Plan, Ppt 13: Substances in Aqueous Solution
(PS5, 19-22 and PS6, 1-7 material)
1. Reminder: Ionic vs. Molecular (non-acid) vs.
Acid
2. Molarity of solutes I
3. Qualitative Issues (types of solutes/solutions)
• Electrolytes (Operational)
• Electrolytes (Conceptual)
• Strong vs. Weak vs Non Electrolyes
4. Molarity Reprise (electrolyte issues)
5. Solubility rules for ionic compounds
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Reminder—Ionic Compounds
• Metal or “NH4” listed first in formula
• Made up of ions (cation and anion type)
• You need to be able to:
- Write the formulas of the cation and anion
- State the number of each ion present in one FU of the compound
• Example: Na3PO4
- Made up of Na+ ions and PO43- ions
- 3 Na+ ions / FU; 1 PO43- / FU
KNOW YOUR
IONS!! (PS3)
• NOTE: NaOH, KOH, etc. are ionic compounds.
Because of the OH- ion (not the cation!), these ionic
compounds are also called “bases”.
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Reminder—Molecular Compounds
• In this class, a compound that is not ionic!
• NO metal first; a nonmetal (unless “NH4”)
• Learned to name binary molecular compounds…
 Non-acids (e.g., sulfur trioxide, SO3), and
 (binary) acids (e.g., hydrochloric acid, HCl)
• …and one class of ternary molecular compounds
 (ternary) acids (e.g., chloric acid, HClO3)
• But there are many other kinds
 Organic compounds:
o C6H12O6 (glucose); CH3COCH3 (acetone);
C8H8O4 (aspirin); most drugs, actually
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Reminder—Acids
(a subset of molecular compounds)
• H is listed first in formula (for acids in this class [simple])
• Made up of molecules, BUT….
• …can be thought of as being made by taking any anion
and adding H+’s until a neutral FU results:
- H+’s & PO43-  H3PO4 (phosphoric acid)
- H+’s & S2-  H2S (hydrosulfuric acid)
KNOW YOUR
(AN)IONS!! (PS3)
• You need to be able to:
- Determine the anion from which the acid is derived
• H+ is often called a “proton” (since an H atom has 1 p and 1 e-; if
you remove the e-, you get H+ [1 p only])
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A solution is a homogeneous mixture
• Solvent is often a liquid
• Solute is “what’s dissolved in” the liquid
– Solute can be solid, liquid, or gas at room temp.
– Solute can be a molecular compound (non-acid or
acid), or an ionic compound
• If solvent is water, solution is called “aqueous”
[indicated by (aq) after solute]
– E.g., NaCl(aq)
• Dissolution is what happens as a solute
dissolves in a liquid
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Operational Definition, Dissolution
• Dissolution: Formation of a homogeneous
mixture (a “solution”)
– one substance “dissolves” in another
– A “solution” is something you can see through (it
is clear, not cloudy) You can assess
• Is milk a solution?
“dissolution”
• Remember the “waviness” in lab
visually.
– If something “dissolves” in a liquid (to any
appreciable extent), it is said to be soluble.
• Some substances are soluble in water, some insoluble
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“Theory” note:
Warning: Dissolution is not the same thing
as Dissociation!!
• Discussion of “dissociation” will come shortly.
• For now just note that dissociation refers to a
process that occurs after dissolution.
– As such, in contrast to dissolution:
You cannot assess
“dissociation” visually.
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Molarity (Worksheet)
Try these now!
a) If 3.12 g of KF is dissolved in some water to
make the final volume equal to 50.0 mL,
what is the molarity of this solution?
•
(using P.T. info: MM of KF = 58.10 g/mol)
b) How many moles of sugar (C6H12O6) are in 5.4
L of a 0.15 M solution?
c) How many liters of a 2.5 M solution of KBr
will contain 3.6 moles of KBr?
PS5, problems19-21 reflect this material
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NOTE: Once molarity is understood,
stoichiometry problems can utilize it!
• “Another way to get to moles” (mol/L x L)
• “Another way to go from moles” (mol  L = M)
• PS6, Q2; Question (d) on Molarity Worksheet:
Plan: mL HCl  L HCl  mol HCl  mol Na2O  g Na2O
Plan: mL HCl  L HCl  mol HCl  mol Na2O  M Na2O
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NOTE: Once molarity is understood,
stoichiometry problems can utilize it!
• More example problems of this type will be
shown and discussed in Ppt 16b (as a
summary application of ideas—titration as an
application of stoichiometry with molarity)
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Back to the qualitative discussion of
solutions
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Some solutes form solutions that conduct
electricity and some do not!
Operational Definition of “Electrolyte”
• Electrolyte: a substance that is soluble in water and
produces a solution with increased electrical
conductivity
– **Can test with a light bulb apparatus or conductivity
meter**
• By definition, electrolytes are a subset of all soluble
substances.
– All electrolytes are soluble, but not all soluble substances
are electrolytes!!
• Soluble substances that do not produce a solution with increased
electrical conductivity are nonelectrolytes
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Some examples
Electrolytes
All soluble ionic compounds
e.g., NaCl, NaNO3, FeCl3,
NiSO4, Ca(OH)2, etc.
All Acids and Molecular Bases:
e.g., HCl, HNO2, HClO4,
NH3, etc.
Nonelectrolytes
Most soluble molecular
substances that are not acids or
bases
• CH3COCH3 (acetone)
• CH3OH (methanol)
• O2
• C12H22O11 (sucrose)
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Conceptual Definition of Electrolyte
(theory)
• Electrical current is “moving charge”
– Need to have “mobile charges” to increase
electrical conductivity
• Electrolyte: a substance for which at least
some of its dissolved formula units end up
turning into ions
– Process is called dissociation(or “ionization”)
– Ions in solution are free to move around, hence
increasing electrical conductivity!
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What happens during dissolution?
(theory)
• See next slide for pics that show the following:
– Molecules will separate from one another, but
generally remain intact (if substance dissolves)
• Too small to scatter light, that’s why solution is “clear”
– Ions will separate from one another once
dissolved (if substance dissolves)
• Too small to scatter light, that’s why solution is “clear”
• Ionic compounds that dissolve will also undergo
dissociation
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Dissolution only
PS6, problem
3 reflects this
material
Dissolution &
dissociation
**Note that molecules remain intact during dissolution; they just
separate from one another (without dissociating)!
**Note that each individual cation (Na+) and anion (SO32-) remain
intact during dissolution, but the ions separate from one another
Dissolution is not the same as Dissociation
(reprise)
• Dissolution applies to any substance that
dissolves, whether molecular or ionic
• Dissociation (in aqueous solution) applies only to
electrolytes; it refers to the production of ions
in solution, after FUs dissolve.
Recall: You cannot assess
“dissociation” visually. Need
“light bulb” or conductivity meter.
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Dissolution ≠ Disscociation
(continued)
• Most molecular substances’ FUs do not
dissociate after they have dissolved
– I.e., Most molecular substances are not
electrolytes
• But FUs of ionic compounds that dissolve DO
also dissociate
– I.e., Soluble ionic compounds are electrolytes
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Electrolytes come in two “types”
• “strong”: ~100% of dissolved FUs are ionized
• “weak”: significantly less than 100% are ionized
• NOTE: “nonelectrolyte” refers to a soluble substance
that is not an electrolyte
– Thus an insoluble substance is technically neither a strong,
weak, nor nonelectrolyte
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Analogous to Fig. 4.14 in Tro. Comparison of Strong,
Weak, and Nonelectrolyte solutions (operationally and
conceptually)
All three
substances are
SOLUBLE, but they
differ in the
fraction of FUs that
exist as separated
ions:
Strong
electrolyte
Weak
electrolyte
Nonelectrolyte
(a) 100% of FUs
are ionized
(b) only a small
fraction of FUs
are ionized
(c) NONE are
ionized
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The Six Common Strong Acids
(memorize)
• HCl
• HBr
• HI
• HNO3
• H2SO4
• HClO4
All others are
weak acids!
• HF
• HNO2
PS5, problem
22 reflects this
• H3PO4
material
• H2CO3
• HClO
• HC2H3O2
• Etc.
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Analogous to Figure 4.10 in Tro. Polar Water
Molecules Interact with the Positive and Negative Ions
of a Salt Assisting in Dissolution (and Dissociation!!)
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NaCl Dissolves (and Dissociates)
(and 100% of FUs separate in solution)
Strong
electrolyte
NaCl(s)  Na+(aq) + Cl-(aq)
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(soluble ionic
compound)
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Analogous to
image on p.
148 in Tro.
HCl is
Completely
Ionized
(Dissociated)
strong electrolyte
[one of the six strong
acids]
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Analogous to
image on p. 148
in Tro.
Acetic Acid
(HC2H3O2)
(weak acid [not one
of the six])
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An
Aqueous
Solution of
Sodium
Hydroxide
(strong electrolyte
[soluble ionic
compound])
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Figure 4.9 (Zumdahl)
A solution of
NH3 in Water
NOTE: This weak
electrolyte does not
technically “dissociate”
to make ions! It
produces ions by
reacting (a bit) with
water molecules:
NH3(aq) + H2O(aq)
NH4+(aq) + OH-(aq)
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Demonstration (if not already done)—
Dissolution is not the same as Dissociation!
• To see if a substance dissolves (e.g., in H2O):
–
–
–
–
Put a small amount of solid in a large test tube
Add water until about 1/3 full
Swirl for awhile: waviness indicates dissolution, and…
…if tube ends up CLEAR (not cloudy or opaque), then all of
the substance dissolved, and the substance is said to be
“soluble”
• If a solution is clear, you cannot tell whether or not is
contains an electrolyte just by looking at it!
– You must test for electrical conductivity (light bulb lights)
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Connecting nanoscopic pictures to
“molarity”
• Molar concentration is like “number density”
– Number of FU’s (regardless of mass) in a given amount
of space
• Molarity = moles of solute per liter of solution
– 0.15 M Na2SO3 means:
• 0.15 moles of Na2SO3 (FUs) per liter of solution,
regardless of the fact that Na2SO3 is a strong electrolyte
– 1.5 M acetone means:
• 1.5 moles of acetone (molecules) per liter of solution
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Molarity of Electrolytes
(remember the meaning of subscripts!)
• What is the concentration of Na+ ions
in a 2.0 M solution of Na2SO3?
Consider: Na2SO3(s)  2 Na+(aq) + SO32-(aq)
2 mol Na2 SO3
2 mol Na+
4 mol Na+
x

 4 M Na +
L
mol Na2 SO3
L
• & the concentration of SO32- ions in the same solution?
2 mol Na2SO3
1 mol SO3 22 mol SO3 2x

 2 M SO3 2L
mol Na2SO3
L
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Molarity of Electrolytes
(remember the meaning of subscripts!)
• [Cl-] in 0.50 M FeCl3?
Consider: FeCl3(s)  Fe3+(aq) + 3 Cl-(aq)
3 mol Cl0.5 M FeCl3 x
 1.5 M Cl mol FeCl3
PS6, problems 4-5 reflect this material
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Some ionic compounds are soluble,
but some are not!
• Those that are are strong electrolytes
• Those that are not, are obviously not strong
electrolytes (their formula units never even
get into solution!)
• Look for patterns:
– NaCl, Na2S, NaClO4, Na3PO4, NaHCO3, Na2C2O4,
and Na2Cr2O7, NaOH, Na2CO3 are all soluble
• Tentative conclusion?
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Some ionic compounds are soluble,
but some are not!
• Look for patterns:
– NaNO3, Ba(NO3)2 , Ru(NO3)2 , Fe(NO3)3, Pb(NO3)4,
Ni(NO3)2 , CuNO3 , and Cu(NO3)2 are all soluble
• Tentative conclusion?
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Common Solubility Rules
(You need to memorize only 1 & 2; 3-6 will be given on Exam 2b)
1. All (common) nitrate salts (salts that contain NO3- as the anion) are soluble.
2. All (common) salts containing an alkali metal ion (Li+, Na+, K+, etc.) or ammonium
ion (NH4+) as the cation are soluble.
3. Most (common) salts containing Cl-, Br-, or I- as the anion are soluble. However,
important exceptions are those containing Ag+, Pb2+, or Hg22+. (i.e., AgBr and
PbCl2 are insoluble.)
4. Most (common) salts containing sulfate (SO42-) as the anion are soluble. However,
important exceptions are CaSO4, SrSO4, BaSO4, PbSO4, and Ag2SO4.
------5. Most (common) salts containing hydroxide(OH-) as the anion are insoluble.
However, Ca(OH)2, Sr(OH)2 , and Ba(OH)2 are slightly soluble.
6. Most (common) salts containing carbonate (CO32-), phosphate (PO43-), chromate
(CrO42-), and sulfide (S2-), as the anion are insoluble. However, CaS,, SrS , and
BaS are soluble and Rule #2 still follows (e.g., Na2CO3 is soluble)
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Table 4.1 Partial, “Reformat”
PS6, problem 7 reflects this material (but used later as well)
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