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Chapter 4
Types of Chemical Reactions and
Solution Stoichiometry
Water (H2O = HOH)
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Solvent - dissolves solute to make solutions
Aqueous means dissolved in water
Water dissolves many substances
Water’s molecular shape: bent
d-
d+
H2O
Water’s polarity
• Causes ionic compounds to “fall apart” in water
when opposite charges become attracted to
one another
• Called hydration
• Some molecules are more soluble in water than
others
Electrical Conductivity
• A solution’s ability to conduct electricity
• The more ions in a solution, the stronger the
electric current will be
• Pure water does NOT conduct electricity
Electrolytes
• STRONG electrolytes: solutions that conduct
electric current very efficiently
• WEAK electrolytes: solutions that conduct a
small electric current
• NONelectrolytes: solutions that permit no
current flow
Strong Electrolytes
• 100% Dissociation (breaking apart)
• Many ions are created in aqueous
solution
• Examples: soluble salts, strong acids,
strong bases
NaCl (s)
H 2O
Na+ (aq) + Cl- (aq)
Arrhenius Acid
• Substance that produces H+ ions
(protons) when dissolved in water
• Proton donor
HA(aq) + H2O(l) --> H3O+(aq) + A-(aq)
Like irrelevant
STRONG Acids
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HCl: hydrochloric acid - produces H+ and ClHNO3: nitric acid - produces H+ and NO3H2SO4: sulfuric acid - produces H+ and HSO4Not in book - HBr (hydrobromic acid), HI
(hydroiodic acid), HClO4 (perchloric acid)
• Characteristics: sour taste (vinegar, citrus fruits),
cause color changes in plant dyes, pH is low (0-7)
Arrhenius Bases
• Substance that produces a hydroxide
(OH-) ion when dissolved in water
Like irrelevant
STRONG Bases
• NaOH: sodium hydroxide - produces Na+ and OH• KOH: potassium hydroxide - produces K+ and OH• Not in book: LiOH, RbOH, CsOH, Ca(OH)2,
Sr(OH)2, Ba(OH)2 (you probably won’t see these)
• Characteristics: bitter taste, feel slippery (soaps
usually contain bases), cause color changes in
plant dyes, pH is high (7-14)
• Group I/II hydroxides
Weak Electrolyte
• Doesn’t completely dissociate in water
• Relatively few ions are produced when
dissolved in water
• Weak acids/bases
CH3COOH
CH3COO- (aq) + H+ (aq)
A reversible reaction can occur in both
directions
Nonelectrolytes
• Create almost no ions when dissolved in water
• Ex: sucrose C12H22O11
Molarity
• A unit representing concentration
• Unit is moles per liter
M = moles of solute
liters of solution
Read as “____ molar”
Molarity Example
• Grams need to be converted to moles
• mL need to be converted to L
• How many mL of solution are necessary if we
are to have 2.48 M NaOH solution that contains
31.52 g of the dissolved solid?
• Answer: 318 mL solution
• See book pg. 137-140
More on Molarity…
• Multiply molarity by number of ions in the
solution to find molarity of specific ions:
• Calculate the molarity of all the ions in the
following:
– 0.25 M Ca(OCl)2
– Anwer: 0.25 M Ca2+, 0.50 M OCl– 2 M CrCl3
– Answer: 2 M Cr3+, 6 M Cl-
Dilution
• We can add water to solutions and
consequently water them down (dilute) them.
*only the amount of water changes, not the
solvent!
M1V1 = M2V2
• M = molarity
• V = volume (units must match for 1 and 2)
Dilution Example
• What volume of 12 M hydrochloric acid must be
used to prepare 600 mL of a 0.30 M HCl
solution?
• Answer: 0.015 L = 15 mL of 12 M HCl
Types of Reactions: Review!
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Combination/Synthesis
Decomposition
Single Replacement
Double Replacement
Combustion
*NEW: Acid-Base Reactions
*NEW: Oxidation-Reduction Reactions
Combination/Synthesis
• One product!
• Two elements = combine and balance
charges
2Al + 3O2 --> 2Al2O3
• Nonmetal oxide + water produces an acid
CO2 + H2O --> H2CO3
Decomposition
• One reactant
• Breaks into its element (diatomics!)
2MgO --> 2Mg + O2
• Metal carbonate --> metal oxide + CO2
Na2CO3 --> Na2O + CO2
• Metal chlorate --> metal chloride + O2
2 Al(ClO3)3  2 AlCl3 + 9 O2
more…
Single Replacement
• Lone metal (nonmetal) + compound
• Metals (or nonmetals) switch
• CHECK ACTIVITY SERIES - lone metal must
be higher or NR
2Al(s) + 3Pb(NO3)2(aq)  3Pb(s) + 2Al(NO3)3(aq)
Double
Replacement
• Two compounds’ metals switch spots
• MUST bubble, form a molecule like water or
form a precipitate (precipitate reaction)
• USE SOLUBILITY RULES!
Solubility Rules
• Slightly soluble = Insoluble and means (s)
• Soluble means (aq)
• MUST BE MEMORIZED!!
Acronyms for Solubility
• CASHN Gia
– C = chlorates, A = acetates, S = sulfates, H =
halogens, N = nitrates, and Gia = group 1 A metals
(all soluble)…EXCEPT…
• For sulfates = Ca, Ba, Sr (remember CBS like TV)
• For halogens: Ca, Ba, Sr + Happy-whats happy? Hg Ag
Pb…mercury, silver and lead…add a -py to the end and
all the first letters spell "happy”
• If it’s not part of CASHN Gia, it’s insoluble.
• SONG?? Or this one…another…
Predict Products/Precipitates
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KCl(aq) + Pb(NO3)2(aq) ->
2KCl(aq) + Pb(NO3)2(aq) -> PbCl2(s) + 2KNO3(aq)
AgNO3(aq) + MgBr2(aq) ->
2AgNO3(aq) + MgBr2(aq) -> 2AgBr(s) +
Mg(NO3)2(aq)
Ca(OH)2(aq) + FeCl3(aq) ->
3Ca(OH)2(aq) + 2FeCl3(aq) -> 2Fe(OH)3(s) +
3CaCl2(aq)
NaOH(aq) + HCl(aq) ->
NaOH(aq) + HCl(aq) -> H2O(l) + NaCl(aq)
Combustion
Organic compound (CHO) + O2 --> CO2 + H2O
Usually releases energy (light or heat)
***NOTE!
• It is assumed from now on that if a reaction isn’t
written out you can (and will) write it out…
Net Ionic Equations
1. Balance the equation
2. Write the complete ionic equation by
breaking up aqueous compounds (ionic)
3. Cancel spectator ions (same on both sides
including state)
4. Rewrite what’s left (reduce coefficients if
necessary)
See example in book pg. 151
Complex Stoichiometry Problems
• We can add limiting
reagent and molarity
to make more
complex
stoichiometry
problems
• Example pg. 152153
Acid-Base Reactions
• Reminder: Arrhenius (irrelevant) acids/bases
• Brønsted-Lowry acids/bases (protons only):
Acids are proton donors, bases are proton
acceptors
• Acid-Base Reactions are called neutralization
reactions because the acid/base becomes
neutralized by the base/acid added to it.
• *OH- will react with both strong and weak acids
acid + base
salt + water
Example
• What volume of a 0.100 M HCl solution is
needed to neutralize 25.0 mL of 0.350 M
NaOH?
Write out reaction: HCl + NaOH --> NaCl + H2O
0.250L X 0.350mol NaOH X 1mol HCl X __1L HCl_
1L
1mol NaOH 0.100 mol HCl
Solve for = 0.0875L of hydrochloric acid needed
Titrations!
• If we know the
concentration and
amount of either the
acid or base (titrant in
the buret), we can
determine the
concentration of the
other (base or acid) by
adding the titrant until
the reaction is
neutralized
(equivalence point)
Titration Setup
• An indicator is added to the solution to be analyzed
(usually phenolpthalein is added to the acid). It will
indicate when the solution is one drop beyond a
specific pH (phenolpthalein is 7). This is called the
endpoint.
• Endpoint is where the color changes vs. equivalence
point, where the number of moles of acid = moles of
base
• Video
Titration Steps
1. Record initial volume of the known
concentration of base/acid in buret
2. Record amount (usually mL) of unknown
concentration acid/base in beaker
3. Add a few drops of indicator to beaker
4. Add titrant until the indicator has just
permanently changed color
5. Record the final volume of the known
concentration of base/acid in buret
6. You will now have the information needed to
calculate the concentration of the unknown
Example
• Use the following data to calculate the
concentration of the HCl solution:
• HCl in beaker with indicator: 10.4mL
• Initial buret reading: 3.5mL 0.100M NaOH
• Final buret reading (pink): 15.4mL 0.100M
NaOH
**write out the reaction first!!
Answer: 0.114M HCl
Book Example
• More complicated problem on page 159
• TIP pg. 160: The first step in the
analysis of a complex solution is to write
down the components and focus on the
chemistry of each one. When a strong
electrolyte is present, write it as
separated ions.
4.9
Oxidation-Reduction Reactions
Oxidation-Reduction Rxns.
(Redox Rxns)
• A chemical reaction where one or more
electrons are transferred
• Photosynthesis, combustion reactions,
sugar/fat/protein oxidation in humans
Covalent Vs. Ionic Bonds
• Covalent - electrons are shared
• Between elements in molecular
compounds (remember two nonmetals
make binary molecular compounds???)
• Ionic - opposite charges (cations and
anions) are attracted like magnets
(strong bond)
Oxidation States (oxidation
numbers)
• Helps show where electrons are transferred in
chemical reactions…RULES…
• Lone and diatomic elements = zero
• Oxygen = -2 except in H2O2 (charge is -1)
• In ionic compounds, ions’ #s are same as their charge
• Do negative first. Charges add to make zero (or a
charge if the compound is charged)
Ion charges: n+ or n- (#before charge)
Oxidation states: +n or -n (# after charge)
Try these…
• NO2
• Cl2
– O: -2 each, N: +2
– Cl: 0
• PbS
• Fe3O4
– S: -2, Pb: +2
– O: -2 each, Fe: +8/3
• SnCl6-2
• This can happen
– Cl: -1 each, Sn: +4
(non whole number
• H+
oxidation numbers)
– H: +1
Purpose
• In reactions, we want to know what elements
gain or lose electrons
• This is shown through assigning oxidation
numbers
• The element that becomes more negative gains
electrons, decreases oxidation state (reduction)
• The element that becomes more positive loses
electrons, increases oxidation state (oxidation)
• One doesn’t happen without the other.
Memorize!
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LEO (lose electrons oxidation) the lion goes
GER (gain electrons reduction)…or…
• OIL (oxidation is lost) RIG (reduction is
gained)
Example: CH4(g) + 2O2(g) --> CO2(g) + 2H2O(g)
1. Complete/balance equation if necessary
2. Assign oxidation states/numbers
3. Determine which element is oxidized/reduced
Also…
• C: -4 -> +4 so it’s oxidized
– Donates e- so REDUCING
AGENT (allows something
else to be oxidized)
• O: 0 -> -2 so it’s reduced
– Accepts e- so OXIDIZING
AGENT (allows something
else to be reduced)
You try…
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PbS(s) + O2(g) --> PbO(s) + SO2(g)
Assign oxidation #s
S: -2 -> +4, O: 0 ->-2
S: oxidized, reducing agent
O: reduced, oxidizing agent
Copper wire reacts with silver nitrate to form silver metal.
What is the oxidizing agent in the reaction?
Balancing Redox Rxns.
• Fancy way of balancing skeleton equations
using electrons
• Two ways:
– By oxidation states (LEARNING.)
– Half-reactions (chapter 18)
Steps to follow
1. Assign oxidation states/numbers
2. Use lines to connect oxidation and reduction
on both sides of the reaction
3. Add coefficients that will neutralize the number
of electrons gained and lost
4. Balance the rest of the equation
5. Make sure your states of matter are
appropriate
Example
Balance the reaction between solid
lead(II)oxide and ammonia gas to produce
nitrogen gas, liquid water, and solid lead.