Chapters 17 and 18: Water and Aqueous Solutions

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Chapters 17 and 18 Water and Aqueous Solutions

• A. The Water Molecule (Ch. 17)
• H2O is polar (has dipoles)
• ex.
• H2O is bent due to its lone e pairs
• H2O covers about 75 of the earths surface

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• Intramolecular bonding bonding within a
  molecule
• ex. covalent bonds
• Intermolecular bonding bonding between
  molecules
• ex. hydrogen bonds
• Hydrogen bonding when a hydrogen atom in one
  molecule bonds with a very electronegativ
  e atom (F, O, N) of another molecule
• ex. H2O

3

• Solvation how H2O molecules arrange
  themselves around ions that are dissolved
  in H2O
• ex.

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• Properties of H2O
• All due to hydrogen bonding (strong)!
• High surface tension
• High specific heat
• High heat of evaporation
• High boiling point

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• High surface tension the inward force (or pull)
  that minimizes the surface
  area of a liquid
• ex.

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• High surface tension (cont.)
• Surfactants decreases surface tension by
  getting (H2O) molecules to relax
  and spread out
• ex. soaps or detergents

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• High specific heat the amount of heat needed
  to raise the temperature of 1 g of a
  substance 1 ?C
• - it takes more energy (heat) to
  heat up H2O compared to other substances
• ex. H2O has a higher specific heat than sand
  and metals
• This explains why sand and metals get hot
  quickly when in the same amount of sun
  (energy/heat), as compared to H2O

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• High heat of evaporation it takes a high amount
  of heat to evaporate H2O molecules
• - evaporation is the same as
  vaporization (l ? g)
• - condensation (g ? l)
• When H2O vaporizes, bonds are broken
• H2O has strong hydrogen bonds, which take a lot
  of heat/energy to break
• High FPs/MPs/BPs vs. other covalent bonds

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• Structure of H2O(l) vs. H2O(s)
• ex. H2O(l) vs. H2O(s)
• Liquid molecules are closer Solid farther apart
• Ice has a bigger volume for the same mass
• Ice has a smaller density (D m / v)
• Ice floats on liquid water

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• B. Solubility of Solutions (Ch. 18)
• Solute the substance being dissolved (s)
• Solvent what the solute is being dissolved in
  (l)
• - usually H2O
• Solution when the solute dissolves in the
  solvent (aq)
• ex. NaCl dissolves in water
• NaCl(s) solute
• H2O(l) solvent
• NaCl(aq) solution

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• Solubility the amount of solute that can
  dissolve in given amount of solvent at a
  given temperature (see solubility curves)
• Miscible liquids that dissolve each other
• Immiscible liquids that do not dissolve each
  other
• Likes dissolve likes
• Polar/Polar dissolves
• Non-Polar/Non-Polar dissolves
• Polar/Non-Polar does not dissolve

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• Factors affecting (to increase) solubility
• Stirring or shaking
• Increase the temperature
• Except gases (dec. temp to inc. solubility)
• Increase the surface area (smaller particles)
• Increase the amount of solvent (H2O)
• Solubility curve a graph that shows the
  different solubilities of many
  substances at different temperatures
  (see worksheet)

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• Saturated solution contains the max. amount of
  solute per solvent at a given temp.
• - on the curve/line
• Unsaturated solution contains less than the
  max. amount of solute per solvent
  at a given temp.
• - below the curve/line
• Supersaturated solution contains more than the
  max. amount of solute per solvent at
  a given temp.
• - above the curve/line

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• Solubility Curve
• ex.

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• C. Molarity (Ch. 18)
• Molarity (M) the concentration of a solution
• - formula M moles / L
• - units moles/L or M (molar)
• - shows how dilute (small amount of
  solute per solvent) or concentrated (large
  amount of solute per solvent)
• - chemists need to know how much
  solute is in a solution for a reaction
• - average solutions are 1 M

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• ex. What is the molarity of a solution that
  contains 2 moles of sugar in 5 liters of water?
• M moles / L
• M 2 moles / 5 L
• M 0.4 moles/L or 0.4 M
• ex. What is the molarity of a solution that
  contains 2 moles of NaCl in 5 liters of water?
• M moles / L
• M 2 moles / 5 L
• M 0.4 moles/L or 0.4 M

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• ex. What is the concentration of a solution that
  contains 0.90 grams NaCl in 100 mL of H2O?
• M moles / L
• 0.90 g NaCl x 1 mole 0.015 moles
• 58.5 g
• 100 mL x __1 L__ 0.1 L
• 1000 mL
• M 0.015 moles / 0.1 L
• M 0.15 M

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• Solving for other variables (moles or liters)
• M moles / L
• moles M x L
• L moles / M
• ex. How many grams are in 1500 mL of a 0.24 M
  solution of sodium sulfate?
• moles M x L (0.24 M) x (1.5 L) 0.36
  moles
• 0.36 moles Na2SO4 x 142.1 g 51.16 g Na2SO4
• 1 mole

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• D. Dilutions (Ch. 18)
• Dilutions make the solution less concentrated
• - make the molarity lower
• - formula M1 x V1 M2 x V2
• - the units of both volumes (V1 and V2)
  must be the same
• - cannot take out dissolved solute from
  a solution so add more solvent (H2O)
• - chemists make a stock solution of a
  high concentrated solution from a high M,
  they can add a specific amount of H2O to
  dilute the solution to the exact M they want

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• ex. What is the new molarity of 100 L of solution
  that was diluted from 50 L of a 5 M solution?
• M1 x V1 M2 x V2
• (5 M) x (50 L) M2 x (100 L)
• 100 L 100 L
• M2 (5 x 50) / 100
• M2 2.5 M

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• ex. How many mL of a 3.0 M solution are needed
  to prepare 1 L of a 1.5 M solution ofMgSO4?
• M1 x V1 M2 x V2
• (3.0 M) x V1 (1.5 M) x (1000 mL)
• 3.0 M 3.0 M
• V1 (1.5 x 1000) / 3
• V1 500 mL

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• To how muchbe added problems
• Solve for V, but subtract the 2 volumes to get
  the volume that should be added
• ex. To how much water should 25 mL of 8 M HCl be
  added to produce a 2 M solution?
• V2 (25 x 8) / 2 100 mL V2
• - 25 mL V1
• Answer 75 mL must be added

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• E. Colligative Properties (Ch. 18)
• Colligative properties depend on the number of
  particle units that are
  dissolved in a solvent
• Particle unit the number of units that each
  substance breaks up into when dissolved
• Covalent compounds 1 particle unit
• ex. sugar (C12H22O11) 1 unit (no ions)
• Ionic compounds of total ions
• ex. NaCl 2 units (1 Na, 1 Cl)
• MgBr2 3 units (1 Mg, 2 Br)
• Al2O3 5 units (2 Al, 3 O)

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• Effects of the solute and solvent
• Solute
• The more moles of solute, the greater the effect
  on the colligative property
• The more particle units (assuming same number of
  moles), the greater the effect on the colligative
  property
• Solvent
• The higher the value of the solvents constant,
  the greater the effect on the colligative
  property (water has a very low constant)

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• Types of colligative properties
• Boiling point elevation
• The boiling point is higher with added solute
  when compared to the pure solvent (just by
  itself)
• ex. Adding salt to water when boiling water
  to cook with (a myth?)
• Freezing point depression
• The freezing point is lower with added solute
  when compared to the pure solvent (just by
  itself)
• ex. Rock salt on icy roads

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• F. Suspensions and Colloids (Ch. 17)
• Heterogeneous solutions not uniform (not the
  same throughout)
• - 2 different phases
• - show the Tyndall effect
• Tyndall effect the scattering of light
• ex.

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• Suspensions particles that settle out upon
  standing (stay at the bottom)
• - large particles (gt 100 nm)
• - particles are visible
• - particles can be filtered out
• - exhibits the Tyndall effect
• ex. Pieces of clay in water
• Snow globes

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• Colloids particles that do not settle out upon
  standing
• - small particles (1-100 nm)
• - particles are not visible
• - particles cannot be filtered out
• - exhibits the Tyndall effect
• - substances contain 2 states of matter
  that are dispersed evenly

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• ex. Colloids
• (p. 491)

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• Solutions particles that do not settle out upon
  standing
• - very small particles (lt 1 nm)
• - particles are not visible
• - particles cannot be filtered out
• - do not exhibit the Tyndall effect
• - substances are homogeneous
• - look at p. 492 for a comparison of
  solutions, colloids, and suspensions