6 Mixtures And Solutions 2008

SOLUTIONS So far we have been talking about substances and their properties. We have also mentioned solubility properties and related them to structure. Now we will take a closer look to mixtures in particular to solutions. Almost every material is a mixture: rocks (oxides and silicates, and so on), air (nitrogen and oxygen and others) sea water (water and salts), a cell (water, carbohydrates, proteins, salts, etc.). The importance of solutions stands on the fact that many chemical processes occur only or just go better and faster in solution. On the other hand, manipulating solutions is many times easier than working with solids. To measure a volume or pumping a liquid is far less messy than weighing or moving a solid. MIXTURES A SUBSTANCE is a material or system with a constant composition that cannot be changed. This means that the substance is the same no matter where it is found. NaCl, H2O, Ne, CO2, and O2 are all substances, because their composition will be the same no matter where you find them. SO2 and SO3 are different substances because they show a different composition and thus have very different properties. All elements and all compounds are defined as substances. A MIXTURE is a system formed by two or more substances that are not chemically united and do not keep fixed proportions to each other. In a compound, the atoms are covalently or ionically bonded and for a given substance the proportion in which they are combined is always the same (This is known as Proust’s law). The properties of a compound are absolutely different to those of the elements that form them. On the other hand, the composition of a mixture can be changed gradually and its properties are similar to those of the substances in it. For example, hydrogen and oxygen are gases, the second one being eight times denser. Mixtures of hydrogen and oxygen can be formed in any proportion and the properties of the mixtures formed (e. g. its density) will be between the properties of the pure gases. Instead, if the two gases combine forming water the particles show absolutely different polarity to that of the elements and water does not resemble them in any way! Most natural systems are mixtures. MIXTURES PURE SUBSTANCES A mixture has no fixed proportion between A pure substance has a constant composition its components with fixed ratios of elements. A mixture can be physically separated intoA compound can be decomposed (broken) pure substances by physical methods by drastic chemical methods Although it is almost physically impossible Just about everything that you can think of to isolate absolutely pure substances, a is probably a mixture. Even the purest of substance is said to be pure if no impurities materials still contain other substances as can be detected using the best available impurities. analytical techniques. Mixtures may exhibit a changing set of Physical properties such as boiling point or physical properties. melting point of pure substances are invariant. For example, pure water boils at For example, mixture of alcohol and water 100 degrees C boils over a range of temperatures.

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HOMOGENEOUS AND HETEROGENEOUS MIXTURES Some other concepts are needed to define what a solution is properly. INTENSIVE PROPERTIES are those properties such as density, boiling point, melting point etc. that do no depend on the amount of substance considered. Weight, volume, mass, are not intensive but extensive properties. PHASE: A phase is any observable region of a material that has its own set of intensive properties. A mixture of sand and sodium chloride has two phases, one of them having the properties of salt, the other one those of sand. In a chocolate chip cookie the dough and the chips have different properties. A spoonful of salt stirred in water forms just one phase because we cannot observe regions with different properties. HETEROGENEOUS SYSTEMS are made up of more than one phase and they can be separated mechanically. The two phases in a sand and water mixture can be separated by filtration and in a salad dressing oil and vinegar by decantation (let stand until oil forms a layer above water and then suck out one of the layers). Notice that ice and water mixed together show two phases despite being the same pure substance. They can be separated just picking up the ice cubes! HOMOGENEOUS SYSTEMS instead, show just one phase. Elements like hydrogen, compounds like sugar, and solutions like salt water, are all considered homogeneous because they are uniform. Each region of a sample is identical to all other regions of the same sample. Some systems seem to be homogeneous to the naked eye but show at heterogeneous when seen through a microscope. Milk is one of such systems. The most powerful the microscope the more we have a chance to find a second phase; so where’s the limit? Chemists agree to consider as homogeneous a system that looks so (optically “empty” or “void”, no “other thing” seen) under an ultramicroscope. An ultramicroscope is a sideways lighted powerful microscope. Very small particles can be seen as tiny points through it because of the Tyndall effect (the same as when you see atmospheric dust when sun beams enter a dark room) These “limiting systems” are called COLLOIDS. Some chemists refer to them as colloidal dispersions although most as colloidal solutions. Albumen (egg white) “solutions” are colloids; gelatine is a colloid too. To separate the components of a homogeneous mixture we profit on the different physical properties of the components (solubility, boiling point etc.) One of the most powerful methods for the separation of complex mixtures is called chromatography.

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CHROMATOGRAPHY

ASSESSING PURITY Impure substances (mixtures in which components in low proportion have to be separated) melt or boil along a range of temperatures and do not have a definite melting or boiling point. Chemists take profit from this fact to decide whether they are dealing with a pure substance or not. To assess purity they usually verify that some physical intensive properties keep reasonably constant after two or three fractionating processes. So, if a solid sample shows the same sharp melting point after two successive recrystallisations it is considered reasonably pure. As we previously said, chromatography is used to assess purity. Gas chromatography allows chemists to detect traces of impurities. The sophisticated technology of the late twentieth century requires

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extremely pure materials and this has fostered the development of new preparative and analytic techniques. SOLUTIONS A solution is a mixture of two or more substances in a single phase. At least two substances must be mixed in order to have a solution. The substance in the smallest amount, the one that dissolves is called the SOLUTE. The substance in the larger amount that surrounds the solute’s particles is called the SOLVENT. In many instances water is the solvent. The gases, liquids, or solids dissolved in water are the solutes. The solvent is the continuous phase: you can move from any point of it to any other point without going through the second phase. The solute is discontinuous: to move from one point to another without going through the solvent is not always possible. Since solutions are mixtures, their compositions may vary over a very wide range. The concentrations may be expressed using a variety of measures. The non-specific terms concentrated and dilute are sometimes used. A concentrated solution has a relatively large (but non-specific) amount of solute dissolved in a solvent. A dilute solution has a smaller quantity of solute dissolved. The table below shows some different classes of solutions

TYPES OF SOLUTIONS Concentrations Solute Less than 50%

Solvent Examples More than 50%

liquid

liquid

alcohol - water

wine beer, vodka acetic acid / water - vinegar

solid

liquid

salt - water

saline (NaCl) solution sugar solution CaCO3 - hard water

gas

liquid

oxygen - water

CO2 - carbonated water NH3 - ammonia solution

gas

gas

gas

solid

hydrogen - platinum

liquid

solid

mercury - another metal dentist’s amalgam

solid

solid

oxygen - nitrogen

copper - tin

air catalyst for hydrogenations Alloy (bronze

SOLUBILITY A solution is said to be SATURATED when it is in equilibrium (no macroscopic changes are detected) with an excess of solute. A definite amount of a solvent can 4

usually dissolve a definite amount of a solute. In some cases (as water and alcohol) solubility is considered as “infinite”: you can add more and more alcohol and it will always dissolve. What is happening in this case is that the solute becomes the solvent and vice versa. The water molecules that originally surrounded the alcohol particles are now in a lesser proportion and they are in turn surrounded by alcohol molecules! We call these two liquids MISCIBLE LIQUIDS. SOLUBILITY is the maximum amount of solute that can be dissolved in a fixed amount of solvent at a given temperature. The solubility of a solute is the concentration of its saturated solution. The conventional reference for solubility is the number of grams of solute that can dissolve in 100 mL of solvent. These values are the amount of solute that will dissolve and form a saturated solution at the temperature listed. The solvent cannot dissolve more solvent at that temperature. The solubility can be increased if the temperature is increased. Clearly there are exceptions such as Ce2(SO4)3 Solubility curves show graphically how the solubility of a substance changes with temperature

A saturated KCl solution at 10oC will have 31 grams of KCl dissolved in 100 grams of water. If there are 40 grams of KCl are in the container, then there will be 9 grams of undissolved KCl remaining in the solid. Raising the temperature of the mixture to 30oC will increase the amount of dissolved KCl to 37 grams and there will be only 3 grams of solid undissolved. The entire 40 grams can be dissolved if the temperature is raised above 40oC. Cooling the hot 40oC solution will reverse the process. When the temperature decreased to 20oC the solubility will eventually be decreased to 34 gram KCl. There is a time delay before the extra 6 grams of dissolved KCl crystallizes. This solution is "supersaturated" and is a temporary condition. The "extra" solute will come out of solution when the randomly moving solute particles can form the crystal pattern of the 5

solid. A "seed" crystal is sometimes needed to provide the surface for solute particles to crystallize on and establish equilibrium. CONCENTRATION OF A SOLUTION The composition of a mixture can be stated in different ways. We usually call the concentration of a solution, the amount of solute dissolved in a given mass or volume of either solvent or total solution. A solution can be concentrated, if the solute is around 10 to 15 % or more of the total mixture, or diluted, in case that it is in a lesser proportion. Concentrated solutions are usually stocked in the lab or in a chemical factory, and they are diluted just before they are used. Concentration can be stated in different forms: Percentage mass to mass (% m / m): grams of solute dissolved in 100 grams of solution. Grams of solute in 100 cm3 of solvent: this is used mainly for solubility data. Percentage mass to volume (% m / v): grams of solute in 100 mL (cm3) of solution. Grams per Litre (g / L or g / dm3 or g.dm-3): grams of solute in one litre of solution Molarity (M or moles dm3): moles of solute per Litre of solution parts per million (p p m): how many parts (grams) are there in 1.000.000 parts of the mixture. Concentration can be stated in any of the forms shown above (or many others). Sometimes it is necessary to change from one form to another one. The density of the solution in g/cm3 must be known in case you must move from a mass in mass to a mass in volume system. WARNING! The concentration of a solute can be given as grams of solute / volume of solution. Do not mess it up with the density of the solution that is grams of solution / volume of solution. The first one tells you how many grams of solute are there in 1 cm3; the second one how much does one cm3 of solution weighs. Of course the question is WHAT DOES MOLE MEAN IN CHEMISTRY? Mr. Mole THE MOLE CONCEPT ATOMIC MASS UNIT AND THE RELATIVE MASSES OF ATOMS AND MOLECULES

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We have studied that the masses of protons and neutrons were sensibly equal to 1. But…Which was the unit used? Evidently it was not the gram but a unit called the “atomic mass unit” with no further explanation. When the mass of a chlorine atom is said to be 35 it means that it has a mass of 35 a.m.u. The Ar or relative atomic mass of chlorine was stated as 35,5 because it was found to be formed by a mixture of isotopes (mainly those with masses 35 and 37) in a given proportion. The current definition of the atomic mass unit is as follows: 1 a.m.u. = 1/12 of the mass of a 12C atom. For reasons beyond the scope of this paper this is preferred to the previous one: 1 a.m.u. is the mass of a hydrogen atom. The values are nevertheless almost the same. It should be stated that 1 a.m.u. = 1,67 x 10-24 grams so 1 gram = 1/1,67x10-24 = 6x1023 a.m.u. The mass of molecules (or formulae in the case of ionic compounds) can be calculated adding the masses of each of the particles in them. The relative molecular mass of H2O is then twice the mass of a hydrogen atom plus the mass of an oxygen atom: Ar (H) = 1 Ar (O) = 16 Mr(H2O) = (2 x 1) + (16 x 1) = 18 AVOGADRO’S NUMBER AND THE MOLE Shoes come in pairs (2 units), eggs are sold by the dozen (12 units), pencils in grosses (144 units) and paper sheets in reams (500 units). Atoms and molecules are manipulated in moles (6 x1023 units). The mole is one of the fundamental concepts in chemistry One mole of a substance is the amount of matter containing 6x1023 particles of that substance. This number (six hundred thousand million million million units) is called Avogadro’s number so we can say that a mole of A is one Avogadro’s number of particles of A. One mole of any substance contains the same number of particles of one mole of any other chemical species. WEIGHING FOR COUNTING How much does one mole of hydrogen molecules weigh? To answer this question we have to multiply the mass in grams of a hydrogen molecule times Avogadro’s number. The mass of a hydrogen molecule (H2) is twice the mass of an atom (Ar(H) = 1) so: Mass of 1 mole (H2) = 2 x 1,67 x 10-24 x 6x 1023 = 2 grams (see previous page) How much does one mole of water molecules weigh? Following the same reasoning Mass of 1 mole (H2O) = 18 x 1,67 x 10-24 x 6 x 1023 = 18 grams It can be clearly seen that the mass in grams of one mole of any substance is numerically equal to its relative atomic (or molecular or formula) mass. To carry out a chemical reaction we must match the number of particles of the different reactants according to the stoichiometric ratios. We cannot count trillions of molecules but we can weigh the mass of the number of particles needed for the reaction.

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Solutions of known molarity are made by weighing out a definite number of grams of solute. The solute is placed in a container that has a calibrated volume, like a volumetric flask. A solvent is added to the flask to bring the solution volume to the calibration mark.

Can you calculate the molarity?

DILUTION Concentrated solutions can be mixed with solvent to make weaker or dilute solutions. This is the kind of thing people do everyday with consumer products like fruit juice. Some concentrated solutions are used as "stock" solutions. Weaker solutions are typically used but the concentrated solutions require less storage space. In dilutions the amount of solvent is increased, but the amount of solute is kept constant. The result is a decreased concentration, but a greater volume. The idea is that the volumes may change but the number of moles does not. This means that the original number of moles and the final number of moles are the same. number moles original = number of moles final . If M stands for molarity and V for volume Moriginal x Voriginal = Mfinal x Vfinal ELECTROLYTES Electrolytes are substances that form solutions that conduct electricity. Pure water does not conduct electricity. Water molecules do not dissociate significantly to form charge carriers (ions). Electrolytes dissolve in water or some other solvent and form ions. Sodium chloride is an electrolyte. When NaCl dissolves in water the crystal is separated into Na1+ and Cl1- ions. The ions are surrounded by solvent molecules. The clumps of solvent molecules and ions are mobile in the solution. Strong electrolytes are substances that convert completely to ions when they dissolve. They are also said to be 100% ionized. The solubility of ionic compounds limits their

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ability to conduct. Some strong electrolytes are covalent compounds like HNO3 or HCl. These compounds also ionize 100% in aqueous solutions.

Weak electrolytes dissolve and only a small fraction of the formula units dissociate to form ions. Solutions of weak electrolytes are poor conductors of electricity. Acetic acid is a weak electrolyte. When 100 molecules of acetic acid dissolve in water 96 of the molecules remain intact and 4 of the molecules ionize. This means that the acetic acid is only 4% ionized or dissociated.

The hydrogen in the -O-H group is acidic. The two oxygens bonded to the same carbon atom attract electrons form the carbon and neighbouring atoms. This draws electrons from the O-H bond and, in turn makes the O-H bond weaker than normal. But the O-H bond is still strong enough to remain intact 96 % of the time. Nonelectrolytes dissolve and all of the formula units remain intact and none of them dissociate to form ions. Solutions of weak electrolytes are poor conductors of electricity. Ethyl alcohol and methyl alcohol are nonelectrolytes. None of the CH3OH or CH3CH2OH molecules dissociate appreciably to form ions. A solution of hexane and dodecane will not conduct electricity because there are no ions formed by the solute and solvent. SOME TIPS FOR SOLVING PROBLEMS • • • • •

The mass of the solution equals the sum of both the mass of the solute and the mass of the solvent Msolution = Msolute + Msolvent The % m / v can be simply calculated the % m / m times the density of the solution. % m / v = (% m / m) x δ 3 The concentration in g / dm is just the % m / v times 10. g / dm3 = (% m / v) x 10 Molarity is calculated dividing the concentration in g / dm 3 into the Mr of the solute. M = g / dm3 : Mr(solute) All this can be summarised as M = [(% m / m) x δ x 10] / Mr(solute)

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PROBLEMS 1- How would you separate a mixture containing: a- water and oil b- water and salt c- water and clay d- water and alcohol e- fourteen amino acids 2- The solubility of copper (II) sulphate depends on temperature as shown in the table Temperature (º C) 0 1 2 3 4 5 6 70 a- Plot the 0 0 0 0 0 0 solubility Solubility (g /100g sv) 1 1 2 2 2 3 4 47 curve 4 7 1 4 9 4 0 b- Calculate the mass of copper (II) sulphate that can be dissolved in 350 g of water at 70º C c- Explain what will happen is the solution is chilled down to 5º C. d- If the crystals are filtered at this temperature calculate the loss of copper (II) sulphate in case the filtrate was discarded. 3- The following table shows the solubility of sodium chloride in water at different temperatures. Temperature (º C) 0 20 40 60 80 100 Solubility (g in 100 g sv) 35, 36, 36, 37, 38, 39,0 5 0 5 0 0 Calculate the minimum amount of water to dissolve 200 g of a sample containing 90 % of copper (II) sulphate and 10 % of sodium chloride at 70ºC and the purity of the copper (II) sulphate crystals obtained by cooling at 5º C. (Plot the data given) 4- Calculate the relative formula (or molecular) mass Mr for the following compounds. Find the relative atomic masses Ar in the periodic table. Round them out to a whole number except for chlorine that is usually considered to be 35,5. sulphuric acid - nitric acid - hydrochloric acid - sodium hydroxide glucose (C6H12O6) - hydrated copper(II) sulphate (CuSO4 · 5 H2O) - aspirin ( C9H8O3) 5- Calculate how many moles and how many molecules of each of the substances of exercise (1) will be in a 100 g mass of each of them. 6- 250 g of hydrated copper (II) sulphate are dissolved in 700 g of water. Calculate the concentration of this solution as % m/m. If the density of the solution is 1,22 g / cm3 express the concentration as % m / v, g / L and molarity . 7- How many positive ions are there in 200 g of the solution prepared in (3)? 8- Which is the molarity of pure water? 10

9- 200 g of the solution in exercise (3) are added with 400 g of water and stirred to homogenise. The volume is then adjusted to 1.000 mL. 20 mL of this dilution are evaporated to dryness in as evaporating dish. Calculate the mass of hydrated copper (II) sulphate that remain in the evaporating dish. 10- How many mL of a 2M solution of HCl must be measured to prepare 0,5 L of a 0,4 M solution? 11- To improve your skill repeat exercises (3) and (4) for methanol (CH4O), considering that the density of the solution obtained is 0,9653 g / cm3. 12- How many ions and how many moles of ions are formed when 6 g of ethanoic acid (C2H4O2) are dissolved in water? Just 4 % of the molecules of ethanoic acid are ionised. Only one hydrogen ion is formed per molecule.

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