Content

• To step
• Part 1 of 3: The basic knowledge of concentrations
• Part 2 of 3: Titration
• Part 3 of 3: Determining the salinity in an aquarium
• Tips
• Warnings

In Chemistry or Chemistry, one solution a homogeneous mixture of two things - one dissolved substance and a solvent or solvent in which the substance is dissolved. Concentration is a measure of the amount of solute in a solvent. There can be many reasons for determining the concentration of a solution, but the chemistry involved is the same whether you are testing the chlorine level in a pool or performing a life-saving analysis on a blood sample. This guide will teach you some fundamental parts of solution chemistry, then walk you through the procedure of a common, practical application - aquarium maintenance.

To step

Part 1 of 3: The basic knowledge of concentrations

1. Notation method of concentrations. A concentration of a substance is the amount of that solute divided by the amount of solvent. However, because there are different ways of expressing the amount of a given substance, it is also possible to represent a concentration in different ways. Here you will find the most common spellings:
   • Gram per liter (g / L.) The mass of a solute in grams dissolved in a given volume of a solution (which is not necessarily the same as the volume of the solvent.) Typically used for solutions of solids in liquid solvents.
   • Molarity (M.) The number of moles of a solute divided by the volume of the solution.
   • Parts per million (ppm.) The ratio of the number of particles (usually in grams) of a solute per one million particles of a solution, multiplied by 10. Typically used for very dilute water solutions (1 L of water = 1000 grams.)
   • Percentage of composite substance. The ratio of particles (again in grams) of a solute per 100 particles of a solution, expressed as a percentage.
2. Know what data you need to find a concentration. Except for molarity (see below), the common ways of writing a concentration as indicated above require you to know the mass of the solute and the mass or volume of the resulting solution. Many chemical problems that require finding the concentration of a solution do not give you this information. If so, you will have to work with what you do know to find out this information.
   • Example: Suppose we need to find the concentration (in grams per liter) of a solution made by dissolving 1/2 teaspoon of salt in 2 liters of water. We also know that 1 teaspoon of salt is about 6 grams. In this case, the conversion is easy - multiply: 1/2 teaspoons x (6 grams / 1 teaspoon) = 3 grams of salt. 3 grams of salt divided by 2 liters or water = 1.5 g / L
3. Learn how to calculate molarity. Molarity requires you to know the number of moles of your solute, but this can be easily deduced if you know the mass of the solute and the chemical formula. Each chemical element has a known "molar mass" (MM) - a specific mass for one mole of that element. These molar masses are found in the periodic table (usually under the chemical symbol and name of the element.) Simply add the molar masses of the components of the solute to get the molar mass. Then multiply the known mass of the solute by the (1 / MM of your solute) to find the amount of your solute in moles.
   • Example: Suppose we want to find the molarity of the above saline solution. Just to recap, we have 3 grams of salt (NaCl) in 2 liters of water. Start by finding out what the molar masses of Na and Cl are by looking at the periodic table. Na = about 23 g / mol and Cl = about 35.5 g / mol. Thus, the MM of NaCl = 23 + 35.5 = 58.5 g / mol. 3 grams of NaCl x (1 mole NaCl / 58.5 g NaCl) = 0.051 mole NaCl. 0.051 mole NaCl / 2 liters water = .026 M NaCl
4. Practice standard exercises on calculating concentrations. The above knowledge is all you need to calculate the concentrations in simple situations. If you know the mass or volume of the solution and the amount of solute added in principle, or you can deduce this from the information provided in the statement, you should be able to measure the concentration of a solution with ease. to calculate. Make practice problems to improve your skills. See the example exercises below:
   • What is the molarity of NaCL in a 400 ml solution, obtained by adding 1.5 grams of NaCl to water?
   • What is the concentration, in ppm, of a solution made by adding 0.001 g of lead (Pb) to 150 L of water? (1 L of water = 1000 grams) In this case, the volume of the solution will increase by a minuscule amount by adding the substance, so you can use the volume of the solvent as the volume of the solution.
   • Find the concentration in grams per liter of a 0.1 L solution made by adding 1/2 mole KCl to water. For this problem, you have to work from front to back, using the molar mass of KCL to calculate the number of grams of KCl in the solute.

Part 2 of 3: Titration

1. Understand when to apply a titration. Titration is a technique used by chemists to calculate the amount of solute present in a solution. To perform a titration, you create a chemical reaction between the solute and another reagent (usually also dissolved). Since you know the exact amount of your second reagent and you know the chemical equation of the reaction between the reagent and the solute, you can calculate the amount of your solute by measuring how much of the reagent you need for the reaction with the solute is complete.
   • So, titrations can be very useful in calculating the concentration of a solution if you don't know how much solute was initially added.
   • If you do know how much of a solute is in the solution, then there is no need to titrate - just measure the volume of your solution and calculate the concentration, as described in Part 1.
2. Set up your titration equipment. To perform accurate titrations you need clean, accurate and professional equipment. Use an Erlenmeyer flask or beaker under a calibrated burette attached to a burette holder. The nozzle of the burette should go into the neck of the flask or beaker without touching the walls.
   • Make sure all equipment is previously cleaned, rinsed with deionized water, and dry.
3. Fill the flask and burette. Accurately measure a small amount of the unknown solution. When dissolved, the substance spreads evenly through the solvent, so the concentration of this small sample of the solution will be the same as that of the original solution. Fill your burette with a solution of a known concentration that will react with your solution. Make a note of the exact volume of the solution in the burette - subtract the final volume to find the total solution used in the reaction.
   • Pay attention: if the reaction between the solution in the burette and the solute in the flask shows no sign of a reaction, you will indicator in the flask. These are used in chemistry to provide a visual signal when a solution reaches the point of equivalence or the end point. Indicators are generally used for titrations examining acid-base and redox reactions, but there are several other indicators as well. Consult a chemistry textbook or look on the internet to find a suitable indicator for your reaction.
4. Begin the titration. Slowly add a solution from the burette (the "titrant") into the flask. Use a magnetic stir stick or a glass stir stick to gently mix the solution while the reaction is underway. If your solution is visibly reacting, you should see certain signs that a reaction is underway - change in color, bubbles, residue, etc. If you are using an indicator, you may see every drop coming through the burette into the rightly flask a color change.
   • If the reaction results in a change in the pH value or potential, you can add pH readers or a potentiometer to the flask to gauge the progress of the chemical reaction.
   • For a more accurate titration, monitor the pH or potential as above, and note each time how the reaction proceeds after adding a small amount of titrant. Plot the acidity of the solution or the potential versus the volume of added titrant. You will see sharp changes in the slope of the curve at the equivalence points of the response.
5. Slow down your titration. As your chemical reaction nears the end point, slow the titration to a dropwise progression. If you are using an indicator, you may notice that the color flashes last longer. Now continue to titrate as slowly as possible until you can determine the exact droplet that will cause your reaction to reach the endpoint. In the case of an indicator, you generally look for the earliest possible sustained color change in the response.
   • Record the final volume in your burette. Subtracting this from the starting volume in the burette, you can find the exact volume of the titrant you used.
6. Calculate the amount of solute in your solution. Use the chemical equation for the reaction between your titrant and the solution to find the number of moles of solute in your flask. Once you find the number of moles of solute, you can simply divide it by the volume of the solution in the flask to find the molarity of the solution, or convert the number of moles to grams and divide by the volume of the solution. solution, to get the concentration in g / L. This requires a little basic knowledge of stoichiometry.
   • For example, suppose we used 25 ml of 0.5 M NaOH in titrating a solution of HCl in water to the equivalence point. The HCl solution had a volume of 60 ml for the titration. How many moles of HCl are there in our solution?
   • To get started, let's take a look at the chemical equation for the reaction of NaOH and HCl: NaOH + HCl> H₂O + NaCl
   • In this case, 1 molecule of NaOH reacts with 1 molecule of HCl with the products water and NaCl. So because you added just enough NaOH to neutralize all of the HCl, the number of moles of NaOH consumed in the reaction will be equal to the number of moles of HCl in the flask.
   • So let's find out what is the amount of NaOH in moles. 25 ml NaOH = 0.025 L NaOH x (0.5 mol NaOH / 1 L) = 0.0125 mole NaOH.
   • Since we deduced from the reaction equation that the number of moles of NaOH consumed in the reaction = the number of moles of HCl in the solution, we now know that there are 0.0125 moles of HCl in the solution.
7. Calculate the concentration of your solution. Now that you know the amount of solute in your solution, it is easy to find the concentration in terms of molarity. Simply divide the number of moles of solute in your solution by the volume of your solution sample (not the volume of the larger amount you took the sample from.) The result is the molarity of your solution!
   • To find the molarity of the example above, divide the number of moles of HCl by the volume in the flask. 0.0125 mole HCl x (1 / 0.060 L) = 0.208 M HCl.
   • To convert molarity to g / L, ppm, or percentage of composition, convert moles of your solute to mass (using the molar mass of your solute.) For ppm and percentage of the compound, you also need to convert the volume of your solution to mass (using a conversion factor such as density, or simply by weighing it), then multiply the result by 10 or 10, respectively.

Part 3 of 3: Determining the salinity in an aquarium

1. Take a water sample from your tank. Record the volume accurately. If possible, measure the volume in SI units such as mL - these are easy to convert to L.
   • In this example we test the water in the aquarium for salinity, the concentration of salt (NaCl) in the water. Suppose we take a water sample for this purpose 3 mL from the aquarium and then set the final answer to be given g / L.
2. Titrate the water sample. Select a titrant that produces a clearly visible reaction in the solute. In this case we use a solution of 0.25 M AgNO₃ (silver nitrate), a compound that produces an insoluble chlorine salt when it reacts with NaCl in the following reaction: AgNO₃ + NaCl> NaNO₃ + AgCl. The salt (AgCl) will be visible as a cloudy white residue that floats and can be separated from the solution.
   • Titrate the silver nitrate from a burette or small injection needle into the aquarium sample until the solution becomes cloudy. With such a small sample it is important to exactly determine how much silver nitrate you have added - study each drop carefully.
3. Continue until the reaction ends. When the silver nitrate stops clouding the solution, you can note the number of ml added. Titrate the AgNO3 very slow and observe the solution closely, especially as the end point approaches.
   • Assume that there are 3mL of the 0.25 M AgNO₃ was necessary for the reaction to come to an end and the water did not cloud further.
4. Determine the number of moles of the titrant. This step is easy - multiply the volume of the titrant you added by the molarity. This will give you the number of moles of titrant used.
   • 3 mL x 0.25 M = 0.003 L x (.25 moles AgNO₃/ 1 L) = 0.000075 mole AgNO₃.
5. Determine the number of moles of your solute. Use the reaction equation to convert the number of moles of AgNO₃ to moles of NaCl. The reaction equation is: AgNO₃ + NaCl> NaNO₃ + AgCl. Because 1 mole AgNO₃ reacts with 1 mole of NaCl, we now know that the number of moles of NaCl in our solution = the number of moles of AgNO₃ that is added: 0.000075 mol.
   • In this case: 1 mole of AgNO₃ reacts with 1 mol of NaCl. But if 1 mole of titrant reacts with 2 moles of our solute, then we would multiply the number of moles of our titrant by 2 to get the number of moles of our solute.
   • In contrast, if 2 moles of our titrant reacts with 1 mole of our solute, then we divide the number of moles of the titrant by two.
   • These rules correspond proportionally to 3 moles of titrant and 1 mole of solute, 4 moles of titrant and 1 mole of solute, etc. as well as 1 mole of titrant and 3 moles of solute, 1 mole of titrant and 4 moles of solute, etc.
6. Convert your solute number of moles to grams. To do this you will need to calculate the molar mass of the solute and multiply it by the number of moles of your solute. To find the molar mass of NaCl, use the periodic table to find and add the atomic weight of salt (Na) and Chloride (Cl).
   • MM Na = 22,990. MM Cl = 35,453.
   • 22,990 + 35,453 = 58.443 g / mol
   • 0.000075 mole NaCl x 58.442 g / mole = 0.00438 mole NaCl.
   • Pay attention: If there is more than one of one kind of molecule in an atom, you have to add the molar mass of that atom several times. For example, if you are the molar mass of AgNO₃, you would add the mass of oxygen three times because there are three oxygen atoms in the molecule.
7. Calculate the final concentration. We have the mass of our solute in grams and we know the volume of the test solution. All we have to do now is divide: 0.00438 g NaCl / 0.003 L = 1.46 g NaCl / L
   • The salinity of sea water is about 35 g NaCl / L. Our aquarium isn't nearly salty enough for marine fish.

Tips

• Although the solute and solvent may exist in different states (solid, liquid, or gas) when separated, the solution formed when the substance dissolves will be in the same state as the solvent state.
• Ag + 2 HNO3 → AgNO3 + NO2 + H2O
• Only use clear plastic or glass.
• Here is an example video: [1]

Warnings

• Store the AgNO3 solution in a closed, dark bottle. It is sensitive to light.
• Be careful when working with strong acids or bases. Make sure there is sufficient fresh air in the room.
• Wear safety glasses and gloves.
• If you want to get the silver back, note the following: Cu (s) + 2 AgNO3 (aq) → Cu (NO3) 2 + 2 Ag (s) Remember that (s) means solid.