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Solutions
Solutions• An aqueous solution is water that
contains dissolved substances.
Solvents and Solutes•In a solution, the dissolving medium is the solvent.
•The dissolved particles in a solution are the solute.
Solutions• A solvent dissolves the solute.
• The solute becomes dispersed in the solvent.
• Solvents and solutes may be gases, liquids, or solids.
• Solutions are homogeneous mixtures.
• Solute particles can be atoms, ions, or molecules.
• If you filter a solution through filter paper, both the solute and solvent pass through the filter.
Solutions
Substances that dissolve most readily in water include ionic compounds and polar covalent compounds.
• Nonpolar covalent compounds do not dissolve in water. Examples of nonpolar covalent compounds include methane (CH4), and compounds found in oil, grease, and gasoline.
• However, oil and grease will dissolve in gasoline.
What types of substances dissolve most readily in water?
Solutions
The Solution Process• A water molecule is polar, with a partial
negative charge on the oxygen atom and partial positive charges on the hydrogen atoms.
• As individual solute ions break away from the crystal, the negatively and positively charged ions become surrounded by solvent molecules and the ionic crystal dissolves.
Solutions
•The process by which the positive and negative ions of an ionic solid become surrounded by solvent molecules is called solvation.
Solvated ions
Surface of ionic solid
The Solution Process
Solutions
The Solution Process• Polar solvents such as
water dissolve ionic compounds and polar compounds.
• Nonpolar solvents such as gasoline dissolve nonpolar compounds.
• This relationship can be summed up in the expression “like dissolves like.”
Electrolytes and Nonelectrolytes electrolyte - a compound that conducts an electric current when it dissolves in water or is melted.
All ionic compounds are electrolytes because they dissociate into ions.
nonelectrolyte - a compound that does not conduct an electric current ever.
nonelectrolytes do not dissociate in water
Electrolytes and Nonelectrolytes•Some polar molecular compounds are nonelectrolytes in the pure state but become electrolytes when they dissolve in water. (ex. Ammonia NH3)• For example, ammonia (NH3(g)) is not an
electrolyte in the pure state.
• Yet an aqueous solution of ammonia conducts an electric current because ammonium ions (NH4
+) and hydroxide ions (OH–) form when ammonia dissolves in water.
NH3(g) + H2O(l) NH4+(aq) + OH–(aq)
Electrolytes and Nonelectrolytes•Not all electrolytes conduct electric current to the same degree.
• In a solution that contains a strong electrolyte, all or nearly all of the solute exists as ions.
• A weak electrolyte conducts an electric current poorly because only a fraction of the solute in the solution exists as ions.
Electrolytes and Nonelectrolytes•Your cells use electrolytes, such as sodium and potassium ions, to carry electrical impulses across themselves and to other cells.
• An electrolyte imbalance can occur if you become dehydrated.
• When you exercise, you can lose water and electrolytes from your body through perspiration.
•The water contained in a crystal is called the water of hydration or water of crystallization.
Hydrates
• hydrate – a compound that contains water of hydration
• anhydrous - A substance that does not contain water.
Hydrates
Heating of a sample of blue CuSO45H2O begins.
After a time, much of the blue hydrate has been converted to white anhydrous CuSO4.
CuSO45H2O(s) CuSO4(s) + 5H2O(g)– heat
+ heat
Each hydrate contains a fixed quantity of water and has a definite composition.
Hydrates
Some Common Hydrates
Formula Chemical name Common name
MgSO47H2O Magnesium sulfate heptahydrate Epsom salt
Ba(OH)28H2O Barium hydroxide octahydrate
CaCl22H2O Calcium chloride dihydrate
CuSO45H2O Copper(II) sulfate pentahydrate Blue vitriol
Na2SO410H2O Sodium sulfate decahydrate Glauber’s salt
KAl(SO4)212H2OPotassium aluminum sulfate dodecahydrate
Alum
Na2B4O710H2O Sodium tetraborate decahydrate Borax
FeSO47H2O Iron(II) sulfate heptahydrate Green vitriol
H2SO4H2O Sulfuric acid hydrate (mp 8.6oC)
Hydrates
Efflorescent Hydrates
• If a hydrate has a vapor pressure higher than the pressure of water vapor in the air, the hydrate will lose its water of hydration, or effloresce.
The water molecules in hydrates are held by weak forces, so hydrates often have an appreciable vapor pressure.
Hydrates
Hygroscopic Hydrates
• Hydrates and other compounds that remove moisture from air are hygroscopic.
Hydrated ionic compounds that have low vapor pressure remove water from moist air to form higher hydrates.
Hydrates
Hygroscopic Hydrates
• Calcium chloride is used as a desiccant in the laboratory.
desiccant - a substance used to absorb moisture from the air and create a dry atmosphere.
Calcium chloride monohydrate spontaneously absorbs a second molecule of water when exposed to moist air.
Hydrates
Deliquescent Compounds
• Deliquescent - compound that removes sufficient water from the air to dissolve completely and form solutions.
Some compounds are so hygroscopic that they become wet when exposed to normally moist air.
Pellets of sodium hydroxide are deliquescent.
For this reason, containers of NaOH should always be tightly stoppered.
The solution formed by a deliquescent substance has a lower vapor pressure than that of the water in the air.
Hydrates
•A suspension differs from a solution because the particles of a suspension are much larger and do not stay suspended indefinitely.
Suspensions
• The particles in a typical suspension have an average diameter greater than 1000 nm.
• By contrast, the particle size in a solution is usually about 1 nm.
suspension - mixture from which particles settle out upon standing.
Suspensions
• A solution is a homogeneous mixture.• Suspensions are heterogeneous because at
least two substances can be clearly identified.
Suspensions•The difference between a solution and suspension is easily seen when the type of mixture is filtered.
The small size of the solute particles in a solution allows them to pass through filter paper.
The particles of a suspension can be removed by filtration.
•Explain why a mixture of sand and water can be separated by filtration, but a mixture of salt and water cannot.
A mixture of sand and water is a suspension, and a mixture of salt and water is a solution. The particles in the sand mixture are much larger than the ions in the salt mixture. The sand particles are too large to pass through filter paper; the ions are not.
Colloids
A colloid is a heterogeneous mixture containing particles that range in size from 1 nm to 1000 nm.
• The particles in a colloid are spread, or dispersed, throughout the dispersion medium, which can be a solid, liquid, or gas.
Colloids
The first substances to be identified as colloids were glues.
Some Colloidal Systems
System
Type ExampleDispersed phase
Dispersion medium
Gas Liquid Foam Whipped cream
Gas Solid Foam Marshmallow
Liquid Liquid Emulsion Milk, mayonnaise
Liquid Gas Aerosol Fog, aerosol
Solid Gas Smoke Dust in air
Solid Liquid Sols, gelsEgg white, jelly, paint, blood, starch in water, gelatin
Colloids
•Colloids have particles smaller than those in suspensions and larger than those in solutions.
• These intermediate-sized particles cannot be retained by filter paper as are the larger particles of a suspension.
• They do not settle out with time.
Colloids
The Tyndall Effect
•You cannot see a beam of sunlight unless the light passes through particles of water (mist) or dust in the air.• These particles scatter the sunlight.
• Similarly, a beam of light is visible as it passes through a colloid.
Colloids
The Tyndall Effect
•The scattering of visible light by colloidal particles is called the Tyndall effect.
Flashlight
Solution Colloid Suspension
Colloids
• Suspensions also exhibit the Tyndall effect.• The particles in solutions are too small to
scatter light.
Flashlight
Solution Colloid Suspension
CHEMISTRY & YOUCHEMISTRY & YOU
•What would be the ideal conditions to see a red sunset?
A misty or foggy evening would be ideal for seeing a red sunset. There would be a large number of particles to scatter the sunlight.
Colloids
Brownian Motion
•Flashes of light, or scintillations, are seen when colloids are studied under a microscope.• This happens because the particles
reflecting and scattering the light move erratically.
ColloidsBrownian Motion
The chaotic movement of colloidal particles, which was first observed by the Scottish botanist Robert Brown (1773–1858), is called Brownian motion.
Colloids
Brownian motion is caused by collisions of the molecules of the dispersion medium with the small, dispersed colloidal particles.• These collisions help prevent the colloidal
particles from setting.
Brownian Motion
Colloids
Coagulation
A colloidal system can be destroyed or coagulated by the addition of electrolytes.
• The added ions neutralize the charged colloidal particles.
• The particles can clump together to form heavier aggregates and settle out from the dispersion.
Colloids
Emulsions
•An emulsion is a colloidal dispersion of a liquid in a liquid.
• An emulsifying agent is essential for the formation of an emulsion and for maintaining the emulsion’s stability.
Colloids
• Oils and greases are not soluble in water. • However, oils and greases readily form a
colloidal dispersion if soap or detergent is added to the water.
Emulsions
Colloids
• One end of a large soap or detergent molecule is polar and is attracted to water molecules.
• The other end of the soap or detergent molecule is nonpolar and is soluble in oil or grease.
• Soaps and other emulsifying agents allow the formation of colloidal dispersions between liquids that do not ordinarily mix.
Emulsions
Colloids
This table summarizes the properties of solutions, colloids, and suspensions.
Properties of Solutions, Colloids, and Suspensions
Property
System
Solution Colloid Suspension
Particle typeIons, atoms, small molecules
Large molecules or particles
Large particles or aggregates
Particle size 0.1–1 nm 1–1000 nm 1000 nm and larger
Effect of light No scattering Exhibits Tyndall effect Exhibits Tyndall effect
Effect of gravity
Stable, does not separate
Stable, does not separate
Unstable, sediment forms
FiltrationParticles not retained on filter
Particles not retained on filter
Particles retained on filter
Uniformity Homogeneous Heterogeneous Heterogeneous
•Calculate the percent by mass of water in washing soda, sodium carbonate decahydrate (Na2CO310H2O).
Hydrates
Solution Formation•Granulated sugar dissolves faster than sugar cubes, and both granulated sugar and sugar cubes dissolve faster in hot tea or when you stir the tea.
Solution FormationComposition of solvent and solute determine whether or not a substance will dissolve.
•Factors that affect how fast a substance dissolves include:
• Agitation (stirring or shaking)
• Temperature
• Particle size of the solute
Solution Formation• AgitationIf the contents of the glass are stirred, the crystals dissolve more quickly.• The dissolving process
occurs at the surface of the sugar crystals.
• Stirring speeds up dissolving because fresh solvent is continually brought in contact with the surface of the solute.
Solution Formation• AgitationAgitation (stirring or shaking) affects only the rate at which a solid solute dissolves.• It does not influence the
amount of solute that will dissolve.
• An insoluble substance remains undissolved regardless of how vigorously or for how long the solvent/solute system is agitated.
Solution Formation
•Temperature also influences the rate at which a solute dissolves.
• Sugar dissolves much more rapidly in hot tea than in iced tea.
• Most solids dissolve faster at higher temperatures.
Temperature
Solution FormationTemperature•At higher temperatures, the kinetic energy of water molecules is greater than at lower temperatures, so the •molecules move faster.
• At higher temperatures, solvent molecules move faster, resulting in more frequent collisions between solute and solvent molecules, resulting in faster dissolving.
Solution Formation• Particle Size of the SoluteThe rate at which a solute dissolves also depends upon the size of the solute particles.
• The smaller particles in granulated sugar expose a much greater surface area to the colliding water molecules.
Solution Formation• Particle Size of the SoluteThe dissolving process is a surface phenomenon.• Smaller particles dissolve
faster because more surface area of the solute is exposed, speeding up the rate of dissolving.
•Which of the following will not speed up the rate at which a solid solute dissolves?
A. Increasing the temperature
B. Stirring the mixture
C. Crushing the solute
D. Decreasing the temperature
•Which of the following will not speed up the rate at which a solid solute dissolves?
A. Increasing the temperature
B. Stirring the mixture
C. Crushing the solute
D. Decreasing the temperature
Solubility•What is happening in this figure?
• Particles move from the solid into the solution.
• Some dissolved particles move from the solution back to the solid.
• When two processes occur at the same rate, no net change occurs in the overall system (equilibrium).
SolubilitySuch a solution is said to be saturated.
• saturated solution contains the maximum amount of solute for a given quantity of solvent at a constant temperature and pressure.
Solubility
•In a saturated solution, a state of dynamic equilibrium exists between the solution and any undissolved solute, provided that the temperature remains constant.
Solubilitysolubility of a substance is the amount of solute that dissolves in a given quantity of a solvent at a specified temperature and pressure to produce a saturated solution.• Solubility is usually expressed as
• grams of solute per 100 g of solvent • (g solute/100 g H2O).• Sometimes the solubility of a gas is expressed in
grams per liter of solution (g/L).
Solubility
unsaturated solution contains less solute than a saturated solution at a given temperature and pressure
• If additional solute is added to an unsaturated solution, the solute will dissolve until the solution is saturated.
Solubilitymiscible liquids dissolve in each other in all proportions.
• Ex: water and ethanol
immiscible - Liquids that are insoluble in each other
• Ex: oil and water
Factors Affecting SolubilitySolubility for most solid substances increases as the temperature of the solvent increases.
Interpret GraphsInterpret Graphs
Temperature (°C)
So
lub
ility
(g
/100
g H
2O)• For a few
substances, solubility decreases with temperature.
Solubilities of Substances in Water at Various Temperatures
Solubility (g/100 g H2O)
Substance Formula 0°C 20°C 50°C 100°C
Barium hydroxide Ba(OH)2 1.67 31.89 — —
Barium sulfate BaSO4 0.00019 0.00025 0.00034 —
Calcium hydroxide Ca(OH)2 0.189 0.173 — 0.07
Potassium chlorate KClO3 4.0 7.4 19.3 56.0
Potassium chloride KCl 27.6 34.0 42.6 57.6
Sodium chloride NaCl 35.7 36.0 37.0 39.2
Sodium nitrate NaNO3 74 88.0 114.0 182
Aluminum chloride AlCl3 30.84 31.03 31.60 33.32
Silver nitrate AgNO3 122 222.0 455.0 733
Sucrose (table sugar) C12H22O11 179 230.9 260.4 487
Hydrogen H2 0.00019 0.00016 0.00013 0.0
Oxygen O2 0.0070 0.0043 0.0026 0.0
Carbon dioxide CO2 0.335 0.169 0.076 0.0
Interpret DataInterpret Data
Factors Affecting Solubility
supersaturated solution contains more solute than it can theoretically hold at a given temperature.
• The crystallization of a supersaturated solution can be initiated if a very small crystal, called a seed crystal, of the solute is added.
Factors Affecting Solubility
•The rate at which excess solute deposits upon the surface of a seed crystal can be very rapid.
The solution is clear before a seed crystal is added.
Crystals begin to form immediately after the addition of a seed crystal.
Excess solute crystallizes rapidly.
Factors Affecting Solubility
The effect of temperature on the solubility of gases in liquid solvents is opposite that of solids.
• Solubility for most gases is greater in cold water than in hot.
Temperature
Factors Affecting Solubility
Changes in pressure have little effect on the solubility of solids and liquids, but pressure strongly influences the solubility of gases.
• Gas solubility increases as the partial pressure of the gas above the solution increases.
Pressure
Factors Affecting SolubilityPressure
Carbonated beverages are a good example.
• These drinks contain large amounts of carbon dioxide (CO2) dissolved in water.
• Dissolved CO2 makes the liquid fizz
Factors Affecting Solubility
Pressure
• The drinks are bottled under a high pressure of CO2 gas, which forces larger amounts of the gas into solution.
Factors Affecting Solubility
Pressure
• When the container is opened, the partial pressure of CO2 above the liquid decreases.
• Immediately, bubbles of CO2 form in the liquid and escape from the open bottle.
Factors Affecting Solubility
•How is the partial pressure of carbon dioxide gas related to the solubility of CO2 in a carbonated beverage?• The relationship is described by Henry’s
law, which states that at a given temperature, the solubility (S) of a gas in a liquid is directly proportional to the pressure (P) of the gas above the liquid.
Pressure
Factors Affecting Solubility
• As the pressure of the gas above the liquid increases, the solubility of the gas increases.
• As the pressure of the gas decreases, the solubility of the gas decreases.
Pressure
•Explain why an opened container of a carbonated beverage is more likely to go flat sitting on the counter than in the refrigerator.
The solubility of a gas in a liquid increases with decreasing temperature. More carbon dioxide will remain in solution at the colder temperature found in the refrigerator.
Factors Affecting Solubility
Factors Affecting Solubility
•What factors affect the solubility of a substance?
Temperature affects the solubility of solid, liquid, and gaseous solutes in a solvent; both temperature and pressure affect the solubility of gaseous solutes.
Key Concepts
•Factors that determine how fast a substance dissolves are stirring, temperature, and surface area.
In a saturated solution, a state of dynamic equilibrium exists between the solution and any undissolved solute, provided that the temperature remains constant.
Temperature affects the solubility of solid, liquid, and gaseous solutes in a solvent; both temperature and pressure affect the solubility of gaseous solutes.