Topic 15 Topic 15 Topic 15: Solutions Table of Contents Topic 15 Topic 15 Basic Concepts Additional...

Topic 15Topic 15

Topic 15: Solutions

Table of ContentsTable of ContentsTopic 15Topic 15

Basic Concepts

Additional Concepts

• Because water is so much a part of life, its properties are easy to take for granted.

Water: The Molecular View

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• If you step back a bit and examine water scientifically, you will find that it is unusual among the compounds found on Earth.

• Water is most often thought of as a liquid.

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• However, solid water, called ice, and gaseous water, called steam or water vapor, also exist in large quantities on Earth.

Water: The Molecular View

• Water is the only substance on Earth that exists in large quantities in all three states.

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• The arrangement of electrons about the central oxygen in the water molecule relates to its three-dimensional geometry.

Geometry of the Water Molecule

• There is a large electronegativity difference between the covalently bonded hydrogen and oxygen.

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• Therefore, the electron pair is shared unequally.

Geometry of the Water Molecule

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• Because of the molecule’s bent structure, thepoles of positive and negative charge in the two bonds do not cancel, and the water molecule as a whole is polar.

• A group of water molecules will orient at the molecular scale because of electrical forces.

Attractions and Order

• The opposites attract and create order among molecules.

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• This effect is especially great at low temperatures.

• Notice that the oxygen on one water molecule is attracted to the hydrogen atoms on other water molecules.

Modeling Water: Hydrogen Bonding

• The connections between the molecules are not full covalent bonds, but they are still fairly strong.

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• The formation of such a connection between the hydrogen atoms on one molecule and a highly electronegative atom on another is called hydrogen bonding.

Modeling Water: Hydrogen Bonding

• Atoms that are electronegative enough to cause hydrogen bonding include oxygen, fluorine, and nitrogen.

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• Notice that water molecules have both intermolecular forces (hydrogen bonds, dashed lines) and intramolecular forces (covalent bonds, solid lines).

Hydrogen Bonding Versus Covalent Bonding in Water

• The oxygens are highly electronegative and strongly attract hydrogen’s electrons.

• As a result, the hydrogen nucleus, a proton, is partially exposed.

• Oxygen-hydrogen groups in molecules tend to promote intermolecular hydrogen bonding.

Water: The Hydrogen Bonding Champion

• In pure water, each water molecule may form hydrogen bonds with four other water molecules.

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• Any molecule that contains O—H bonds has the potential to form hydrogen bonds.

Water: The Hydrogen Bonding Champion

• Oxygen is sufficiently electronegative to attract hydrogen’s single electron strongly so that the hydrogen atom almost becomes an exposed proton.

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States of Water

• The intermolecular hydrogen bonds hold the water molecules together strongly enough that they cannot readily escape into the gaseous state at ordinary temperatures.

• Water occurs primarily in the liquid and solidstates on Earth, rather than as a gas.

States of Water

• That is why water has such a high boilingpoint for such a small molecule, 100°C.

• You know that if you drop an ice cube into a glass of water, the ice floats.

Ice Floats

• You also know this means that the density of the solid water is less than that of liquid water.

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• As water cools from 60°C, its volume decreases and its density increases.

• The water molecules move less rapidly, and they are able to be drawn closer together by dipole-dipole attractions.

Ice Floats

• The volume of the water decreases because the molecules pull together.

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• Meanwhile, the mass of water stays the same, so density increases.

• You can account for this if you know what is happening to the molecular arrangement.

Ice Floats

• Below 4°C, the water molecules are beginning to approach the solid state, which is highly organized.

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Ice Floats

• The water molecules begin to form the open arrangement.

• The arrangement results from hydrogen bonding and is the most stable structure for the molecules in or near the solid state.

Surface Tension

• Have you ever watched water drip from a faucet? Each drop is composed of an enormous number of water molecules— roughly 2 x 1021.

• The observation that this large number of molecules can hold together as a unit to form a single drop is more evidence for the presence of intermolecular forces.

Surface Tension

• A water molecule forms a drop because of surface tension, which is the force needed to overcome intermolecular attractions and break through the surface of a liquid or spread the liquid out.

Surface Tension

• The higher the surface tension of a liquid is, the more resistant the liquid is to having its surface broken.

• The net inward force makes the surface of the drop contract and seem to toughen, behaving like a sort of skin.

Capillarity

• These examples of the rising of liquids in narrow tubes are illustrations of capillarity, or, as it is sometimes called, capillary action.

• Capillarity results from the competition between the interparticle attractive forces between the molecules of liquid and the attractive forces between the liquid and the tube that contains it.

Capillarity

• In the case of a glass tube, water molecules can form hydrogen bonds to the oxygen atoms in the silicon dioxide that makes up the glass.

• This attractive force between the water and the SiO2 draws the water up the walls of the tube.

• Because the water molecules are also attracted to each other, more water rises upward.

Capillarity

• If the tube is narrow, like the capillary tubes used in blood tests, the liquid will be drawn high into the tube because nearly all the molecules are close to the walls and are thus strongly attracted to them.

Specific Heat

• Specific heat measures the amount of heat, in joules, needed to raise the temperature of 1 g of substance by 1°C.

• This means that water must absorb or release more heat for its temperature to change by one Celsius degree than any of the other substances.

Specific Heat

• The specific heat of water is 4.18J/g°C. (The unit is read “joules per gram per degree Celsius.”)

• In order to raise the temperature of 1 g of water by 1°C, you must add 4.18 J of heat.

• On the other hand, you must remove 4.18 J of heat from a 1-g sample of water to lower its temperature by 1°C.

Water Evaporation/Condensation

• Recall that vaporization is the change in state of liquid to gas, and that the amount of heat required to vaporize a quantity of a liquid is called the heat of vaporization.

• During the vaporization of a liquid, interparticle forces must be overcome.

• Therefore, vaporization of a liquid is an endothermic (energy-absorbing) process.

• Condensation, the formation of a liquid from a gas, is an exothermic (energy-releasing) process.

• Evaporation is vaporization from the surface of a liquid.

• Evaporation of water is one of the mechanisms by which your body regulates its temperature.

Water: The Super Solvent

• Most of the water on Earth is not pure, but rather is present in solutions.

• Water is difficult to keep pure because it is an excellent solvent for a variety of solutes.

Water: The Super Solvent

• Water is such a versatile solvent that it is sometimes called the universal solvent.

• Its ability to act as a solvent is one of its most important physical properties.

• As you will see, it is again the attraction of water molecules for other molecules, as well as for one another, that accounts for these solvent properties.

Water Dissolves Many Ionic Substances

• Salt, like a great many ionic compounds, is soluble in water.

• The salt solution is also an excellent conductor of electricity.

• This high level of electrical conductivity is always observed when ionic compounds dissolve to a significant extent in water.

Water Dissolves Many Ionic Substances

• Your model of water and its interactions explains why salt and many other ionic compounds dissolve in water and why the solutions conduct electricity.

A Model of the Dissolving of NaCl

• Remember that ionic solids are composed of a three-dimensional network of positive and negative ions, which form strong ionic bonds.

• The process by which the charged particles in an ionic solid separate from one another is called dissociation.

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• You can represent the process of dissolving and dissociation in shorthand fashion by the following equation.

Water Dissolves Many Covalent Substances

• Water is not only good at dissolving ionic substances. It also is a good solvent for many covalent compounds.

• Consider the covalent substance sucrose, commonly known as table sugar, as an example.

• You have probably observed that this substance, with the formula C12H22O11, dissolves in water. In fact, it is highly soluble.

• It’s possible to dissolve almost 200 g of sugar in 100 mL of water.

• Take a look at the molecular structure of a sucrose molecule.

• Notice that the structure has a number of O—H bonds.

Solutions: Basic ConceptsSolutions: Basic ConceptsSolutions: Basic ConceptsSolutions: Basic ConceptsTopic 15Topic 15 Water Dissolves Many Covalent

Substances

• As you learned earlier, if a molecule containsO—H bonds, it will tend to be polar and it can form hydrogen bonds.

Like Dissolves Like

• Although water dissolves an enormous variety of substances, both ionic and covalent, it does not dissolve everything.

• The phrase that scientists often use when predicting solubility is “like dissolves like.”

• The expression means that dissolving occurs when similarities exist between the solvent and the solute.

Concentrated Versus Dilute

• Chemists never apply the terms strong and weak to solution concentrations.

• As you’ll see in the next chapter, these terms are used in chemistry to describe the chemical behavior of acids and bases.

• Instead, use the terms concentrated and dilute.

Unsaturated Versus Saturated

• Another way of providing information about solution composition is to express how much solute is present relative to the maximum amount the solution could hold.

• If the amount of solute dissolved is less than the maximum that could be dissolved, the solution is called an unsaturated solution.

• Such a solution, which holds the maximum amount of solute per amount of the solution under the given conditions, is called a saturated solution.

• An interesting third category of solution is called a supersaturated solution.

• Such solutions contain more solute than the usual maximum amount and are unstable.

• They cannot permanently hold the excess solute in solution and may release it suddenly.

• Supersaturated solutions, as you might imagine, have to be prepared carefully.

• Generally, this is done by dissolving a solute in the solution at an elevated temperature, at which solubility is higher than at room temperature, and then slowly cooling the solution.

Effect of Temperature on Solubility

• Temperature has a significant effect on solubility for most solutes.

• The solubilities of some solutes, such as sodium nitrate and potassium nitrate, increase dramatically with increasing temperature.

Effect of Temperature on Solubility

• Other solutes, like NaCl and KCl, show only slight increases in solubility with increasing temperatures.

• A few solutes, like cerium(III) sulfate, Ce2(SO4)3, decrease in solubility as temperature increases.

Basic Assessment QuestionsBasic Assessment Questions

Question 1

Name the three compounds that form hydrogen bonds.

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Answer

Sample answer: water, alcohols, biological molecules containing O—H bonds.

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Question 2

Define capillarity.

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Answer

Competitive forces between molecules of a liquid and the tube that contains the liquid.

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Question 3

Describe why water is sometimes called the universal solvent.

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Answer

Water is a versatile solvent because of its attraction to other molecules and its polarity.

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Solutions: Additional ConceptsSolutions: Additional Concepts Solutions: Additional ConceptsSolutions: Additional Concepts Topic 15Topic 15

Additional Concepts

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Molarity• Concentration units can vary greatly. • They express a ratio that compares an

amount of the solute with an amount of the solution or the solvent.

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• For chemistry applications, the concentration term molarity is generally the most useful.

• Molarity is defined as the number of moles of solute per liter of solution.

• Molarity = moles of solute/liter of solution

Molarity

• Note that the volume is the total solution volume that results, not the volume of solvent alone.

• Suppose you need 1.0 L of the salt solution mentioned above.

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MolarityTopic 15Topic 15

• In other words, it must have a molarity of 0.15.

• In order to be at the same concentration as the salt in the patient’s blood, it needs to have a concentration of 0.15 moles of sodium chloride per liter of solution.

• To save space, you refer to the solution as 0.15M NaCl, where the M stands for “moles/liter” and represents the word molar.

• Thus, you need 1.0 L of a 0.15-molar solution of NaCl. How are you going to prepare it?

• Assuming you’re making an aqueous solution, you need to know only three things when working quantitatively: the concentration, the amount of solute, and the total volume of solution needed.

Preparing 1 L of an NaCl Solution

• How would you prepare 1.0 L of a 0.15M sodium chloride solution?

• First, determine the mass of NaCl to add to a 1.0-L container.

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• The 0.15M solution must contain 0.15 moles of NaCl per liter of solution.

Preparing 1 L of an NaCl SolutionTopic 15Topic 15

• The proper setup, showing the conversion factors, is as follows.

Preparing 1 L of an NaCl Solution

• Then carry out cancellations and calculate the answer.

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Preparing 1 L of an NaCl SolutionTopic 15Topic 15

• The result means you need to measure 8.8 g of NaCl, add some water to dissolve it, and then add enough additional water to bring the total volume of the solution to 1.0 L.

Preparing a Different Volume of a Glucose Solution

• How would you prepare 5.0L of a 1.5M solution of glucose, C6H12O6?

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• You need to determine the number of grams of glucose to add to a 5.0-L container.

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• The 1.5M solution must contain 1.5 mol of glucose per liter of solution.

• The proper setup, showing the conversion factors, is as follows.

• Cancel units and carry out the calculation.

Topic 15Topic 15 Preparing a Different Volume of a

Glucose Solution

Solutions: Additional ConceptsSolutions: Additional Concepts Solutions: Additional ConceptsSolutions: Additional Concepts Topic 15Topic 15

• The mass of glucose required is 1400 g. • Weigh this mass, add it to a 5.0-L container,

add enough water to dissolve the glucose, and fill with water to the 5.0-L mark.

Calculating Molarity

• You add 32.0 g of potassium chloride to a container and add enough water to bring the total solution volume to 955 mL. What is the molarity of this solution?

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• You are given that there are 32.0 g of solute per 955 mL of solution, so this relationship can be expressed in fraction form with the volume in the denominator.

• Therefore, the initial part of the setup is as follows.

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Calculating Molarity• Determine that the molar mass of KCl is

74.6 g/mol by adding the atomic masses of K and Cl and applying the unit grams/mole to the sum.

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• The conversion factor that must be used to convert from grams to moles of KCl is 1mol KCl/74.6 g KCl.

• Next, to convert milliliters to liters, given that there are 1000 mL solution/L solution, use that conversion factor in the setup.

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Calculating Molarity• Cancel units and carry out the calculation,

using the setup just developed.

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Freezing-Point Depression

• A solution always has a lower freezing point than the corresponding pure solvent.

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• If you are interested only in aqueous solutions, this means that any aqueous solution will have a freezing point lower than 0°C.

• The amount that the freezing point is depressed relative to 0°C depends only upon the concentration of the solute.

Boiling-Point Elevation

• You have just learned that the freezing point of a solution is lower than the freezing point of the pure solvent.

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• It turns out that the boiling point of a solution is higher than the boiling point of the pure solvent.

Boiling-Point Elevation

• For aqueous solutions this means that the solution boiling point will be greater than 100°C, assuming standard atmospheric pressure.

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• The solute must also be nonvolatile; that is not able to evaporate readily.

Osmosis

• The flow of solvent molecules through a selectively permeable membrane, driven by concentration difference, is called osmosis.

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OsmosisTopic 15Topic 15

• As osmosis begins, water molecules diffuse more rapidly from the water into the sucrose solution than they diffuse from the sucrose solution into the water.

OsmosisTopic 15Topic 15

• As a result, the sucrose solution gains water, becomes more dilute, and the volume rises.

Osmosis• The increasing height of the

sucrose solution exerts a pressure that opposes the diffusion of water molecules from left to right.

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• Eventually, the pressure becomes high enough that the rates of diffusion of water in both directions become the same.

Osmosis• Adding extra pressure to the

side with the sucrose solution can cause the water diffusion to move in the opposite direction, forcing water out of the solution and producing pure water.

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• This process, called reverse osmosis, can be used to purify water.

Colloids

• Sometimes, mixtures are partway between true solutions and heterogeneous mixtures.

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• Such mixtures, called colloids, contain particles that are evenly distributed through a dispersing medium, and remain distributed over time rather than settling out.

• The major difference between a colloid and a solution is the size of the solute particles.

Colloids

• Colloid particles are generally clumps that are ten to 100 times larger than typical ions or molecules dissolved in solutions.

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• Because of their relatively large particle size, colloids play important roles in a variety of processes.

• Some biological molecules, such as proteins, are large enough that their behavior is often best understood using a colloid model.

Molality

• The molality (m) of a solution is equal to the number of moles of solute per kilogram of solvent.

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Calculating Molality

• What is the molality of a solution that contains 16.3 g of potassium chloride dissolved in 845 g of water?

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• Convert the mass of solute to moles.

Calculating MolalityTopic 15Topic 15

• The solvent mass, 845 g, must be expressed in kilograms.

Calculating Molality

• Substitute the known values into the equation for molality and solve.

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Additional Assessment QuestionsAdditional Assessment Questions

A solution is made by dissolving 17.0 g of lithium iodide (LiI) in enough water to make 387 mL of solution. What is the molarity of the solution?

Question 1 Topic 15Topic 15

0.328M

AnswerTopic 15Topic 15

Calculate the molarity of a water solution of CaCl2, given that 5.04 L of the solution contains 612 g of CaCl2.

What is the molality of the solution formed by mixing 104 g of silver nitrate (AgNO3) with 1.75 kg of water?

0.350m

Suppose that 5.25 g of sulfur (S8) is dissolved in 682 g of the liquid solvent carbon disulfide (CS2). What is the molality of the sulfur solution?

0.0300m

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