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Types of Reactions
Ch. 4 Part II
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Types of Chemical Reactions (3)
Precipitation
Acid-Base
Reactions
Reactions
Oxidation-Reduction
Reactions
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Precipitation Reactions

A precipitation reaction can also be called a double
displacement reaction.

When ionic compounds dissolve in water, the resulting
solution contains the separated ions.

If a solid forms from a combination of selected ions in
solution, the solid must contain an anion part and cation part,
and the net charge on the solid must be zero.

There are some simple solubility rules to help predict the
products of reactions in aqueous solutions.
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Simple Rules for Solubility

1. Most nitrates (NO3-) salts are soluble.

2. Most salts containing the alkali metal ions (Li+, Na+, K+, Cs+,
Rb+) and the ammonium ion (NH4+) are soluble.

3. Most chloride, bromide, and iodide salts are soluble. Notable
exceptions are salts containing the ions Ag+, Pb+, and Hg2+

4.Most sulfate salts are soluble. Notable exceptions are BaSO4,
PbSO4, HgSO4, and CaSO4

5.Most hydroxide salts are only slightly soluble. The important
soluble hydroxides are NaOH and KOH. The compounds
Ba(OH)2, Sr(OH)2, and Ca(OH)2 are slightly soluble.

6.Most sulfide (S2-), carbonate (CO32-), chromate (CrO42-), and
phosphate (PO43-) salts are only slightly soluble.

Page 144
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Solubility Memorization
Requirements for AP Exam
 All
sodium, potassium, ammonium, and nitrate
salts are soluble in water.
**Memorization of “solubility rules” is beyond the scope of this
course and the AP Exam.**
Rationale: Memorization of solubility rules does not deepen
understanding of the big ideas.
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Predicting Reaction Products

To begin, focus on the ions in solution before any reaction
occurs.

Complete and balance the following reactions, determining in
each case, if a precipitate is formed.

KCl(aq) + Pb(NO3)2(aq)

AgNO3(aq) + MgBr2(aq)

Ca(OH)2(aq) + FeCl3(aq)

NaOH(aq) + HCl(aq)
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Describing Reactions in Solution

Formula Equation (overall reaction – does not identify ions)
2AgNO3(aq) + Na2S(aq)

Complete Ionic Equation (includes ions in solution)
2Ag+(aq) + 2NO3-(aq) + 2Na+(aq) + S2-(aq)

Ag2S(s) + 2NaNO3(aq)
Ag2S(s) + 2NO3-(aq) + 2Na+(aq)
Net Ionic Equation (identifies species that undergo chemical change)
2Ag+(aq) + S2-(aq)
Ag2S(s)
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Writing Equations for Reactions
Example

Write the formula, complete ionic, and net ionic forms for
each of the following equations.

Aqueous nickel (II) chloride reacts with aqueous sodium
hydroxide to give a nickel (II) hydroxide precipitate and aqueous
sodium chloride.

Solid potassium metal reacts with water to give aqueous
potassium hydroxide and hydrogen gas.

Aqueous sodium hydroxide reacts with aqueous phosphoric acid
to give water and aqueous sodium phosphate.
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Practice Problems

Page 173 # 37, 39, 41, and 43

15 Minutes
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
Stoichiometry of Precipitation
Reactions
Steps:

1. Identify the species present in the combined solution, and determine
what reaction occurs.

2. Write the balanced net ionic equation for the reaction.

3. Calculate the moles of reactants.

4. Determine which reactant is limiting.

5. Calculate the moles of product or products, as required.

6. Convert to grams of other units, as required.
**No two examples are alike, the PROCESS to solving will be similar**
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Solution Stoichiometry Example

What mass of Fe(OH)3 is produced when 35. mL of a 0.250 M
Fe(NO3)3 solution is mixed with 55 mL of a 0.180M KOH
solution?

Individual Practice Problem Page 173 #49 and 50
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Check for Understanding
Precipitation Stoichiometry
Please complete the following on a sheet of paper to turn in.
 When
aqueous solutions of Na2SO4 and Pb(NO3)2
are mixed, PbSO4 precipitates. Calculate the
mass of PbSO4 formed when 1.25 L of 0.0500 M
Pb(NO3)2 and 2.00 L of 0.0250 M Na2SO4 are
mixed.
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Acid-Base Reactions


The Bronsted-Lowry Definitions

Acid is a proton donor.

Base is a proton acceptor.
An acid-base reaction is often called a neutralization
reaction.
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
Acid-Base Example
How many mL of a 0.800M NaOH solution is needed to just neutralize
40.00mL of a 0.600M HCl solution? (Hint: Write down the reaction and convert to
moles, and relate moles of acid to moles of base.)

Individual Example Problems 65 page 171 – 8 Minutes
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Acid-Base Titrations

Volumetric analysis


Titration


Involves delivery (from a buret) of a measured volume of a solution of
known concentration (the titrant) into a solution containing the
substance being analyzed (analyte).
Equivalence point (Stoichiometric point)


A technique for determining the amount of a certain substance by
doing a titration.
Point in titration where enough titrant has been added to react exactly
with the analyte
Indicator


A substance added at the beginning of the titration that changes color
at (or very near) the equivalence point
Point at which indicator actually changes color is the endpoint of the
titration.
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Acid-base titration (Standardizing)
 You
wish to determine the molarity of a solution of
sodium hydroxide. To do this, you titrate a 25.00mL
aliquot of your sample, which has had 3 drops of
phenolphthalein indicator added so that it is pink,
with 0.1067 M HCl. The sample turns clear after the
addition of 42.95 mL of the HCl. Calculate the
molarity of your NaOH solution.

Individual Example Problems 63 and 65 page 174 – 10 Minutes
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Oxidation Reduction Reactions

Reactions in which one or more electrons are transferred
(redox reactions)

Example:
H2(g) + Cl2(g)
2HCl(g)
Electrons are transferred from the hydrogen to the chlorine.

Table 4.2 Rules for Assigning Oxidation States (pg 156)

Remember: The sum of the oxidation states in a neutral
compound must equal zero.
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Oxidation States

Assign oxidation states to each of the atoms in the following
compounds:
Hint: Check that oxidation states sum up to the charge on the compound/ion.

CaF2

C2H6

H2O

ICl5

KMnO4

SO42-
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Practice- Oxidation States
Fe2O3 + 2Al
Al2O3 + 2Fe
THE IRON GAINED ELECTRONS. IT HAS BEEN REDUCED.
THE ALUMINUM LOST ELECTRONS. IT HAS BEEN OXIDIZED.
OIL – Oxidation Involves Loss (of electrons)
RIG- Reduction Involves Gain (of electrons)
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Tricky…Tricky…

Oxidizing agent


Something that is reduced (it causes something else to oxidize)
Reducing agent

Something that is oxidized (it causes something else to reduce)
Fe2O3 + 2Al

Iron was reduced.

Aluminum was oxidized.

Iron (II) oxide was the oxidizing agent.

Aluminum was the reducing agent.
Al2O3 + 2Fe
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OIL RIG Practice

For each reaction, identify the atoms that undergo reduction
or oxidation (and charges). Also, list the oxidizing and
reducing agents.
2H2(g) + O2(g)
2H2O(g)
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
2AgCl(s) + H2(g)
2H+(aq) + 2Ag(s) + 2Cl-(aq)
+ Balancing Oxidation-Reduction
Equations (Half-Reaction Method)

Steps: Acidic

1. Write separate equations for the oxidation and reduction halfreactions.

2. For each half-reaction

Balance all the elements except hydrogen and oxygen.

Balance oxygen using H2O.

Balance hydrogen using H+.

Balance the charge using electrons (e-).

3. If necessary, multiply one or both balanced half-reactions by an
integer equalize the number of electrons transferred in the two
half-reactions.

4. Add the half-reactions, and cancel identical species.

5. Check that the elements and charges are balanced.
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Balancing Oxidation-Reduction
Equations (Half-Reaction Method)

Steps: Basic

The same as acidic and then…

6. Add a number of OH- ions to both sides so that you just balance
excess H+ ions.

7. H+ and OH- will form H2O on the side with excess H+. Free OHwill appear on one side of the equation.

8. Double check atoms and charges!
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Balancing Redox Example - Acid
Cu(s) + HNO3(aq)

Copper is being oxidized:

Nitrogen is being reduced:
Cu2+(aq) + NO(g)
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Balancing Redox Reactions - Base
Cr2O72-(aq) + 2NO(g) + 6H+(aq)
2Cr3+(aq) + 2NO3-(aq) + 3H2O(l)
**Already balanced in acid**
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Practice Time!
Page
Be
174, # 67, 69, 71, 73, and 75
sure to show your work!
+ Chemistry of Copper Lab - _______

Pre-lab: REQUIRED FOR ENTRY INTO LAB!!



Read lab handout thoroughly.
Set-up data organization
 What data are you collecting for each step?
Complete procedure component of lab
 Safety:



Safety goggles at ALL times!
Concentrated acids will be used!!
 Can cause severe burns
 No touching your face!
 Wash your hands after completing lab.
 Read thoroughly…carefully pour acids and remove hot beakers
from hot plate with tongs! (A hot beaker looks like a cold one!)
Avoid congestion around the fume hood (Step 5 should be done in
fume hood!).
Keep lab stations cleaned through multistep procedure.
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Copper Lab Partners

Group 1:

Group 2:

Group 3:
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Step 1: Copper Lab Task
 Label
beaker with tape
 Block
 Student group initials
+ Bleach Redox Lab
 PRE-LAB
 Read
DUE ___________!!
handout thoroughly! (in black binder)
 Complete
the pre-lab questions
 Hint: Dilution, Concentration, Stoichiometry
 Required
 Also
for entry into lab.
for _________…
 Begin working on Lab Report for Cu Lab
 Continue practice with Precipitation, Acid-Base, and
redox reactions.
 Begin reading Ch.5 Gases
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Today’s Agenda

Review Expectations for Tomorrow

Check Pre-Lab for Bleach Redox Lab.

Review Bleach Redox Lab

Take final copper mass reading for Lab C.

Complete/Collect Data for the Bleach Redox Lab
 Watch Time.
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Due for Tomorrow…

Copper Lab Report (Heavily Qualitative)

Date when data was collected

Read Ch. 5 Gases

Begin completing Bleach Redox Lab Report

Due _________ at the BEGINNING of class.
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Bleach Redox Lab

Pre-lab

Safety:

Bleach is a 5% solution of sodium hypochlorite. This solution is a
corrosive liquid; it can cause skin burns. The solution reacts with
acid and with heat to evolve chlorine gas. The solution is
moderately toxic by ingestion and inhalation. Keep away from
skin and clothing. Hydrochloric acid solution is toxic by ingestion
or inhalation. It is corrosive to skin and eyes.

CAUTION: Adding hydrochloric acid to bleach may cause
chlorine gas to be given off. Step 6 is completed in the fume
hood!

Wear your goggles at ALL times! Do NOT touch your face. Wash
your hands with soap before leaving the lab.
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Data Collection

Collect mass of copper. This may be placed in the trash and
the cleaned beaker back on the drying rack.

Complete the 3 titrations of bleach. READ INSTRUCTIONS
CAREFULLY!

ALL glassware must be away from the edge of the tables.

Practice caution in the lab! Be aware of your surroundings.

When finished clean lab station and return to seat and begin
working on the reports or reading Ch. 5.
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Up Next…. GASES!! Ch. 5