1. Unit 2
Water and Solutions

2. Water and Heterogeneous systems
Objectives:
Explain surface tension
Describe the structure of ice and liquid water.
Distinguish between a suspension and a solution
Identify the characteristics of a colloid.

3. Water in the liquid state:
Water is a polar molecule. d+
H2O d-

4. Hydrogen bonding: the negative end of the water molecule is attracted to the positive ends of other molecules- resulting in hydrogen bonding. O H
Hydrogen
bond d- d+

6. Water molecules “stick”
together to create
surface tension
to support light
weight objects.
It also tends to hold a drop
of a liquid in a spherical
shape.
Surfactants interfere with hydrogen bonding reducing surface tension.
This attraction between water molecules slows the tendency of water evaporation.
What_s_So_Special_About_Water_.asf
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

7. Suspensions: a mixture from which particles settle out upon standing.
Have very large particles
Can be filtered
Ex. Muddy water
some medicines (antibiotics)
Suspensions.asf

8. Colloids Have medium sized particles Can not be filtered.
Many colloids are cloudy or milky in appearance.
Show Tyndall effect (the path of light is visible)
Colloids.asf

9. Types of colloids Colloid type Dispersion medium Dispersed substance
Examples
Aerosol
Gas
Liquid
Solid
Solid emulsion

10. Types of colloids Colloid type Dispersion medium Dispersed substance
Examples
Aerosol
Gas
Liquid
Clouds, fog,aerosol spray
Solid
Smoke, dust
Foam
Shaving cream, whipped cream
Emulsion
Mayonaisse, milk
Sol
Paint, ink
Solid aerosol
Marshmallow, styrofoam
Solid emulsion
gel
Solid sol
pearl

11. Classwork: Complete the table on the last slide, giving examples for the different types of colloids.

12. Solutions Objectives: Distinguish between a solvent and a solute
Describe what happens in the solution process
Explain what electrolytes are.
Identify the factors that determine the rate at which a solute dissolves.
Identify the factors that determine the mass of solute that will dissolve in a given mass of solvent.

13. Definitions Solution = Solute + Solvent
Solution - homogeneous mixture
Solute - substance being dissolved
Solvent - present in greater amount
Aqueous (water)
Tincture (alcohol)
Amalgam (mercury)
Organic Polar or Non-polar
Dental filling
Solution = Solute + Solvent

14. Types of Solutions Solute Solvent Solution Gaseous Solutions
liquid
air (nitrogen, oxygen, argon gases)
humid air (water vapor in air)
Liquid Solutions
solid
carbonated drinks (CO2 in water)
vinegar (CH3COOH in water)
salt water (NaCl in water)
Solid Solutions
dental amalgam (Hg in Ag)
sterling silver (Cu in Ag), alloys
Solutions are not limited to gases and liquids; solid solutions also exist.
• Amalgams, which are usually solids, are solutions of metals in liquid mercury.
• Network solids are insoluble in all solvents with which they do not react chemically; covalent bonds that hold the network together are too strong to be broken and are much stronger than any combination of intermolecular interactions that might occur in solution.
• Most metals are insoluble in all solvents but do react with solutions such as aqueous acid or base to produce a solution; in these cases the metal undergoes a chemical transformation that cannot be reversed by removing the solvent.
Charles H.Corwin, Introductory Chemistry 2005, page 369

15. Solvation – the process of dissolving:
solute particles are surrounded by solvent particles
solute particles are separated and pulled into solution
Solution formation- The polarity or composition of the solute and the solvent will determine…Whether a substance will dissolve and how much
The interactions that determine the solubility of a substance in a liquid depend on the chemical nature of the solute whether it is ionic or molecular) rather than on its physical state (solid, liquid, or gas).
• Two examples:
1. Forming a solution of a molecular species in a liquid solvent
2. Formation of a solution of an ionic compound
Solutions of molecular substances in liquids
– London dispersion forces, dipole-dipole interactions, and hydrogen bonds that hold molecules to other molecules are weak.
– Energy is required to disrupt these interactions, and unless some of that energy is recovered in the formation of new, favorable solute-solvent interactions, the increase in entropy on solution formation is not enough for a solution to form.
– For solutions of gases in liquids, the energy required to separate the solute molecules is ignored (H2 = 0) because molecules are already separated – it is necessary to only consider the energy required to separate the solvent molecules (H1) and the energy released by new solute-solvent interactions (H3).
1. Nonpolar gases are most soluble in nonpolar solvents because H1 and H3 are both small and of similar magnitude.
2. Nonpolar gases are less soluble in polar solvents than in nonpolar solvents because H1 >> H3.
3. Solubilities of nonpolar gases in water increase as the molecular mass of the gas increases.
– All common organic liquids, whether polar or not, are miscible; the strengths of the intermolecular attractions are comparable, the enthalpy of solution is small, and the increase in entropy drives the formation of a solution.
– If predominant intermolecular interactions in two liquids are very different from one another, they may be immiscible, and when shaken with water, they form separate phases or layers separated by an interface.
– Only the three lightest alcohols are completely miscible with water; as the molecular mass of the alcohol increases, so does the proportion of hydrocarbon in the molecule, which leads to fewer favorable electrostatic interactions with water
Hydrophilic and hydrophobic solutes
– A solute can be classified as hydrophilic, meaning that there is an electrostatic attraction to water, or hydrophobic, meaning that it repels water.
1. Hydrophilic substance is polar and contains O–H or N–H groups that can form hydrogen bonds to water; tend to be very soluble in water and other strongly polar solvents
2. Hydrophobic substance may be polar but usually contains C–H bonds that do not interact favorably with water; essentially insoluble in water and soluble in nonpolar solvents
– The difference between hydrophilic and hydrophobic substances has substantial consequences in biological systems.
– Vitamins can be classified as either fat soluble or water soluble.
1. Fat-soluble vitamins (Vitamin A) are nonpolar, hydrophobic molecules and tend to be absorbed into fatty tissues and stored there.
2. Water-soluble vitamins (Vitamin C) are polar, hydrophilic molecules that circulate in the blood and intracellular fluids and are excreted from the body and must be replenished in the daily diet.

16. Dissolving of NaCl - - - O + + + + Cl- Na+ hydrated ions H
Timberlake, Chemistry 7th Edition, page 287

17. Electrolytes Electrolytes - solutions that carry an electric current
Electrolyte — any compound that can form ions when it dissolves in water
– When strong electrolytes dissolve, constituent ions dissociate completely, producing aqueous solutions that conduct electricity very well.
– When weak electrolytes dissolve, they produce relatively few ions in solution, and aqueous solutions, of weak electrolytes do not conduct electricity as well as solutions of strong electrolytes.
– Nonelectrolytes dissolve in water as neutral molecules and have no effect on conductivity.
strong electrolyte
weak electrolyte
nonelectrolyte
NaCl(aq) Na+ + Cl-
HF(aq) H+ + F-
Timberlake, Chemistry 7th Edition, page 290

18. Factors determining rate of solution. 1
Factors determining rate of solution Stirring (agitation) moves fresh solvent into contact with the solute. 2. Smaller pieces increases the amount of surface area of the solute.
3. Higher temperature makes the molecules of the solvent move faster; speeds up dissolving.
Higher Temperature ALSO Usually increases the amount that will dissolve (an exception is gases ).

19. Miscible means that two liquids can dissolve in each other
water and antifreeze
water and ethanol
Partially miscible- slightly
water and ether
Immiscible means they can’t
oil and vinegar

20. Solids tend to dissolve best when:
They are heated
They are stirred
Crushed into smaller particles
Gases tend to dissolve best when:
The solution is cold
The pressure is high
Thermal pollution may result from industry using water for cooling

21. Solubility- is the maximum amount of substance that will dissolve at a specific temperature. The units for solubility are: grams of solute/100 grams solvent
Saturated solution- Contains the maximum amount of solute dissolved NaCl = 36.0 g/100 mL water
Unsaturated solution- Can still dissolve more solute (for example 28.0 grams of NaCl/100 mL)
Supersaturated- solution that is holding (or dissolving) more than it theoretically can; Very unstable, can crystalize.

22. SUPERSATURATED SOLUTION
Solubility
UNSATURATED SOLUTION
more solute dissolves
SATURATED SOLUTION
no more solute dissolves
SUPERSATURATED SOLUTION
becomes unstable, crystals form
Maximum amount of a solute that can dissolve in a solvent at a specified temperature and pressure is its solubility.
– Solubility is expressed as the mass of solute per volume (g/L) or mass of solute per mass of solvent (g/g) or as the moles of solute per volume (mol/L).
– Solubility of a substance depends on energetic factors and on the temperature and, for gases, the pressure.
• A solution that contains the maximum possible amount of solute is saturated.
• If a solution contains less than the maximum amount of solute, it is unsaturated.
When a solution is saturated and excess solute is present, the rate of dissolution is equal to the rate of crystallization.
• Solubility increases with increasing temperature — a saturated solution that was prepared at a higher temperature contains more dissolved solute than it would contain at a lower temperature, when the solution is cooled, it can become supersaturated.
increasing concentration

24. Solubility vs. Temperature for Solids
a) 80g NaNO3 at 45C. The solution is ?
140
unsaturated KI
130
b) 100g NaNO3 at 45C. The solution is ?
120
gases
solids
saturated
NaNO3
110
c) 120g NaNO3 at 45C. The solution is ?
100
KNO3 90
supersaturated 80
d) How much more NaNO3 can you add to a solution with 40g of NaNO3 at 45C until it becomes saturated?
HCl
NH4Cl 70
Solubility (grams of solute/100 g H2O) 60
NH3
KCl
60 g 50
“Solubility Curves for Selected Solutes”
Description: This slide is a graph of solubility curves for 10 solutes. It shows the number of grams of solute that will dissolve in 100 grams of water over a temperature range of 0cC to 10 cC.
Basic Concepts
The maximum amount of solute that will dissolve at a given temperature in 100 grams of water is given by the solubility curve for that substance.
When the temperature of a saturated solution decreases, a precipitate forms.
Most solids become more soluble in water as temperature increases, whereas gases become less soluble as temperature increases.
Teaching Suggestions
Use this slide to teach students how to use solubility curves to determine the solubilities of various substances at different temperatures. Direct their attention to the dashed lines; these can be used to find the solubility of KClO3 at 50 cC (about 21 g per 100 g of H2O). Make sure students understand that a point on a solubility curve represents the maximum quantity of a particular solute that can be dissolved in a specified quantity of solvent or solution at a particular temperature. Point out that the solubility curve for a particular solute does not depend on whether other solutes also are present in the solution (unless there is a common-ion effect; this subject usually is covered at a later stage in a chemistry course).
Questions
Determine the solubilities (in water) of the following substance at the indicated temperatures: NH3 at 50 oC; KCl at 90 oC; and NaNO3 at 0 oC.
Which of the substances shown on the graph is most soluble in water at 20 oC? Which is lease soluble at that temperature? For which substance is the solubility lease affected by changes in temperature?
Why do you think solubilities are only shown between 0 oC and 100 oC?
In a flask, you heat a mixture of 120 grams of KClO3 and 300 grams of water until all of the KClO3 has just been dissolved. At what temperature does this occur? You then allow the flask to cool. When you examine it later, the temperature is 64 oC and you notice a white powder in the solution. What has happened? What is the mass of the white powder?
Compare the solubility curves for the gases HCl, NH3, and SO2) with the solubility curves for the solid solutes. What generalizations(s) can you make about the relationship between solubility and temperature?
According to an article in an engineering journal, there is a salt whose solubility in water increases as the water temperature increases from 0 oC to 65 oC. The salt’s solubility then decreases at temperatures above 65 oC, the article states. In your opinion, is such a salt likely to exist? Explain your answer. What could you do to verify the claims of the article?
e) What is the solubility of NH3 at 25C? 40 30
NaCl
58 g
KClO3 20
e) 100 g of water are saturated with NaNO3 at 30C. If the solution is heated to 60C, how much more can be dissolved? 10
SO2
20 g
LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World , 1996, page 517

25. Classwork: solubility graph handout

26. Concentration of Solutions
Objectives:
Solve problems involving the molarity of a solution.
Describe the effect of dilution on the total moles of solute in solution.
Define percent by volume and percent by mass solutions.

27. Concentration is...
a measure of the amount of solute dissolved in a given quantity of solvent
A concentrated solution has a large amount of solute
A dilute solution has a small amount of solute
These are qualitative descriptions
But, there are ways to express solution concentration quantitatively (NUMBERS!)

28. Concentrated vs. Dilute
Notice how dark the solutions appears.
Notice how light the solution appears.
Small amount of solute in a large amount of solvent.
Lots of solute, in a small amount of solvent.

29. Making solutions
Pour in a small amount of the solvent, maybe about one-half
Then add the pre-massed solute (and mix by swirling to dissolve it)
Carefully fill to final volume.

31. Molarity: a unit of concentration
Molarity = n (moles of solute)
V (liters of solution)
Abbreviated with a capital M, such as 6.0 M
Units M or mol/L
This is the most widely used concentration unit used in chemistry.

32. How many grams of sodium chloride, NaCl, do you need to prepare 250
How many grams of sodium chloride, NaCl, do you need to prepare 250.mL of a 0.5M NaCl solution? Molar mass NaCl= 58.5 g/mol

33. 2. What volume of a 1. 0M NaCl solution can you prepare if you have 45
2. What volume of a 1.0M NaCl solution can you prepare if you have 45.0g of NaCl? Molar mass NaCl = 58.5g / mol?
Classwork : p 54 # 19-21