Download Chapter 4: Aqueous Reactions and Solution

Chapter 4: Aqueous Reactions and Solution
Stoichiometry
The topics in this chapter will further our knowledge of types of chemical
reactions and our abilities to predict the products of and write balanced
chemical equations for a variety of chemical reactions. We will also review
solution terminology, properties of solutions and the concept of molarity.
We will then be able extend our ability to do stoichiometric calculations to
those involving reactions that take place in solutions. Our goals are to:
1.
2.
3.
4.
5.
6.
develop an understanding of and the ability to describe the
nature of electrolytes and non-electrolytes in aqueous solutions.
become proficient at recognizing reaction types and be able to
predict products for common chemical reactions: precipitation,
acid-base and simple oxidation-reduction.
develop an understanding of and learn to use the activity series
to predict products of single replacement redox reactions.
become proficient at writing net-ionic equations.
understand molarity (review): solution preparation and dilutions.
become proficient at stoichiometric calculations involving
solutions.
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A REVIEW of solution terminology and symbols:
Solution: A homogeneous mixture containing a solute dissolve in a
solvent. For aqueous solutions, water is the solvent.
In a balanced equation we use phase symbols to indicate the physical
state of the reactants/products. There are 4 phase symbols:
1. (s): solid
2. (l): liquid
3. (g): gas
4. (aq): aqueous. This means dissolved in water to form a solution.
By definition any soluble compound will have the (aq) phase
symbol in a reaction that takes place in water.
Soluble = (aq)
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All Soluble Compounds have Electrolytic Properties!
Non electrolyte
Weak electrolyte
Strong electrolyte
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Details: Strong Electrolytes in Water
• Aqueous solutions of these conduct electricity well.
 These are compounds that separate completely (or almost completely)
into ions when dissolved in water.
(Exist as ≈ 100% ions in solution)
They include:
1. All soluble ionic compounds, including the soluble strong bases: These
substances dissociate (separate) into their ions in water.
1. We will explore the general solubility guidelines for ionic compounds in lab.
You will be expected to memorize a list of these to be handed out later.
2. Strong soluble bases: (MEMORIZE) NaOH, LiOH, KOH, CsOH, RbOH,
Ba(OH)2, Ca(OH)2 and Sr(OH)2.
2. All strong acids: These are molecular substances that ionize completely in
water. (MEMORIZE)
1. HCl
2. HBr
3. HI
4. HNO3
5. HClO4
6. H2SO4 (Only one H is ionized completely. We will explore this later.)
7. HClO3 (included in your text as a 7th strong acid)
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Let’s explore the special properties of water that
result in the “Hydration of Ions”
The ions from electrolytes become hydrated (surrounded by
water molecules) as they move into the water.
KMnO 4 (s) ⎯ ⎯2⎯→ KMNO 4 (aq) = K + (aq) + MnO 4 − (aq)
H O
Text Question 4.13
Many ionic solids dissolve in water as strong electrolytes, that is, as
separated ions in solution. What properties of water facilitate this process?
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Details: Weak Electrolytes in Water
• Aqueous solutions of these conduct electricity poorly
compared to strong electrolyte solutions.
 These are compounds where a relatively small percentage separate into
ions when dissolved in water. These are molecular substances that
ionize partially in water. About 5% (or less) of the molecules separate
into ions, most remain in the molecular form.
They include:
A. All weak acids (For Chem 1A, any acid that is NOT strong). A few examples
are:
1. HCN
2. HF
3. CH3COOH
4. Most other acids that contain the carboxylic acid group, –COOH
B. All weak bases
1. NH3
2. Compounds that contain –NHx
Examples: CH3NH2 and (CH3)2NH
Weak electrolytes form a chemical equilibrium in solution:
HCN(aq) ⇌ H+(aq) + CN–(aq) about 1% of the HCN is ionized.
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Weak Electrolytes
Hydrofluoric acid:
Ammonia:
Text Question 4.17
Formic acid, HCOOH, is a weak electrolyte. What solute particles are
present in an aqueous solution of this compound?
Write the chemical equation for the ionization of HCOOH in water. Use
the correct type of arrow in the equation to show equilibrium.
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Details: Non Electrolytes in Water
• Aqueous solutions of these DO NOT conduct electricity
 These compounds are soluble, but do not form ions when
dissolved in water.
Some examples include:
1. Sugars
2. Alcohols (What is the “functional group?)
3. Glycols (ethylene glycol is what we commonly know as antifreeze)
“Give it some thought”: What property of these types of molecules
enables them to be soluble in water? (This will be explored in
great detail in Chemistry 1B.)
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Reaction Types in this Chapter
You must be familiar with and be able to predict products
for the following types of reactions:
1.
2.
3.
4.
5.
Precipitation
Acid-Base Ionization/Dissociation Reactions
Acid-Base Chemical Reactions
Gas Formation
Oxidation-Reduction
You will be expected to be able to write complete (molecular) and
net ionic chemical equations. This will require you to be
disciplined, committing yourself to memorizing certain
information and doing much practice. To be successful you must
develop the ability to recognize an acid, a base and an ionic
compound from their formulas. An understanding of the nature of
strong, weak and non-electrolytes in water is also needed. For
oxidation-reduction reactions, you will need to become proficient
at the use of the activity series.
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Net-ionic Equations (for all reaction types)
• Net-ionic equations are used to describe the actual chemical
species involved in an aqueous reaction. To write net-ionic
equations you must identify ALL the STRONG ELECTROLYTES!
• Steps:
A. Write the “molecular” or “complete” equation:
1. Write the correct formulas for reactants and products with the
correct phase symbols: (s), (l), (g), (aq).
2. Balance the equation.
B. Use the molecular equation to obtain the net-ionic
equation:
1. Identify all the STRONG ELECTROLYTES.
 Search only the (aq) species for the strong electrolytes!
2. Rewrite the strong electrolytes as separated ions. All other
species are rewritten without modification.
3. Cancel any ions that appear on both sides of the reaction.
These are called spectator ions.
4. Rewrite the equation, omitting the spectator ions. Check that
the final net-ionic equation is balanced.
•
With enough practice most students develop the ability to go directly from the
complete (step A2) to the net-ionic equation (step B4).
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1. Precipitation or Exchange (Metathesis) Reactions
•
Precipitation



•
Here an insoluble, ionic precipitate is formed when two solutions
containing dissolved ions are mixed.
A precipitate has the phase symbol (s).
These are sometimes called “double replacement reactions”.
You must know the solubility guidelines to predict products.
Examples
a) AgNO3(aq) + NaCl(aq) —> AgCl(s) + NaNO3(aq)
b) CuSO4(aq) + 2 NH4F(aq) —> CuF2(s) + (NH4)2SO4(aq)
c) 2HI(aq) + Pb(NO3)2(aq) —> PbI2(s) + 2 HNO3(aq)
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*
*Add to list of soluble compounds: Compounds containing Group 1A cations and
compounds containing the ammonium ion. No important exceptions
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Examples: Precipitation or Exchange
(Metathesis) Reactions
1. When aqueous solutions of sodium carbonate and iron (III) chloride are
mixed, a precipitate forms.
What is the identity of the precipitate?
Write the net-ionic equation for the reaction.
2. Suggest two soluble salts that when mixed together in aqueous solution
will produce calcium phosphate.
Write the balanced net-ionic equation for the reaction.
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Examples: Precipitation or Exchange
(Metathesis) Reactions
Text Question 4.21
Will precipitation occur when the following solutions are mixed? If so,
write a balanced molecular chemical equation for the reaction. (Include
phase labels)
a) Na2CO3 and AgNO3
b) NaNO3 and NiSO4
c) FeSO4 and Pb(NO3)2
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Examples: Precipitation or Exchange
(Metathesis) Reactions
Text Question 4.23
Name the spectator ions in any reactions that may be involved when
each of the following pairs of solutions are mixed.
a) Na2CO3(aq) and MgSO4(aq)
b) Pb(NO3)2(aq) and Na2S(aq)
a) (NH4)3PO4(aq) and CaCl2(aq)
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2. Acid-Base Ionization Reactions
• Acid: the simplest definition of an acid is “a compound that
ionizes to form hydrogen ions (H+) in aqueous solution.”
Ionize: The process of forming ions by reacting with water. Acids
are NOT ionic substances, the ions form when the acid is added
to and then reacts with water.
ACID IONIZATION CHEMICAL EQUATIONS:
Example: Hydrochloric acid
HCl(aq) —> H+(aq) + Cl–(aq)
• 100% ionized since HCl is a strong acid/electrolyte
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Acid Ionization
• We write H+(aq) to show the proton is ionized from the acid
and solvated by the water.
• In reality the proton is always attached to water, forming the
hydronium ion, H3O+(aq).
• When written this way, we see that HCl donates a proton to H2O.
Acid: a BETTER DEFINITION is “a proton (H+) donor”.
Nitric acid:
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Weak acid ionization, only a few % of the acid
molecules ionize
Weak acid ionization is written with a double arrow to
indicate less than 100% ionization.
Examples:
Oxalic acid: H2C2O4(aq) ⇄ H+(aq) + HC2O4–(aq)
Acetic acid: CH3CO2H(aq) ⇄ H+(aq) + CH3CO2–(aq)
The strong acids we have written ionization equations for thus far are monoprotic acids.
Let’s consider sulfuric acid, a strong diprotic acid:
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Bases
• Base: the simplest definition of a base is “a compound that
produces hydroxide ions (OH—) in aqueous solution.”
Strong Bases:
These substances dissociate 100% into ions when dissolved
in water. The word dissociate is used instead of ionize since
strong bases are ionic in nature. The ions simply separate from
each other (dissociate) when the substance is dissolved in water.
Example: NaOH(aq) —> Na+(aq) + OH–(aq)
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Weak Bases:
A weak base undergoes a hydrolysis reaction with water. This
reaction produces hydroxide ions. At any given moment,
only a small percent of the base molecules have undergone
hydrolysis. The most common weak base is ammonia.
• In the reaction shown above, NH3 is accepting a proton from
water. In the process, a hydroxide ion is produced.
Base: a BETTER DEFINITION is “a proton (H+) acceptor”.
Ammonia:
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3. Acid-Base Reactions
Often called Neutralization Reactions
Acids: compounds that can donate a proton, H+, to a base.
Bases: compounds that can accept a proton, H+, from an acid.
Reaction: Transfer the proton(s) from the acid to the base to make the
products. (This definition does NOT limit neutralization reactions to water
solutions.)
•

Example reaction for strong bases:
Ba(OH)2(aq) + 2 HNO3(aq) —> Ba(NO3)2(aq) + 2 H2O(l)

Example reaction for weak bases:
NH3(aq) + HNO3(aq) —> NH4NO3(aq)
The insoluble hydroxide salts can also act as bases when in the presence
of an acid. The reaction causes the insoluble base to seem to dissolve. In
reality, this is a chemical reaction, not simply a physical change of
dissolving.

Al(OH)3, Fe(OH)2, Mg(OH)2, etc.

Example reaction:
Mg(OH)2(s) + 2 HCl(aq) —> MgCl2(aq) + 2 H2O(l)
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Examples: Acid-Base (Neutralization)
• To predict the products of an acid-base reaction classify the
base as a hydroxide (contains OH–) or as ammonia or an
amine (contains –NHx):
a) Base contains hydroxide: —> products are water and a salt
(The salt is almost always soluble, but there are exceptions!)
b) Base is ammonia or an amine (–NHx): —> product is a salt
containing an –NH+x+1 ion (no water produced). For example, if
the base is NH3, an ammonium (NH4+) salt is produced.
(Remember, ammonium salts are always soluble.)
Text Question 4.39a&c
Complete and balance the following molecular equations, and then
write the net ionic equation for each:
a) HBr(aq) + Ca(OH)2(aq) —>
c) Al(OH)3(s) + HNO3(aq) —>
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Examples: Acid-Base (Neutralization)
•
Predict the products of the following reactions. Complete and balance
the molecular equations, and then write the net ionic equation for
each. (React all acidic protons on the acid, complete neutralization.)
a) H2SO4(aq) + Fe(OH)3(s) —>
b) CH3COOH(aq) + NH3(aq) —>
c) CH3NH2(aq) + HCl(aq) —>
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4. Gas-Forming Reactions
• There are several reactions that occur between an acid and a
base that produce a gas as one of the products. We call these
Acid-base Reactions with Gas Formation.
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4. Gas-Forming Reactions
• Carbonates (or bicarbonate) + acid —> CO2(g) + H2O(l) + salt
• Sulfides + acid —> H2S(g) + salt
• For these reaction types, questions on exams will involve only
the carbonate/bicarbonate gas forming reactions.
• Predict the products of and balance the following reactions:
(React all acidic protons on the acid, complete neutralization.)
a) HI(aq) + NaHCO3(aq) —>
b) H2SO4(aq) + CaS(s) —>
c) CH3COOH(aq) +Al2(CO3)3(s) —>
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5. Oxidation and Reduction (redox) Reactions
Oxidation and reduction involve the transfer of electrons from
one species to another.
• Oxidation: The loss of electrons
• Reduction: The gain of electrons
One cannot occur without the other!
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Recognizing Redox Reactions: Assigning
Oxidation Numbers to Atoms
I.
For an atom in its elemental form the oxidation number is
zero.
II. For any monatomic ion the oxidation number equals the
charge of the ion.
III. For nonmetals the oxidation number is usually negative.
a) Oxygen is usually -2 in all compounds.
b) Fluorine is -1 in all compounds.
c) Hydrogen is +1 when bonded to nonmetals and -1 when bonded
to metals (metal hydrides).
IV. The sum of the oxidation numbers for all atoms is zero for
neutral compounds or equals the charge for polyatomic ions.
Assign oxidation numbers to each atom in the following:
(a) Na3N
(b) CO2
(c) CH4
(d) Mn(NO3)2
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Example: Redox Reactions
• Determine which of the following reactions are redox:
a) O2(g) + 2 H2(g) —> 2H2O(g)
b) 2KClO3(s) —>2 KCl(s) +3 O2(g)
c) Na2CO3(aq) + CaCl2(aq) —> 2NaCl(aq) + CaCO3(s)
d) CH4(g) + 2O2(g) —> 2H2O(g) + CO2(g)
e) NH3(g) + HCl(g) —> NH4Cl(s)
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One type of Redox Reaction: Oxidation of Metals
• Elemental metals can be oxidized in the presence of a salt or acid.
The general form is:
M + AX —> MX + A where M is a metal and AX is a salt or acid.
• (Often called “single-replacement reactions”)
Examples:
Cu(s) + 2 AgNO3(aq)—> Cu(NO3)2(aq) + 2 Ag(s)
What’s the balanced net ionic chemical equation?
Mg(s) + 2 HCl(aq) —> MgCl2(aq) + H2(g)
What’s the balanced net ionic chemical equation?
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Activity Series of Metals/Ions
Reactivity of Metal Ion Increases
Reactivity of Metal Increases
(Used to predict if a single replacement redox reaction will take place.)
Any metal on the left will
be oxidized (will react)
with any ion on the right
that is below the metal in
the table.
Noble Metals
(don’t react with acids!
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Examples: Oxidation Reduction Reactions
• Predict the products (if any) of the following reactions. If a
reaction occurs, write the net-ionic equation.
a) Mg(s) + FeCl2(aq) —>
b) HCl(aq) + Cu(s) —>
c)
Zn(s) + CuSO4(aq) —>
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Examples: Oxidation Reduction Reactions
Text Question 4.53b&d
Write balanced molecular and net ionic equations for the
reactions of
b) chromium with hydrobromic acid
d) aluminum with formic acid, HCOOH
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Solution Concentration: Definition of Molarity
•
•
•
Molarity is the most common unit of concentration in chemistry.
Molarity is defined as: (moles solute)/(1 L of solution)
Solutions of accurate and precise molarity are made using volumetric
flasks.
1. Calculate the molarity of NiCl2 in a solution prepared
by dissolving 0.435 g of NiCl2(H2O)4 in a 250-mL
volumetric flask and adding water “to the mark”.
a) What are the concentrations of each ion in
solution?
2. How many grams of Ba(OH)2 are needed to prepare a
500.0 mL solution that is 0.012 M in OH– ion?
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Dilution of Solutions
Often solutions must be diluted before they can be used.
Dilution of a 0.0400 M solution
The formula for dilution is:
C1V1 = C2V2 = amount of solute
Where:
C1 is the concentration of the stock solution (before
dilution)
V1 is the volume of the stock solution to be diluted
C2 is the concentration after dilution
V2 is the volume of the diluted solution
Referring to the video, concentration after dilution of
2.00 mL of a 0.0400 M solution diluted to 500.0
mL:
C2 = C1V1/V2 = (0.0400 M)(2.00 mL)/(500.0 mL)
= 1.60x10-4 M
Question: In lab you are asked to prepare 250.0 mL of
a 0.0250 M FeCl3 solution from a 0.520 M stock
solution. How would you prepare this solution?
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Using Molarity with Stoichiometry
Since M x V = moles solute, we can use M and volume of a reactant to
find moles of reactant. With moles we can do reaction stoichiometry!
Chapter 3
Chapter 4
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Stoichiometry of Aqueous Solutions
Problem: Help! An acid spill! A student spills 500-mL of 3.4 M sulfuric
acid on the floor. Dr. Larson quickly dumps 250 g of sodium
bicarbonate on the acid to neutralize it. Was all the acid neutralized?
Solve as a limiting reactant problem!
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Acid Spill Solution Using a Reaction Table
Reaction
2 NaHCO3
+ H2SO4
—>
Na2SO4
2 H2O
2 CO2
Initial
250 g
500 mL @ 3.4
M
given amounts
0
0
0
Initial (mol)
2.976
1.70
use molar mass
or MxV
0
0
0
Moles of Reaction
Available (mol)
1.488
1.70
= mol/coeff
0
0
0
Change
–2.976
-1.488
= limiting mol
*coeff
+1.488
+2.976
+2.976
Final (mol)
0
0.212
= initial +
change
+1.488
+2.976
+2.976
Final
0
60 mL @ 3.4 M
NOT
neutralized
use molar mass
or V = mol/M
211 g
53.6 g
131 g
How much more NaHCO3 is needed to neutralize the remaining acid?
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Titrations (Covered in lab as well)
• In a titration you are trying to determine the concentration (or
amount) of a solute in a solution of unknown concentration.
• A standard solution is used to react with the solute of
interest.
• The equivalence point is when stoichiometrically equivalent
quantities are brought together (balanced eqn. needed)
• The reaction can be of various types: acid-base, gasformation, precipitation or redox.
 HCl(standard) + NaOH(unknown) —>
 HNO3(standard) + Na2CO3(unknown) —>
 Al2(SO4)3(standard) + BaCl2(unknown) —>
 MnO4–(standard) + C2O42– (unknown) —> Mn2+ + CO2
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Titration Examples
1. What is the molarity of a barium chloride solution that requires 14.5 mL of
0.412 M Al2(SO4)3 to titrate 20.0 mL of the solution to the equivalence
point?
2. Text Question 4.102: Tartaric acid, H2C4H4O6, has two acidic hydrogens.
The acid is often present in wines and precipitates from solution as the wine
ages. A solution containing an unknown concentration of the acid is titrated
with NaOH. It requires 22.62 mL of 0.2000 M NaOH solution to titrate both
acidic protons in 40.00 mL of the tartaric acid solution. Write a balanced net
ionic equation for the neutralization reaction, and calculate the molarity of
the tartaric acid solution.
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Example Problems from the Text
Text Question 4.27
You know that an unlabeled bottle contains a solution of one of the following: AgNO3,
CaCl2, or Al2(SO4)3. A friend suggests that you test a portion of the solution with Ba(NO3)2
and then with NaCl solutions. Explain how these two tests together would be sufficient to
determine which salt is present in the solution.
Text Question 4.29
Which of the following solutions has the largest concentration of solvated protons: (a) 0.1
M LiOH, (b) 0.1 M HI, (c) 0.5 M methyl alcohol (CH3OH)? Explain.
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Example Problems from the Text
Text Question 4.87
A solution of 100.0 mL of 0.200 M KOH is mixed with a solution of 200.0 mL
of 0.150 M NiSO4.
a) What precipitate forms?
b)
Write the balanced chemical equation for the reaction that occurs.
c)
d)
e)
What is the limiting reactant?
How many grams of the precipitate form?
What is the concentration of each ion that remains in solution?
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Example Problems from the Text
Text Question 4.90
A 0.5895-g sample of impure magnesium hydroxide is dissolved in 100.0 mL
of 0.2050 M HCl solution. The excess acid then needs 19.85 mL of 0.1020 M
NaOH for neutralization. Calculate the percent by mass of magnesium
hydroxide in the sample, assuming that it is the only substance reacting with
the HCl solution.
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Example Problems from the Text
Text Question 4.112
The mass percentage of chloride ion in a 25.00-mL sample of seawater was
determined by titrating the sample with silver nitrate, precipitating silver chloride.
It took 42.58 mL of 0.2997 M silver nitrate solution to reach the equivalence point
in the titration. What is the mass percentage of chloride ion in the seawater if its
density is 1.025 g/mL?
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