Unit 5 Solids, Liquids

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Unit 5Solids, Liquids Solution Chemistry
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The Three Phases of Matter

• In solids, the molecules are locked into
  position.
• In liquids, the molecules are in close contact,
  but they can move around one another.
• In gases, the molecules are separated by large
  distances.

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The Kinetic-Molecular Theory

• The kinetic-molecular theory of matter states
• Particles of matter (atoms and molecules) are
  always in motion.
• We measure this energy of motion(kinetic energy)
  as temperature.
• If temperature increases, theparticles will gain
  more energy and move even faster.
• Molecular motion is greatest in gases, less in
  liquids, and least in solids.

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Intermolecular Forces

• The forces of attraction between molecules are
  called intermolecular forces.
• Intermolecular forces hold the molecules together
  in liquids and solids.
• Intermolecular forces vary in strength but are
  generally weaker than bonding forces.

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Polar Molecules

• Polar molecules are those with an uneven
  distribution of charge.
• Polar molecules have oppositely-charged ends
  called dipoles.
• A dipole is represented by an arrow with its
  head pointing toward the negative pole and a
  crossed tail at the positive pole.
• Water is a highly polar molecule.

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Dipole-Dipole Forces

• Dipole-Dipole Forces exist between 2 polar
  molecules. The side ofone dipole attracts
  the side of another.
• Polar molecules can also induce a temporary
  dipole in a neighboring non-polar molecule.

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Hydrogen Bonding

• When a very electronegative atom (N, O, or F) is
  bonded to hydrogen, it strongly pulls Hs
  electron toward it.
• The exposed protonof H acts as a very strong
  center of charge, attracting all the electron
  clouds from neighboring molecules.
• The resulting hydrogen bonding is the strongest
  intermolecular force that can occur in pure
  substances.

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Ion-Dipole Forces

• In Ion-Dipole Forces, ions from an ionic compound
  are attracted to the dipole of polar molecules.
• These forces are very strong, and are especially
  important in aqueous solutions of ionic compounds.

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Dispersion Forces

• Fluctuations in electron distribution can cause
  a temporary dipole.
• The attractive forces caused by these temporary
  dipoles are called dispersion forces (aka
  London Forces.) All atomsand molecules have
  them.
• They are the weakest intermolecular force, but
  will increase with increasing atomic mass.

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Types of Intermolecular Forces
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Intermolecular Forces in Action Surface Tension

• In a liquid, surface moleculeshave a higher
  potential energythan interior molecules.
• Surface Tension Liquidstend to minimize their
  surfacearea, and the surface behaveslike a
  membrane or skin.

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Intermolecular Forces in Action Viscosity

• Viscosity the resistance of a liquid to flow.
• Stronger intermolecular forces higher
  viscosity.
• Higher temperature lower viscosity.

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Intermolecular Forces in Action Meniscus

• Cohesive Forces the attraction between
  molecules in a liquid.
• Adhesive Forces the attraction between liquid
  molecules and the surface of a tube.
• Meniscus the curved surface of a liquid in a
  thin tube due to the competition between
  adhesive and cohesive forces.

Mercury Convex Meniscus Cohesive gt Adhesive
Water Concave Meniscus Adhesive gt Cohesive
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Intermolecular Forces in Action Capillary Action

• Capillary Action the vertical movement of a
  liquid through a solid tube.
• Adhesive forces draw the liquid up thetube.
  Cohesive forces pull along thosemolecules not in
  direct contact withthe tube walls.

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Vaporization

• Vaporization Phase change from liquid to gas.
• Evaporation higher-energy particles at the
  surface of a liquid escape and enter the gas
  phase.
• Boiling bubbles of gas appear throughout a
  liquid. Will not occur below a certain
  temperature (the boiling point.)
• A volatile liquid is one that vaporizes easily.

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Condensation

• Condensation Phase change from gas to liquid.
• Some vapor molecules will get captured back into
  the liquid when they collide with it.
• Also, some may stick together toform droplets of
  liquid, particularlyon surrounding surfaces.
• Vaporization is endothermic, condensation is
  exothermic.

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Sublimation Deposition

• Sublimation the phase change from solid
  directly to gas.
• Deposition the phase change from gas directly
  to solid.

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Melting Freezing

• Melting (also called fusion) phase change from
  solid to liquid.
• Freezing phase change from liquid to solid.
• Melting Freezing happen at the same
  temperature.
• Melting (fusion) is endothermic,Freezing is
  exothermic.

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Heating Curve

• Flat sections of the graph represent phase
  changes.
• During a phase change, adding more heat just
  causes a more rapid phase change. It does not
  raise the temperature of the substance.

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Phase Diagram

• Shows phases and phase changes for various
  temperature/pressure conditions.
• Areas represent phases.
• Lines represent phase changes.
• Triple point the temperature/pressure
  condition where all three states exist
  simultaneously.

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Solids

• There are two main types of solids
• Crystalline Solids Made up of crystals.
  Particles are arranged in an orderly,
  geometric, repeating pattern.
• Amorphous Solid Particles are arranged
  randomly (ex. rubber, plastic, glass.)

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Types of Crystalline Solids
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Comparing Ionic and Molecular Compounds

• Molecular compounds have relatively weak forces
  of attraction between molecules, but ionic
  compounds have a strong attraction between ions.
  This causes some differences in their properties

Ionic
Molecular
molecules
formula units
very high melting points
low melting points
hard, but brittle
usually gas or liquid
Ex NaCl, CaF2, KNO3
Ex H2O, CO2, O2
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Crystalline Solids the Crystal Lattice

• Crystalline lattice the orderly arrangement of
  the particles within a crystalline solid.
• the smallest unit that shows the pattern of
  arrangement for all the particles is called the
  unit cell.

Sodium Chloride crystal lattice (many Na and Cl
atoms) Formula Unit NaCl
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The Metallic Bond

• In metals, overlapping orbitals allow the outer
  electrons of the atoms to roam freely throughout
  the entire metal.
• These mobile electrons form a sea of electrons
  around the metal atoms, which are packed
  together in a crystal lattice.
• A metallic bond results from the attraction
  between metal atoms and the surrounding sea of
  electrons.

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Properties of Metals

• The characteristics of metallic bonding gives
  metals their unique properties, listed below
• malleability (can be hammered into thin sheets)
• ductility (can be pulledor extruded into wires)
• luster (shiny appearance)
• electrical conductivity
• thermal (heat) conductivity

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Network Covalent Solids

• Network Covalent Solid covalent bonds hold
  atoms together in a continuous network. There
  are no individual molecules, the entire crystal
  may be considered a macromolecule.(ex. diamond,
  graphite, quartz.)
• Diamond and graphite are examples of allotropes -
  forms of a chemical element that differ in their
  molecular structure.

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Solutions, Colloids, and Suspensions

• Mixtures are classified into 3 types based on
  particle size

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Suspensions

• A suspension is a mixture in which the particles
  are so large that they settle out unless the
  mixture is constantly stirred or agitated.
• Particles in a suspension are over 1000 nm in
  diameter1000 times as large as atoms.

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Colloids

• A colloid is a mixture in which the particles are
  intermediate in size between those in solutions
  and suspensions.
• Many colloids look like solutions because
  theirparticles cannot be seen.
• But colloids exhibit theTyndall effect the
  scattering of light by colloidal particles.
  (example a headlight beam on a foggy night.)

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Solutions

• A solution is a homogeneous mixture of two or
  more substances in a single phase.
• Solvent the dissolving medium.
• Solute the substance being dissolved.
• Solutions may exist as gases, liquids, or
  solids.
• Alloys solid solutions inwhich the atoms of
  two or more metals are uniformly mixed (i.e.
  brass, 14 K gold.)

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Solute-Solvent Interactions

• Like dissolves like
• Polar solutes are soluble in polar solvents.
• Non-polar solutes are insoluble in polarsolvents
  but soluble in non-polar solvents.
• Liquids that dissolve in one another are
  miscible. Liquids that dont dissolve each other
  are immiscible.

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Aqueous Solutions

• Solutions in which H2O is the solventare called
  aqueous solutions.
• H2O is highly polar. It pulls ions apart
  surrounds them, dissolving the crystal.
• Water can dissolve so many substances, it is
  known as the universal solvent.

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Dissolving a Solid in a Liquid

• The following 3 factors affect the rate at
  whicha solid dissolves in a liquid
• Surface area a solute dissolves faster if
  surface area is increased.
• Agitating a solution shaking or stirring makes
  dissolving faster.
• Increased temperature most substances
  dissolvefaster in hot water.

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Solubility of a Gas

• There are 2 main factors that affectthe
  solubility of a gas in a liquid
• Temperature decreasing the temperature
  generally increases the solubility of a gas.
• Pressure Increasing the pressure increases the
  solubility of a gas.
• Effervescence The rapid escape of a gas from a
  liquid due to a decrease in pressure.

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Saturated, Unsaturated, and Supersaturated
Solutions

• A saturated solution contains the maximum amount
  of dissolved solute.
• A solution that contains less than the maximum
  amount of solute is unsaturated.
• A supersaturated solution contains more than the
  maximum amount of solutethat can be dissolved
  under existing conditions.

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Solubility

• Solubility the amount of a substance that will
  dissolve in a specific amount of solvent at a
  certain temperature.(Ex The solubility of NaCl
  is 36 g per 100 g of water at 20C.)
• The solubility of solidsgenerally increases
  withincreasing temperature.

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Concentration

• The concentration of a solution is a ratio of
  solute to solvent.
• Concentrated means that there is a relatively
  large amount of solute in a solvent.
• Dilute means that there is a relatively small
  amount of solute in a solvent.

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Molarity

• Molarity is the number of moles of solute in
  one liter of solution.
• The symbol for molarity is M.

moles of solute
Molarity (M)
liters of solution
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Molarity CalculationsSample Problem A

• You have 3.50 L of solution that contains 90.0 g
  of NaCl. What is the molarity of that solution?
• Solution
• First, you need to convert from g to mol of
  solute.
• Then, plug your answer into the molarity
  equation

1 mol NaCl
90.0 g NaCl
1.54 mol NaCl

58.5 g NaCl
1.54 mol
mol solute
0.440 mol/L


Molarity (M)
3.50 L
L solution
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Molarity CalculationsSample Problem B

• You have 0.8 L of a 0.5 M HCl solution. How many
  moles of HCl does this solution contain?
• Solution
• First, write out M as moles/Liter
• Then, multiply by the L of solution

0.5 mol HCl
mol solute

Molarity (M)
1 L
L solution
0.5 mol HCl
0.8 L
0.4 mol HCl

1 L
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Molarity CalculationsSample Problem C

• How many grams of solute are needed to make 2.50
  L of a 1.75 M solution of Ba(NO3)2?
• Solution
• Multiply Molarity by L of solution to get mol of
  solute
• Then, convert from mol to g by using molar mass

1,140 g Ba(NO3)2
1.75 mol Ba(NO3)2
261.3 g Ba(NO3)2
2.50 L

1 L
1 mol Ba(NO3)2
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Diluting a Solution

• When an existing solution is diluted by adding
  additional solvent, the concentration changes.
• The concentration times volume of the diluted
  solution is equal to the concentration times
  volume of the original solution.

M1V1 M2V2
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Diluting a SolutionSample Problem

• During lab, you prepare 100.0 mL of a 3.00 M KNO3
  aqueous solution. If you pour all of the
  solution into a 500.0 mL flask and dilute it with
  water, what will be the final concentration of
  the new solution?
• Solution

M1V1 M2V2
(3.00 M)
(100.0 mL)
M2
(500.0 mL)

(3.00 M)
(100.0 mL)
0.600 M
M2


(500.0 mL)
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Molality

• Molality is the concentration of asolution
  expressedin moles of solute per kilogram of
  solvent.
• The symbol for molality is m.

moles of solute
molality (m)
kilograms of solvent
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Molality CalculationsSample Problem A

• Find the molal concentration of a solution
  prepared by dissolving 17.1 g of sucrose
  (C12H22O11) in 125 g of water.
• Solution
• First, you need to convert from g to mol of
  solute.
• Convert g of water to kg by dividing by 1000.
• Then, plug the numbers into the molality
  equation

1 mol C12H22O11
0.0500 mol C12H22O11
17.1 g C12H22O11

342.0 g C12H22O11
0.0500 mol
mol solute
0.400 mol/kg


molality (m)
0.125 kg
kg solvent
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Molality CalculationsSample Problem B

• How many grams of solute are needed to make a
  1.50 m solution of HNO3 in 2.00 kg H2O?
• Solution
• Multiply molality by kg of H2O to get mol of
  solute
• Then, convert from mol to g by using molar mass

1.50 mol HNO3
63.0 g HNO3
189 g HNO3
2.00 kg H2O

1 kg H2O
1 mol HNO3
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Percent Concentration

• Mass Percent mass of solute in 100 parts
  solution by mass.
• Ex. If a solution is 0.9 by mass, then there
  are 0.9 grams of solute in every 100 grams of
  solution.
• Mass of Solute Mass of Solvent Mass of
  Solution.

Mass solute (g)
x 100
Percent by Mass
Mass solution (g)
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Mass Percent CalculationSample Problem

• How much sucrose (C12H22O11), in g, is
  containedin 355 mL (12 ounces) of a soft drink
  that is 11.5 sucrose by mass? (Assume a density
  of 1.04 g/mL)
• Solution
• First, use density to change from mL to g of
  solution
• Then, multiply by the mass of solution by the
  mass (as a decimal)

M
or, M DV
(1.04 g/mL)(355 mL)
369 g
D
V
369 g
(0.115)
42.4 g C12H22O11

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Mole Fraction Mole Percent

• Mole fraction (xsolute) the ratio of moles of
  solute to total moles of solute and solvent.
• Mole percent mole fraction x 100.

solute (mol)
Mole fraction
solute solvent (mol)
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Mole Fraction Mole PercentSample Problem

• 17.2 g of ethylene glycol (C2H6O2) is dissolved
  in 0.500 kg H2O. Calculate mole fraction and
  mole .
• Solution
• First, convert mass to moles for solute and
  solvent
• Then, find mole fraction
• Finally, multiply by 100 to get mole percent

1 mol C2H6O2
17.2 g C2H6O2

0.277 mol C2H6O2
62.0 g C2H6O2
1000 g H2O
1 mol H2O
0.500 kg H2O

27.8 mol H2O
1 kg H2O
18.0 g H2O
0.277 mol
solute (mol)
9.86 x 10-3


0.277 27.8 mol
solute solvent (mol)
0.986
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Colligative Properties

• Colligative properties depend on concentration
  of the solute particles but not on their
  identity.
• Colligative properties include vapor-pressure
  lowering, freezing-point depression,
  boiling-point elevation, and osmotic pressure.
• The ratio of moles of particles in solution to
  moles of formula units dissoved is called the
  Vant Hoff factor (i)

Moles of particles in solution
i
Moles of formula units dissolved
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Vapor Pressure

• When the rate of vaporization and the rate of
  condensation become equal in a closed container,
  dynamic equilibrium has been reached.
• The pressure of a gas in dynamic equilibrium with
  its liquid is called vapor pressure.

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Vapor-Pressure Lowering

• A nonvolatile substance has little tendency to
  become a gas under existing conditions.
• The vapor pressure of a solvent is lower when a
  nonvolatile substance is added to it.
• This causes the solutionto be liquid over a
  largertemperature range,lowering the freezing
  point and raising the boiling point.

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Freezing-Point Depression

• Freezing-Point Depression the lowering of the
  freezing point of a solution compared to the
  puresolvent proportional to the molal
  concentration of the solution.

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Boiling-Point Elevation

• Boiling-Point Elevation the raising of the
  boiling point of a solution compared to the pure
  solvent proportional to the molal concentration
  of the solution.

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Electrolytes

• Electrolyte a substance that conducts
  electricity when it is dissolved in water.
• Ionic compounds are generally good electrolytes.
  Molecular compounds are not.

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Dissociation

• Dissociation is the separation of ions that
  occurs when an ionic compound dissolves.
• Examples
• NaCl(s) ? Na(aq) Cl-(aq)
• CaCl2(s) ? Ca2(aq) 2 Cl-(aq)

H2O
1 mole
1 mole
1 mole
H2O
1 mole
1 mole
2 moles
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DissociationSample Problem

• a.) Write the equation for the dissociation of
  aluminum sulfate in water.
• b.) How many moles of ions are produced by
  dissolving 1 mol of Al2(SO4)3?
• c.) Assuming complete dissociation, what is the
  Vant Hoff factor of Al2(SO4)3?

H2O
Al2(SO4)3(s)
Al3(aq)
SO42-(aq)
2
3
5 moles
Moles of particles in solution
5
5
i


Moles of formula units dissolved
1
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Effect of Dissociation

Because of the forces of attraction between
them, ions cluster together somewhat, producing
measured (actual) Vant Hoff factors that are
somewhat lower than expected.
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vant Hoff FactorsSample Problem

• Predict the vant Hoff factors for each of the
  following aqueous solutions. Which solution will
  have the highest boiling point?
• Solution Expected Vant Hoff Factor
• 1.0 M C12H22O11
• 1.0 M NaCl
• 1.0 M MgCl2
• 1.0 M P4O10
• 1.0 M Fe(NO3)3
• 1.0 M CuSO4

1
2
3
Greatest boiling point elevation
1
4
2
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Calculating Freezing Point Depression

• The amount that the freezing point is lowered for
  solutions is given by the equation
• DTf is the change in temperature of the freezing
  point in oC.
• i is the vant Hoff factor for the solute.
• m is the molality of the solution in mol/kg.
• Kf is a constant for the solvent.Kf for water is
  -1.86 oC/m.

DTf imKf
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Freezing Point DepressionSample Problem

• Calculate the freezing point of a 1.3 m saline
  (aqueous sodium chloride) solution.
• Solution
• The freezing point of pure water is 0.00 oC
• So, the new freezing point would be

DTf imKf
(-1.86 oC/m)
- 4.8 oC
DTf
(1.3 m)
(2)
- 4.8 oC
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Calculating Boiling Point Elevation

• The amount that the freezing point is lowered for
  solutions is given by the equation
• DTb is the change in temperature of the boiling
  point in oC.
• i is the vant Hoff factor for the solute.
• m is the molality of the solution in mol/kg.
• Kb is a constant for the solvent.Kb for water is
  0.512 oC/m.

DTb imKb
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Boiling Point ElevationSample Problem

• How many grams of ethylene glycol (C2H6O2) must
  be added to 1.0 kg of water to produce a solution
  that boils at 105.0oC?
• Solution
• First, solve for molality
• Then use molality to solve for grams of C2H6O2

5.0 oC
DTb
9.8 m

or, m
DTb imKb
(0.512 oC/m)
(1)
iKb
9.8 mol C2H6O2
62.0 g C2H6O2
1.0 kg H2O
610 g C2H6O2

1 kg H2O
1 mol C2H6O2
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To Summarize

• Watch the official unit summary here
• http//www.youtube.com/watch?vVTmfQUNLlMY