POLAR BONDS AND MOLECULES

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POLAR BONDS AND MOLECULES

• NOTES 16.3

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Covalent Bonds

• bond in which two atoms share a pair of
  electrons.
• Single bond 1 shared pair of electron
• Double bonds 2 shared pairs
• Triple bonds 3 shared pairs

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Bond Polarity

1. bonding pairs of electrons in covalent bonds are
   pulled between the nuclei of the atoms sharing
   the electrons.
2. When bonds are pulled equally the bond is a
   nonpolar covalent (occurs between like atoms.)

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• 3. When a covalent bond occurs between different
  atoms then electrons are shared unequally which
  is a polar bond.

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• This leads to partial charges on the atoms (?- or
  ? )
• b. The polarity of a bond can be shown with an
  arrow pointing to the negative side.
• c. Table 16.4 p. 462

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Polar Molecules

• when one end of a molecule is slightly negative
  and other is slightly positive.
• - Water is polar because the way the bonds
  cause the molecule to bend. You have partial
  charges on two sides of molecule.

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• - CO2 is not polar because the double bonds
  keep the molecule linear and the charges cancel.

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Attractions

• molecules are attracted to each other through a
  variety of forces.

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• 1. These forces are responsible for determining
  whether a compound is a gas, liquid or solid.
• 2. Van der Waals forces consist of two possible
  types.

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• a. Dispersion ? caused by the movement of
  electrons. Increases as the of electrons
  increase.
• b. Dipole Interactions ? occurs when polar
  molecules are attracted to one another.

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• 1. Hydrogen bonds occur when H already in a polar
  compound bond with a partial negative of another
  molecule.
• - extremely important in determining the
  properties of water and biological molecules
  such as proteins.

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Intermolecular Attractions

• The physical properties of a compound depend on
  the type of bonding it displays.

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WATER AND AQUEOUS SYSTEMS

• NOTES 17.1

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Water Molecule

1. It is a triatomic molecule with two polar
   covalent bonds (H O).
2. It has a bent shape leading to a partially (d)
   and (d) ends to the molecule.

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• 3. Because of the polarity the molecule will
  form Hydrogen bonds with other water molecules.

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Water Molecule -- Polarity

• - Water is polar because the way the bonds
  cause the molecule to bend. You have partial
  charges on two sides of molecule.

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• H bonding gives H2O many of its properties.
• 1) high surface tension
• 2) low vapor pressure
• 3) high specific heat
• 4) high boiling point

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Surface Properties

1. Water molecules experience an uneven attraction
   the molecules are hydrogen-bonded on only one
   side of the drop. The molecules pull toward the
   body of the liquid.

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• 2. This pull is called surface tension.
• 3. A liquid that has strong intermolecular forces
  has high surface tension.
• 4. You can decrease surface tension by adding a
  surfactant, an agent such as soap that interferes
  with the hydrogen bonds.

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Specific Heat Capacity

• water has a high specific heat capacity.
  Ability to absorb heat without changing
  temperatures.

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WATER VAPOR AND ICE

• NOTES 17.2

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Evaporation and Condensation

1. The hydrogen bonds of water helps hold the
   molecules together, and therefore requiring a
   high amount of energy to break the bonds to turn
   to a vapor.

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• 2. The less hydrogen bonding the easier to
  vaporize.
• 3. The reverse of evaporation is condensation,
  when water condenses it releases energy (heat).
• 4. Temperatures in the tropics would be much
  higher if water didnt absorb heat.

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• 5. Temperatures in the polar regions would be
  much lower if water vapor did not release heat
  when condensing out of the air.

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Ice

• A typical liquid cools, it contracts slightly.
  Its density increases because its volume
  decreases. The solid would sink because its
  density is higher than the liquid.

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• As water cools it acts like a typical liquid,
  until it reaches
• 4 oC then the density begins to decrease. As
  Ice forms at 0 oC the volume expands and it has
  lower density than the surrounding water.

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AQUEOUS SOLUTIONS

• NOTES 17.3

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Solvents and Solutes

1. Chemically pure water never exists in nature,
   because water dissolves so many substances
   (Universal solvent).
2. Water samples containing dissolved substances are
   called aqueous solutions.

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• In a solution, the dissolving medium is the
  solvent.
• In a solution, the particles dissolved are the
  solutes.
• 3. Solutions are homogeneous mixtures.

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• 4. Solvents and solutes can be solids, liquids,
  and gases.
• 5. Ionic compounds and Polar covalent molecules
  dissolve most readily in water, but nonpolar
  covalent do not.

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The Solution Process

1. As you place a solute in a solvent the particles
   begin to collide with one another. The solvent
   attracts the solute particles until substance is
   dissolved.

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• 2. In some ionic compounds, the solvent cant
  break the ionic bonds and the salt doesnt
  dissolve.
• 3. Polar solvents dissolve ionic and polar
  molecules, nonpolar solvents dissolve nonpolar
  compounds.

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Solute added to solution

• Raises boiling point
• Salt in water
• Lowers freezing point
• Salt on road

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Electrolytes Nonelectrolytes

• 1. Compounds that conduct an electric current in
  aqueous solution or molten state are
  electrolytes.
• - All ionic compounds are electrolytes.

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• Compounds that do not conduct electric current
  are nonelectrolytes.
• - Most molecular compounds and compounds of
  carbon are nonelectrolytes.

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• 3. Some very polar molecular compounds are
  nonelectrolytes in pure state, but electrolytes
  in an aqueous state.
• 4. You can have strong or weak electrolytes.
  Depends on how well the solute dissolves into
  ions.

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Water of Hydration

1. Water molecules are an integral part of crystal
   structure this is called water of hydration.
   Also, called a hydrate.
2. Effloresce ? the ability to lose the water
   hydration.

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• 3. Hygroscropic ? the ability to remove water
  from the air.
• a. These are used as drying agents (desiccants).
• Ex. Silica gel
• b. Agents that became wet from solutions from
  H2O in air when exposed to air are deliquescent.

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HETEROGENEOUS AQUEOUS SYSTEMS

• NOTES 17.4

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Suspensions

• mixtures from which particles settle out of
  solution upon standing.
• Differs from solution because component parts are
  much larger.

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Colloids

• contain particles that are intermediate in size
  between suspensions and solutions.
• 1. The properties of colloids differ from both
  suspensions and solutions.

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• 2. Colloids are cloudy in appearance when
  concentrated but clear to almost clear when
  diluted.
• 3. Particles do not settle of a mixture.
• 4. Colloids exhibit the Tyndall Effect, the
  scattering of visible light.

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• Emulsions ? dispersions of liquids in liquids.
  An emulsifying agent is essential for the
  formation of an emulsion.
• (Ex. Mayo ? vinegar, oil, and egg)

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PROPERTIES OF SOLUTIONS

• NOTES 18.1

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Solution Formation

• 1. Solutions are homogeneous mixtures and can be
  solids, liquids or gases.

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• Factors that affect how fast a substance
  dissolves.
• a. agitation
• b. temperature
• c. surface area ? the smaller the particle, the
  faster it dissolves.

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Solubility

• the amount that dissolves in a given quantity of
  a solvent at a given temperature to produce a
  saturated solution.
• Particles can move from solid to a solvated state
  and back to a solid again.

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• This is a saturated solution (contains the
  maximum amount of solvent.)
• A solution that contains less solute than a
  saturated solution is unsaturated.

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• 2. Two liquids are said to be miscible if they
  dissolve in each other.
• (ex. Water ethanol)
• 3. Liquids that are insoluble in each other are
  immiscible
• (ex. Oil Vinegar)

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Factors Affecting Solubility

• 1. Temperature
• a. for most solids as temperature increases
  solubility increases.
• b. For most gases as temperature
  decreases solubility increases.

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• 2. Pressure
• a. Gas solubility increases as the partial
  pressure of gas above the solution increases (Ex.
  Carbonated drinks ? contain dissolved CO2 in H2O)
  and decreases as pressure decreases.

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• 3. Supersaturated solution ? contains more solute
  than it should theoretically continue to hold.
• a. Crystallization of the solution can occur by
  adding a small crystal (seed crystal).

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CONCENTRATIONS OF SOLUTIONS

• NOTES 18.2

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Molarity

• 1. Concentration is a measure of the amount of
  solute that is dissolved in a given quantity of
  solvent.
• a. dilute solution ? low concentration
• b. concentrated ? high concentration

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• 2. Molarity (M) ? the number of moles of a solute
  dissolved per liter of solution.
• a. also known as molar concentration and read as
  molar.
• b. To calculate the molarity of any solution -
  calculate the number of moles in 1 L of the
  solution.

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• Molarity moles of solute
• Liters of solution

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• 3. Making Dilution
• a. by adding solvent to a solution you can lower
  its molarity.
• 1) moles of solute do not change.

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• 4. Percent Solutions
• a. If both solute and solvent are liquids, a
  convenient way to make a solution is to measure
  volumes.
• 1)If 20 mL of pure alcohol is diluted with
  water to a total volume and 100 mL the final
  solution is 20 alcohol by volume.

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• Percent volume
• Volume of Solute x 100
• Solution Volume

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COLLIGATIVE PROPERTIES OF SOLUTIONS

• NOTES 18.3

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Decrease in Vapor Pressure

1. Properties of solutions differ from those of the
   pure solvent.
2. Properties that depend on the number of particles
   dissolved in a given mass of solvent are called
   colligative properties.

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• 3. 3 Important colligative properties of
  solutions
• a. vapor pressure lowering
• b. boiling point elevation
• c. freezing point depression

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• 4. A solution that contains a nonvolatile solute
  always has a lower vapor pressure than the
  solvent.
• (Volatile ? easily vaporized.)

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Boiling Point Elevation

1. By adding a nonvolatile solute would increase the
   boiling point.
2. Attractive forces occur between the solvent and
   solute therefore you need more energy to overcome
   these forces.

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Freezing Point Depression

1. When a substance freezes the particles of the
   solid take on an orderly pattern. The presence
   of a solute disrupts this. Therefore more energy
   must be withdrawn.
2. The more solute you add the lower the freezing
   point. (Ex. Salt and water)

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Percent by Mass
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• A solute in solution is the number of grams of
  solute dissolved in 100 g of solution.
• Percent by Mass
• __Mass of solute__ x 100
• Mass of solute mass of solvent

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Example

• by mass
• 10 g NaOH___ x 100
• 10 g NaOH 90 g H2O
• 10 NaOH

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Molarity (M)

• Number of moles of solute in one liter of
  solution.
• Molarity
• of moles of solute of liters of solution

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Example

• If 0.500 moles of NaOH is dissolved in 1.00 L of
  solution, what is the molarity that is produced?
• M 0.500 mole NaOH
• 1.00 L
• 0.500 M NaOH

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• Moles of solute
• M1 x V1 M2 x V2

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Molality (m)

• Concentration of a solution expressed in moles of
  solute per kilogram of solvent.
• Molality of moles solute__
• Mass of solvent (kg)

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Example

• If 8.50 g of ammonia, which is one half mole of
  ammonia, is dissolved in exactly 1 kg of water,
  what molality is produced?
• m 0.500 mole NH3
• 1 kg H2O
• 0.500 m NH3