Ions in Solution
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Transcript of Ions in Solution
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Ions in Solution
Chapter 14
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I. Ionic Compounds in Aqueous Solution (Aqueous - water is solvent) A. Theory of Ionization 1. Faraday - current causes ions to form a. Electrolytes b. Nonelctrolytes 2. Arrhenius - ionization of molecules
in water produces ions
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1. The solution process for ionic compounds
a. Hydration - solution process with water as solvent
b. Factors affect # of water molecules needed for hydration:
1) size of ion 2) charge of ion
B. Dissolving Ionic Compounds
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2. Heat of solution for ionic compounds heat of hydration - energy released
when ions become surrounded by water
a) Exothermic - releases heat ; negative heat of solution
b) Endothermic - absorbs heat ; positive heat of solution
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3. Dissociation - separation of ions when an ionic compound dissolves
NaCl ---> Na+(aq) +Cl-(aq)
1 mol 1mol 1 mol CaCl2 ---> Ca +2
(aq) + 2Cl -(aq)
1 mol 1mol 2 mol
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C. Ionic Equations and Precipitation Reactions
1. Reactions in Solution a. Precipitate (ppt) - insoluble
substance formed through a chemical reaction in a solution
b. Some double replacement reactions produce ppt; others form a gas or water.
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c. Solubility Table 1. i - insoluble - forms a ppt 2. ss - slightly soluble - formation
of a slight ppt 3. s - soluble - no ppt forms
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2. Writing Ionic Equations
a. Write formula for compound.
sodium chloride = NaCl b. Write the compound as ions:
NaCl becomes Na+ + Cl-
c. Check solubility table to determine if a ppt forms
d. If all combinations give ‘s’ - reaction is NR
e. If one combination gives either ‘i’ or ‘ss’ - then a reaction takes place
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e. Overall ionic equation includes all ions those that form a ppt and those that are referred to as ‘spectator ions’ because they do not form a ppt
f. Net ionic equation includes only those ions that form a ppt; cancel out the spectator ions on both sides of the equation.
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Examples:
1. Write the overall ionic equation and the net ionic equation that occurs when aqueous solutions of zinc nitrate and ammonium sulfide are combined.
2. A solution of sodium sulfide is combined with a solution of iron(II) nitrate. Write the net ionic equation for any reaction that occurs.
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II. Molecular Electrolytes (Polar covalent molecules can form electrolytes)
A. The solution process for molecular electrolytes
1. Polar molecules in water - opposite dipoles attract - if strong enough bond breaks and the molecule is separated into simpler charged parts
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2. Ionization - formation of ions from solute molecules by the action of the solvent
[Dissociation: ionic compounds ---- Ionization: polar compounds]
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B. The Hydronium Ion
1. H+ is only a proton, smaller than any other ion - it is attracted to others so strongly it does not have any independent existence
2. H + + H2O ---> H3O +
hydrogen ion water hydronium ion
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C. Strong and Weak Electrolytes
1. Strong - 100% ions 2. Weak - low concentration of ions
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III. Properties of Electrolyte Solutions
A. Conductivity of Solutions 1. Strong-weak: degree of ionization 2. Concentrated-dilute; amount of
solute-solvent 3. Ionization of H2O
2H2O ---> H3O+ + OH-
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B. Colligative Properties of Electrolyte Solutions
1. Electrolytes affect colligative properties more than nonelectrolytes
Example: Compute the bp and fp for a solution made by adding21.6 g of NiSO4 to100 g of water.
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2. Theory vs Reality a)Theory - electrolytes reduce fp by 2,3
times - depending on # of ions b) Reality - reduces more than
nonelectrolytes , but not as much as predicted c) Reason - because ions are attracted to
each other in water - more concentrated solutions have higher attraction for each other because they are closer together
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3. “Ideal Solution” - dilute enough that the ions have the expected activity
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IV. Colligative Properties of Solutions
A. Definition - a property that depends on the number of solute particles but is independent of their nature 1. Nonelectrolytes - 1 solute particle 2. Electrolytes - # of solute particles
dependent on # ions
• NaCl: 2 AgNO3: 2
• MgCl2: 3 K3PO4: 4
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1. Vapor Pressure Lowering - the tendency for molecules to escape from a liquid to a gas is less in a solution than a pure solvent
2. Freezing Point Depression - solution has a lower fp than solvent ∆ tf = Kfm
∆ tf- freezing point change Kf- molal freezing point constant m - molality of the solution
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Example: What is the fp of water in a solution of 17.12 g C12H22O11 and 200 g of water?
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3. Boiling Point Elevation - solution has a higher bp than solvent
∆ tb = Kbm
∆ tb - change bp Kb - molal boiling pt constant m - molality
Example: What is the bp of a solution that is made by adding 20 g C12H22O11 in 500g H20?
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C. Determination of Molar Mass of a Solute
1. Determine Δtf(Δ tb)
2. Determine m Δ t = Km
3. If ionic divide by number of particles4. Calculate moles of solute
m X kg of solvent5. Molar mass = mass of solute
moles of solute
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Example: When 1.56 g of an unknown , nonelectrolyte solute is dissolved in 200 g H2O, the ∆ tf = -0.453 Co. Determine the molar
mass.