Title: Unit 5 Solids, Liquids
1Unit 5Solids, Liquids Solution Chemistry
2The 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.
3The 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.
4Intermolecular 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.
5Polar 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.
6Dipole-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.
7Hydrogen 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.
8Ion-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.
9Dispersion 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.
10Types of Intermolecular Forces
11Intermolecular 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.
12Intermolecular Forces in Action Viscosity
- Viscosity the resistance of a liquid to flow.
- Stronger intermolecular forces higher
viscosity. - Higher temperature lower viscosity.
13Intermolecular 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
14Intermolecular 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.
15Vaporization
- 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.
16Condensation
- 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.
17Sublimation Deposition
- Sublimation the phase change from solid
directly to gas. - Deposition the phase change from gas directly
to solid.
18Melting 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.
19Heating 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.
20Phase 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.
21Solids
- 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.)
22Types of Crystalline Solids
23Comparing 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
24Crystalline 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
25The 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.
26Properties 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
27Network 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.
28Solutions, Colloids, and Suspensions
- Mixtures are classified into 3 types based on
particle size
29Suspensions
- 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.
30Colloids
- 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.)
31Solutions
- 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.)
32Solute-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.
33Aqueous 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.
34Dissolving 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.
35Solubility 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.
36Saturated, 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.
37Solubility
- 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.
38Concentration
- 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.
39Molarity
- 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
40Molarity 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
41Molarity 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
42Molarity 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
43Diluting 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
44Diluting 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)
45Molality
- 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
46Molality 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
47Molality 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
48Percent 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)
49Mass 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
50Mole 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)
51Mole 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
52Colligative 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
53Vapor 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.
54Vapor-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.
55Freezing-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.
56Boiling-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.
57Electrolytes
- Electrolyte a substance that conducts
electricity when it is dissolved in water. - Ionic compounds are generally good electrolytes.
Molecular compounds are not.
58Dissociation
- 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
59DissociationSample 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
60Effect 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.
61vant 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
62Calculating 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
63Freezing 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
64Calculating 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
65Boiling 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
66To Summarize
- Watch the official unit summary here
- http//www.youtube.com/watch?vVTmfQUNLlMY