of the unknown solution and put it in a flask with the indicator. Then you add the standard solution in a buret, which is a special tube for adding solution slowly and measuring the volume added at the end. These days, it might be easier to use a plastic squeeze bottle instead of a buret. You put the standard solution in the squeeze bottle, get the mass of the bottle, do the titration, and then mass the bottle again. Now you know exactly how much standard was added!

Example \(\PageIndex{1}\)

We have a solution of HCl whose concentration is known imprecisely (~2.5 M). (We made this solution in the previous section on molarity.) We want to determine the concentration more precisely. We have a solution of NaOH that is known to be 5.1079 M. We place 100.00 ml of the HCl solution in a flask with a drop of an indicator that will change color when the solution is no longer acidic. Then we add NaOH slowly until the indicator color changes. At this point, we have added 46.67 ml NaOH. Calculate the precise concentration of the HCl.

To answer, we need to know that the reaction is

\[HCl + NaOH \rightarrow NaCl + H_{2}O\]

So the ratio is 1 HCl:1 NaOH. We calculate the number of moles of NaOH added:

\[(5.1079\; mol/L)(46.67\; mL)= 238.4\; mmol\]

This is also the number of moles of HCl in the original 100.00 mL of solution, because the reaction ratio is 1:1. To calculate the concentration of the HCl solution, we just divide the number of moles of HCl by the volume.

\[(238.4\; mmol)/(100.00\; mL) = 2.384\; M\]

We could do this in one step using dimensional analysis:

\[(46.67\; \cancel{mL\; NaOH}) \left(\dfrac{5.1079\; \cancel{mol\; NaOH}}{1000\; \cancel{mL}}\right) \left(\dfrac{1\; mol\; HCl}{1\; \cancel{mol\; NaOH}}\right) \left(\dfrac{1}{100.00\; mL}\right)=2.384\; M\]

Now, we diluted 250 mL of the original stock solution to 1.00 L to make this solution. What is the concentration of the 10 M HCl, precisely? First we calculate the moles of HCl in the whole "2.5 M" solution, which is equal to the number of moles of HCl in quantity of stock solution ("10 M") that was used to make it (250ml).

\[(1.00\; \cancel{L\; of\; diluted\; solution}) \left(\dfrac{2.384\; mol\; HCl}{1\; \cancel{L\; of\; diluted\; solution}}\right) \left(\dfrac{1}{0.250\; L\; of\; conc.\; solution}\right)=9.536\; M\]

Why so low? HCl is a gas, and can evaporate out of solution. The stock solution must have been pretty old.

Contributors and Attributions

Emily V Eames (City College of San Francisco)

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Acid–base back titration

By Sandrine Bouchelkia and Jo Haywood

• Five out of five

Experiment and develop learners’ quantitative analysis skills

This resource accompanies the article The essential guide to teaching quantitative chemistry in  Education in Chemistry , where you will find tips, common misconceptions and further ideas for teaching reacting masses and limiting reagents. 

Use the technician notes to prepare the practical and find the experimental procedure in the student worksheet.

Download this

Use the experiment and accompanying resources, including answers, to develop your post-16 learners' quantitative chemistry skills.

Learning objectives

• Apply your knowledge of mole calculations to practical experimental data.
• Write balanced equations for practical experiments.
• Review practical techniques.

Learners will address all three objectives throughout the experiment. They should successfully carry out the practical, write relevant balanced equations and utilise these to carry out mole calculations. Learners will draw their findings together in a conclusion and evaluation, comparing their experimental and theoretical data. Find the answers in the teacher notes and use the accompanying mole calculator spreadsheet to check learners’ calculations.

Scaffolding

Some learners will be keen to work through the whole task independently. Others will benefit from you checking the calculation steps as they go along, ensuring they have correctly balanced the equations and found the masses and moles before they write their conclusion and evaluation. Use the PowerPoint  (also available as a  pdf ) to model each step of the calculation. Remove the example tables in the student sheet to encourage learners to draw their own or leave them in for support.

As an extension, ask learners to titrate the original HCl against the NaOH to check the concentration is correct. They can also investigate the effect of CO 2 dissolving on the pH of water .

Safety and hazards

• Read our standard health and safety guidance  and carry out a risk assessment before running any live practical.
• Eye protection must be worn.
• The flask containing hydrochloric acid and calcium carbonate may get warm.
• Instruct learners to take care not to spill solutions, particularly phenolphthalein, on their skin. If they do get any on their skin, rinse well.
• Fill the burette at eye level.

Do the experiment individually or as a pair/small group. Each learner/group will require:

• Two 100 ml conical flasks
• 25 ml or 50 ml pipette
• 10 ml pipette
• Pipette fillers sized to match pipettes
• Mass balance measuring three decimal places (0.001 g)*
• Clamp and retort
• Cotton wool
• Safety equipment: safety glasses

*If a mass balance measuring three decimal places is not available then use a mass balance measuring two decimal places (0.01 g). This will be less accurate though.

• Hydrochloric acid, 1.00 mol dm -3
• Calcium carbonate chips (approx 1.00 g)
• Sodium hydroxide, 0.400 mol dm -3
• Phenolphthalein indicator solution in dropper bottles

More resources

• Develop learners’ practical skills by ensuring your feedback is meaningful with the  Common Practical Assessment Criteria (CPAC) tracker .
• Use the  Quantitative chemistry starters to practice skills relating to concentration, mass and gas calculations with your 16–18 learners. 
• Bring together organic chemistry and moles calculations with the  Aspirin screen experiment .
• Review your 14–16 year-old learners’ understanding of key ideas, words and phrases relating to quantitative chemistry by downloading the  scaffolded, partially scaffolded and unscaffolded worksheets .
• Solidify you knowledge with the  Quantitative chemistry online CPD course  and download the accompanying resources  too.
• Prepare tables for your results before you start your experiments.
• Using the pipette, accurately measure out 50 ml of 1.00 mol dm -3 hydrochloric acid and add this to the conical flask. Place a piece of cotton wool in the top of the flask.
• Measure the mass of the flask, cotton wool and hydrochloric acid. Record this in the table.
• Using the three decimal places (dp) mass balance, accurately weigh out between 1.000 and 1.500 g (1.00 to 1.50 g for the two dp mass balance) of calcium carbonate and record the mass in the table.
• Add the calcium carbonate to the hydrochloric acid and swirl the flask until all of the calcium carbonate has reacted.
• Measure the final mass of the flask and contents and record in the table.
• Fill the burette with sodium hydroxide solution (0.400 mol dm -3 ) using a funnel and pouring at your eye level.
• Using a 10 ml pipette, measure out 10 ml of your reaction mixture into a clean 100 ml conical flask.
• Add a few drops of phenolphthalein to the flask containing your reaction mixture. The mixture should remain colourless.
• Titrate your mixture with the sodium hydroxide from your burette, recording how much sodium hydroxide you require to turn the mixture pink.
• Rinse the conical titration flask with distilled water.
• Repeat steps 6–11 until you have two results, excluding the rough titration, which are within 0.05 ml of each other (or you run out of reaction mixture).
• Calculate the average titre to use in the calculations.

Acid–base back titration student sheet

Acid–base back titration teacher notes, acid–base back titration technician notes, acid–base back titration calculation slides, acid–base back titration mole calculator, additional information.

Resource created by Jo Haywood. Technician notes adapted by Sandrine Bouchelkia.

• 16-18 years
• Practical experiments
• Presentation
• Teacher notes
• Technician notes
• Maths skills
• Practical skills and safety
• Investigation
• Applying scientific method
• Acids and bases
• Quantitative chemistry and stoichiometry
• Asking scientific questions
• Observing and measuring
• Recording data
• Interpreting data
• Reactions and synthesis

Specification

• C5.1b describe the technique of titration
• Students should be able to: describe how to carry out titrations using strong acids and strong alkalis only (sulfuric, hydrochloric and nitric acids only) to find the reacting volumes accurately
• (HT) Calculate the chemical quantities in titrations involving concentrations in mol/dm³ and in g/dm³.
• C5.4.7 describe and explain the procedure for a titration to give precise, accurate, valid and repeatable results
• C5.3.6 describe and explain the procedure for a titration to give precise, accurate, valid and repeatable results
• AT f: Use acid–base indicators in titrations of weak/strong acids with weak/strong alkalis.
• Titrations of acids with bases.
• Students should be able to perform calculations for these titrations based on experimental results.
• 11. be able to calculate solution concentrations, in mol dm⁻³ and g dm⁻³, including simple acid-base titrations using a range of acids, alkalis and indicators. The use of both phenolphthalein and methyl orange as indicators will be expected.
• 4. use laboratory apparatus for a variety of experimental techniques, including:
• titration, using burette and pipette
• 6. use acid-base indicators in titrations of weak/strong acids with weak/strong alkalis
• di) use of laboratory apparatus for a variety of experimental techniques including: i) titration, using burette and pipette
• f) use of acid–base indicators in titrations of weak/ strong acids with weak/strong alkalis
• d) the techniques and procedures used when preparing a standard solution of required concentration and carrying out acid–base titrations
• carrying out a back titration
• quantitative stoichiometric calculations
• Back titration is used to find the number of moles of a substance by reacting it with an excess volume of reactant of known concentration.
• The resulting mixture is then titrated to work out the number of moles of the reactant in excess.
• From the initial number of moles of that reactant the number of moles used in the reaction can be determined, making it possible to work back to calculate the initial number of moles of the substance under test.
• A back titration is useful when trying to work out the quantity of substance in a solid with a low solubility.
• Stoichiometry is the study of quantitative relationships involved in chemical reactions.
• The ability to balance and interpret equations enabling calculations to be carried out involving any of the above skills/techniques is an important part of chemistry at this level and is examinable in both the Unit and Course assessments.
• process and analyse data using appropriate mathematical skills as exemplified in the mathematical appendix for each science
• consider margins of error, accuracy and precision of data
• follow written instructions
• make and record observations
• keep appropriate records of experimental activities
• present information and data in a scientific way
• evaluate results and draw conclusions with reference to measurement uncertainties and errors
• comment on experimental design and evaluate scientific methods
• use laboratory apparatus for a variety of experimental techniques including: titration, using burette and pipette
• use acid-base indicators in titrations of weak/strong acids with weak/strong alkalis
• carry out experimental and investigative activities, including appropriate risk management, in a range of contexts
• evaluate methodology, evidence and data, and resolve conflicting evidence
• (f) acid-base titrations
• PRACTICAL: Back titration, for example, determination of the percentage of calcium carbonate in limestone
• (f) relationship between grams and moles
• (g) concept of concentration and its expression in terms of grams or moles per unit volume (including solubility)
• (j) concept of stoichiometry and its use in calculating reacting quantities, including in acid-base titrations
• 5.3.4 demonstrate understanding of the method of back titration, for example by determining the purity of a Group II metal oxide or carbonate.
• determine the purity of a Group II metal oxide or carbonate by back titration;
• use volumetric flasks to prepare standard solutions of various volumes and safely use a burette and a pipette with acid-base indicators to carry out titrations of weak or strong acids with weak or strong alkalis;
• Calculations involving excess of one reactant.
• Apparatus used in volumetric analysis.
• Correct titrimetric procedure.
• Acid-base titrations.

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