Water: The Universal Solvent

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Water The Universal Solvent

• Polarity makes water special.
• O atom is an electron hog in the covalent bond
  it forms with 2 H atoms.
• Water has a positive end and a negative end.
• This makes it a good solvent for ionic solids.

Slightly negative
O
H
H
Slightly positive
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Waters Properties

• Hexagonal crystal shape
• Molecule is polar.
• Hydrogen bonding
• Ice floats.
• Expands during freezing until -4.0 º C.
• Solid form is less dense than liquid
• Surface tension
• Water beads on smooth surfaces.
• Insects walk on water surfaces.

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Surface tension

• Force that pulls adjacent parts of a liquid
  surface together.
• The higher the attractive forces between
  particles in the liquid, the higher the surface
  tension.
• Hydrogen bonds make water have higher surface
  tension than most liquids.

Soap
Water droplet
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Solutions

• Same as a homogeneous mixture
• Solvent substance that takes in, or dissolves,
  another substance
• Solute substance that is taken in, or dissolved
  by the first substance
• The 3 phases of matter can be paired together in
  9 ways to make solutions.

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Solutions and phases of matter
Solute (dissolved) (in) Solvent Example
Gas Gas O2 in N2 - Air
Gas Liquid CO2 in water
Gas Solid Hydrogen fuel storage
Liquid Gas Water in air
Liquid Liquid Acid in water (vinegar)
Liquid Solid Amalgam dental fillings (Hg in Ag)
Solid Gas Mothballs in air
Solid Liquid Salt water
Solid Solid Metal alloys - jewelry
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Factors which affect solution rate and solubility

• Solution rate
• Temperature
• Particle size
• Movement
• Solubility
• Temperature
• Pressure

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Solubility

• Solids Increased temp increases solubility.
• Gases Solubility in water decreases when
  temperature increases.
• Gases Increased pressure increases solubility.
• More fizz in a soda bottle if you store with lid
  on (increased pressure) and in the refrigerator
  (decreased temperature).

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Concentration

• Relationship between solute and solvent
• Concentration volume solute/volume solvent
• Relative concentration of solutions
• Unsaturated More solute can be dissolved at a
  given temp. and pressure. May be dilute or
  concentrated.
• Saturated Maximum solute a solvent will
  dissolve at a given temp. and pressure.
• Supersaturated More than the maximum solute a
  solvent will dissolve at a given temp. and
  pressure. (Usually obtained by heating a
  saturated solution to very high temp and cooling
  slowly.)

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Suspensions and Colloids

• Remember A solution is a homogeneous mixture.
• Some mixtures are not solutionsIt depends on the
  size of the particles.
• Large particles usually dont make homogeneous
  mixtures.
• The names of 2 types of mixtures with large
  particles are suspensions and colloids.

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Suspensions

• A mixture in which some particles will settle out
  is a suspension.
• Gravity will eventually make large particles
  sink.
• Particles are large enough to be filtered out.
• Examples Muddy water and salad dressing.
• Particles are large enough to reflect light, so a
  beam of light will be visible in a suspension.

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Comparing mixtures

• Solution homogeneous, small particles
• Suspension heterogeneous, large particles
• Colloid in-between suspension and solution.
• Is not completely homogeneous.
• Examples fog, homogenized milk, toothpaste.

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Colloids

• How can you tell is a mixture is a colloid?
• A beam of light will be visible when shone
  through a colloid (like a suspension).
• Particles do not separate or settle out after
  standing.

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Colloids 5 types

• Sol
• Solid is dispersed in solid or liquid.
• Examples Clay, toothpaste, red glass
• Gel
• Long particles that trap liquid inside.
• Examples Jam, jelly, gelatin dessert, hair gel
• Aerosol
• Solid is dispersed in air.
• Examples Fog, some hair sprays
• Foam
• Gas is dispersed in liquid.
• Examples Marshmallows, whipped cream
• Emulsion
• Liquid is dispersed in liquid.
• Examples Milk, butter, some cosmetics

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Solutes Electrolytes

• Electrolyte Substance that dissolves in water
  to give a solution that conducts electricity
• Mobile charged particles conduct electricity
  (ions)
• Soluble ionic compounds are electrolytes
• (Na and Cl-)
• Many acids are electrolytes
• (HCl gtH3O and Cl-)

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Solute-Solvent Interactions

• Like dissolves like
• Polar dissolves polar Water NaCl
• Nonpolar dissolves nonpolar CCl4 motor oil
• Immiscible liquids (oil and vinegar)
• Do not mix
• Liquid solutes and solvents that are not soluble
  in each other
• Miscible liquids (gas and oil)
• Mix
• Liquids that dissolve freely in one another in
  any proportion

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Solubility Rules
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Henrys LawSolubility of gas in liquid

• The solubility of a gas is directly proportional
  to the partial pressure of that gas on the
  surface of the liquid.
• Soda bottle
• High pressure at the surface while the bottle is
  closed, so lots of CO2 in the liquid
• Open bottle, pressure on surface lowers to room
  atmosphere and CO2 leaves the liquid
• High pressure High gas concentration
• Low pressure low gas concentration

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Phase changes depend on forces holding molecules
together
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Vaporization

• Enough energy is added to overcome the forces
  holding molecules together in the liquid (or
  solid) state
• Boiling, evaporation or sublimation

Covalent bond
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Intermolecular forces Generalizing properties

• Low boiling point particles are more likely to
  leave liquid solution
• Weaker IM forces lower boiling point
• Lower boiling point more vapor higher vapor
  pressure
• High boiling point slow evaporation
• If IM forces are the same, look at formula
  weight. Heavier molecules have higher boiling
  points.
• Strength of IM forces
• Hydrogen bondgtdipole-dipolegtLondon dispersion

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Boiling Point ? Vapor Pressure?
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Dispersion Forces

• Weakest forces attracting molecules together
• Dependent on momentary polarity induced by
  protons and electrons

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Dispersion Forces

• More electrons more polarizability
• Heavier molecules have stronger IMFs because they
  are more polarizable. (If they have same type of
  forces)

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Dipole-Dipole Forces

• Between molecules made up of atoms with differing
  electronegativities Typically molecules
  containing polar covalent bonds

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Dipole-Dipole forces, cont.

• As with dispersion forces, when comparing two
  substances having dipole-dipole interactions as
  their strongest IMFs, the most polarizable will
  have stronger IMFs.
• HCl (36.5 amu, 17 electrons in Cl) should have
  higher boiling point than HBr (81 amu, 35
  electrons in Br)

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Hydrogen bonds

• These are NOT bonds
• Attraction between H and N,O or F

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Intermolecular Forces
Interacting molecules or ions
No
Yes
Yes
No
No
Yes
No
Yes
London Forces only Ex. Ar(l), I2(s)
Dipole-Dipole Ex. H2S
Hydrogen Bonding Ex. NH3, H2O
Ion-dipole Forces Ex. KBr in H2O
Ionic bonding Ex. NaCl
Chart from Chemistry The Central Science by
Brown-LeMay et. al
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Phases of matter Comparison
Property Solid Liquid Gas
Particles Closely packed High density (ButWater is different!) More densely packed than in gas Most compressible-least densely packed
Particle movement Vibrate weakly around fixed positions Lowest kinetic energy Can change positions with other particles Can change positions with other particles Highest kinetic energy
Intermolecular forces Most effective (strongest) Stronger than in gases Least effective (weakest)
Shape and volume Both definite Definite volume only No definite shape or volume
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Changing states

• Equilibrium When there is no net change in a
  system.
• Dynamic equilibrium
• When a vapor is in equilibrium with its liquid as
  one molecule leaves the liquid to become a vapor,
  another molecule leaves the vapor to become a
  liquid. In other words, an equal number of
  molecules will be found moving in both
  directions.

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Boiling Point

• Vapor pressure Pressure exerted by a vapor
  Pressure of the liquid at given temperature
• Liquid boils when its vapor pressure equals
  pressure of the atmosphere.
• Boiling is the conversion of a liquid to vapor
  within the liquid as well as at its surface.
• Boiling point is the temperature at which the
  equilibrium vapor pressure of the liquid equals
  the atmospheric pressure.
• Volatile liquids are liquids that evaporate
  readily.

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Boiling Point, cont.

• High elevation Low atmospheric pressure
• Low atmospheric pressure lower boiling point
• High pressure in pressure cooker increased
  boiling point, faster cooking
• If pressure above liquid increases, the liquid
  temperature rises until it matches the new
  pressure and boils again.

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Separation by Distillation

• Distillation is the separation of liquid
  substances according to their different boiling
  points.
• As a liquid mixture is heated, the substance with
  the lower boiling point will vaporize first.
• Distillate Condensed liquid substance

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Phase DiagramA phase diagram is a graph of
pressure vs. temperature that shows the
conditions under which phases of matter exist.
Critical temp (Tc) Above this, the substance
cannot exist in the liquid state.
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Phase Diagrams Density

• Negative liquid/solid slope shows density of
  solid is LESS than
• liquid (like H2O). See previous slide.
• Most substances will have a positive slope of
  this line since most solids are more dense than
  the liquid

http//wine1.sb.fsu.edu/chm1045/notes/Forces/Phase
/Forces06.htm
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Four major "points" on a phase diagram

• Triple point, TP - All three phases can exist in
  equilibrium at this temperature and pressure.
• (The solid-liquid line and the liquid-vapor line
  meet.)
• Normal boiling point, Tb - The temperature at
  which the vapor pressure of a liquid is equal to
  standard atmospheric pressure.
• (Standard atmospheric pressure line crosses the
  liquid-vapor line.)
• Normal melting point, Tm - The temperature at
  which the vapor pressure of the solid and the
  vapor pressure of the liquid are equal.
• (Standard atmospheric pressure line crosses
  the solid-liquid line.)
• Critical temperature, Tc - The temperature above
  which no amount of pressure will liquefy a vapor.
  
• (The liquid-vapor line becomes vertical.)

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Kinetic Energy and Equilibrium Vapor Pressure

• In the beginning
• particles condensing to liquid phase
• particles evaporating to gas phase
• Increase temp Increase kinetic energy
• Now, more molecules have enough energy to leave
  the liquid.
• More vapor molecules higher vapor pressure
• Equilibrium will soon be established, but at a
  higher vapor pressure.

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Colligative properties

• vant Hoff factor i
• i number of ions one unit of substance will
  dissociate into in solution
• CaCl2 dissociates into Ca2, Cl- and Cl-, so i
  3.
• Boiling Point Elevation
• ?T kbmi
• ?T increase in solution boiling point
• kb boiling point elevation constant for the
  solvent
• m molality

(for now)
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• Freezing Point Depression
• ?T kfmi
• ?T decrease in solution freezing point
• kf freezing point depression constant for the
  solvent
• m molality
• Vapor Pressure Lowering
• P XPº
• P vapor pressure of the solution
• X mole fraction of the solvent
• Pº vapor pressure of the pure solvent
• Add solute Solutions vapor pressure goes down.
• Lower vapor pressure Raise boiling point

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Vapor Pressure Raoults Law

• Vapor gas formed by the boiling or evaporation
  of a liquid or a solid
• Vapor pressure of a liquid is the pressure
  exerted by the vapor in equilibrium with its
  liquid.
• A solution containing a nonvolatile solute has a
  lower vapor pressure than the pure solvent.

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Raoults Law

• Psolvent XsolventPsolvent
• Psolvent Vapor pressure of solvent in the
  solution
• Xsolvent Mole fraction of solvent in solution
• Psolvent Vapor pressure of the pure solvent
• As the mole fraction (solvent) goes up, its
  vapor pressure also goes up proportionally.
• If the solute is nonvolatile, then Psolvent
  Psolution
• Lowering of vapor pressure is defined
• ?PsolventPsolvent Psolvent OR ? Psolvent
  XsolutePsolvent