October 10th, 2012 - The Bronx High School of Science

Report
1)
2)
Do Now:
Compare Homework with a partner, how
did you do?
With your partner, write a list of steps on
how to solve a limiting reagent problem
to add to your notebook.
•
Define the word solution.
–
•
How do you differentiate between the
solute and the solvent in a solution?
–
•
The solute is dissolved into the solvent
Which is present in the smallest amount,
the solute or the solvent?
–
•
A homogenous mixture of two or more
substances.
Solute
How do you determine if a solution is
considered aqueous?
–
Solutions in which water is the solvent.
A high percentage of naturally occurring
chemical reactions occur with water as the
solvent; why is water such an excellent
solvent?
Why is water considered a polar molecule?
How does this benefit water a solvent?
•
How do we know, prior to testing for electricity,
whether or not a substance will contain
electrolytes?
–
•
If a substance forms ions in solution. IE: NaCl.
How does a nonelectrolyte differ from an
electrolyte?
–
–
May dissolve in water, but does not dissociate into ions.
Electrolytes conduct electricity, non electrolytes do
not.
Ionic Compounds in Water
•
Summarize your observations of what
occurred on the molecular level when solid
NaCl was added to water.
–
–
IONIC COMPOUND DISSOLVES IN WATER, IONS
DISSOCIATE.
EACH ION IS SURROUNDED BY SEVERAL WATER
MOLECULES (AQUEOUS IONS “(aq)”
»
•
•
THESE IONS ARE CONSIDERED “SOLVATED”
Illustrate this concept in your notebook.
How is the electric current created?
When a molecular compound (ie: CH3OH) dissolves in
water, the solution usually consists of intact
molecules dispersed homogeneously throughout the
solution.
There is nothing in solution to transport electric
charge and therefor most molecular compounds are
non-electrolytes
**Important exceptions: ie: HCl, NH3
•
•
Compounds whose aqueous
solutions conduct electricity
well are called “Strong”
Electrolytes. (exist in solution
mostly as ions)
Compounds whose aqueous
solution conduct electricity
poorly are called weak
electrolytes
NaCl(aq)  Na+(aq) + Cl-(aq)
•
•
Why is a single arrow used to show this
dissociation?
Soluble ionic compounds, strong acids, and
soluble strong bases are considered strong
electrolytes.
 Molecular
compounds that produce small
concentration of ions when dissolved = weak
electrolytes.
 Ie:
acetic acid, HC2H3O2, is primarily present in solution as
molecules. Approx. 1 percent is present as ions.
 **Note: DO NOT confuse the extent to which an electrolyte dissolves with whether it is strong
or weak.For
example, HC2H3O2 is extremely soluble in water but
is a weak electrolyte. In contrast, Ba(OH)2 is not very soluble,
but the amount of the substance that does dissolve dissociates
almost completely. Therefore, Ba(OH)2 is a strong electrolyte.
 How
does this dissociation differ from the previous
dissociation of NaCl?
 Why
is the double arrow important in this equation?
 Do
Now:
 Summarize
yesterday’s lesson.
•
•
How can we classify this type of reaction?
Define the word precipitate.
–
Insoluble solid formed by a reaction in solution
Ie: Pb(NO3)2 + 2KI → ??
–

–
Define the word solubility.
•
The amount of a particular substance that can be
dissolved in a given quantity of solvent at that
temperature.
When does a substance become regarded as
insoluble?
•
Substance with a solubility of less than 0.01 mol/L
•

N: NITRATES
A: ACETATES
G: GROUP 1 (ALKALI METALS)
S: SULFATES (except PMS and CaStroBear)
A: AMMONIUM
G: GROUP 17 (except PMS)
C: CARBONATES ( except for G1, Ammonium)
A: ALCOHOLS (except for G1, CaStroBear, Ammonium)
P: PHOSPHATES (except for G1, Ammonium)
S: SULFIDE (except for G1, CaStroBear, Ammonium)

•
Are the following compounds soluble or insoluble
and why?
–
–
–
–
–
–
–
KNO3
Li3PO4
AgCl
NH4OH
Ba3PO4
Hg2S
Na2CO3
•
Predict a general formula for an “exchange”
or metathesis reaction.
•
•
How would you describe an “exchange”
reaction?
•
•
•
AX + BY → AY + BY
Swapping of ions in solution
Precipitation and acid-base reactions exhibit
this pattern.

Consider the following:
2KI(aq) + Pb(NO3)2(aq) → PbI2(s) 2KNO3(aq)
–
–
•
•
Both Reactants are colorless solutions. When mixed,
they form a bright yellow precipitate of PbI2 and a
solution of KNO3.
Final Product contains solid PbI2, Aqueous K+ ions
and Aqueous NO3- ions.
HOW MIGHT WE DIFFERENTIATE BETWEEN THE
MOLECULAR, COMPLETE IONIC, and NET IONIC
EQUATIONS?

•
Molecular: lists all species in complete chemical
forms
•
•
2KI(aq) + Pb(NO3)2(aq) → PbI2(s) 2KNO3(aq)
Complete Ionic: lists all strong electrolytes in rxn
as ions
•
•
•
•
Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq) → PbI2(s) + 2K+(aq) +2NO3(aq)
Only strong electrolytes dissolved in solution are written
in ionic form.
Weak electrolytes and non-electrolytes are written in
complete chemical form

•
Lists only ions which are not common on both sides
of the equation
Pb2+(aq) + 2NO3-(aq) + 2K+(aq) + 2I-(aq) → PbI2(s) + 2K+(aq) +2NO3-(aq)
•
Which ions can we remove from each side?
–
•
2NO3- and 2K+
Predict what the net ionic equation will look like?
•
Pb2+(aq) + 2I-(aq) → PbI2(s)
•

•
•
•
How do we define spectator?
Why are the ions we removed considered
spectator ions?
Formulate a step-by-step “how to” list for
creating the net-ionic equation.
1) Write a balanced molecular equation
2) Rewrite equation to show ions that were present after
dissociation (only strong electrolytes)
3) Identify and cancel “spectator” ions

Complete the worksheet with a partner, be
prepared to instruct the class on how you
reached your answers.
 Define

the word Acid.
Substances that are able to ionize in aqueous
solutions to form H+
 Using
this information, how might we define
the word base?

Proton acceptors
 How
might we differentiate between a
monoprotic acid and a diprotic acid?



Monoprotic ionize to form 1 H+ ion.
Diprotic ionizes to form 2 H+ ions.

 How
soluble are strong acids and strong
bases?

Very. They completely ionize in solution.
 Strong

Group 1A metal hydroxides, Ca(OH)2, Ba(OH)2 and
Sr(OH)2
 Strong

Bases:
Acids:
HCl, HBr, HI, HClO3, H2SO4, and HNO3
 Write
the ionization of a strong acid.
 How
does the ionization of a weak acid/base
compare to that of strong acids/bases.

Partially ionized in solution.
 HF(aq)
is a weak acid; most acids are weak
acids.
 Write the ionization of a weak acid.

October 22nd, 2012
 Do
Now:
 How
do you determine a weak acid/base.
 Write
the Net Ionic Equation for the
following reaction:
 When aqueous solutions of sodium phosphate and
calcium chloride are mixed together, an insoluble
white solid forms.
Neutralization Reactions and Salts
 Write
a generalized equation for a
neutralization reaction.

Acid + Base  Water + Salt
 Define

the word Salt.
Any ionic compound whose cation (+) comes from
a base and anion (-) comes from an acid.
 Example:

Mg(OH)2(s) is a suspension and HCl is added.
Write the net ionic equation.
 MgOH2(s)
 
+ 2H+ (aq)  Mg2+ (aq) + 2H2O (l)
Neutralization Reactions with Gas
Formations

How do sulfides act as bases?
 There are many bases besides OH- that react with H+
to form molecular compounds.
 The reaction of sulfides with acids give rise to H2S in
gaseous form.

Write the Net Ionic equation of Sodium Sulfide reacting
Hydrochloric Acid.


2H+ (aq) + S2- (aq)  H2S (g)
Carbonates and hydrogen carbonates (or
bicarbonates) will form CO2 (g) when treated with an
acid.
 Write the net ionic equation of sodium bicarbonate
reacting with Hydrochloric acid.


H+ (aq) + HCO3- (aq)  H2O (l) + CO2 (g)

BOMBS AWAY!
Oxidation-Reduction Reactions
OXIDATION
REDUCTION

Oxidation-Reduction
 How
is it determined which substance is
undergoing oxidation vs. reduction?


Loss of electrons = Oxidation
Gain of electrons = Reduction
 Why

are oxidation numbers useful?
Oxidation numbers help us keep track of
electrons during chemical reactions.
 How
do we determine the oxidation numbers
of atoms?

RULES!
I lost an electron!
I’M POSITIVE!
Hey! Why so
Serious? What are
you looking for?
Are you sure??
RULES FOR OXIDATION STATES
 For
an atom in its elemental state, the
oxidation number is:
 For any monatomic ion, the oxidation
number equals:
 The oxidation number of oxygen is usually:

The major exception is in peroxides (containing
the O22- ion)
 The
oxidation state of Hydrogen is:
 The oxidation number of Flourine is:
 The sum of oxidation numbers in polyatomic
a polyatomic ion:

Practice determining oxidation
state:
 Complete
the worksheet handed out!
Identifying Oxidation Vs.
Reduction
 PRACTICE
SHEET!
 Do
Now:
Find the oxidation states of each of the elements
in the following compounds:
1. P2O5
2. NaH
3. C2O724. SnBr4
5. BaO2

How can we generalize a reaction between a
metal or an acid and a salt?

Write a general equation for this type of
reaction. Define this type of reaction.

Given this example: Zinc reacting with Hydrochloric acid,
predict the products and determine which is being reduced
and which is being oxidated.
 Write
the molecular and net ionic equation
for the following reaction:

Zinc reacting with hydrobromic acid.
 Why
is writing the net ionic equation helpful
when discussing oxidation and reduction
reactions?

Complete 1 through 3 on
a the worksheet given
to you
It is possible for a metal to be oxidized in the
presence of a salt.
Fe(s) + Ni(NO3)2(aq)
 Predict
the products and write a molecular
and net ionic equation for the reaction
above.
 Will
a metal always be oxidized in the
presence of a salt? Explain your answer.
The activity series lists metals in order of
decreasing oxidation.
 How
can we differentiate between active
metals and noble metals?
 How
do we determine if a reaction will occur
using the activity series?
Answer the aim in
a summary
paragraph.
 Do
1.
2.
Now:
It is said that HClO4 is a strong acid, whereas
HClO2 is a weak acid. What does this mean in
terms of the extent to which they ionize in
solution?
Classify each of the following as a nonelectrolyte, weak electrolyte, or strong
electrolyte in water:
1.
2.
3.
4.
H2SO3
C2H5OH (ethanol)
NH3
KClO3
 How
can we differentiate between strong
and concentrated solutions?
 Define
 How
concentration
do we express concentration of a
solution?
Molarity (symbol M) expresses concentration of
solution (number of moles solute in a liter of
solution)
 Calculate
the molarity of a solution made by
dissolving 23.4 grams of sodium sulfate in
enough water to form 125 mL of solution.
 Calculate
the molarity of a solution made by
dissolving 5.00 grams of glucose in sufficient
water to form exactly 100 mL of solution.
 Calculate
the number of grams of solute in
0.250 L of 0.150 M KBr.
When an ionic compound dissolves, the
relative concentration of ions depend on the
chemical formula of a compound.
 In
a 1.0 M solution of NaCl, what are the
concentration of Na+ ions and Cl- ions?
 In
a 1.0 M solution of Na2SO4, what are the
concentrations of the Na+ ions and the SO42ions?
 What
is the molar concentration of K+ ions in
a 0.015 M solution of potassium carbonate?
 Which
will have the greatest concentration
of potassium ion: 0.20 M KCl, 0.15 M K2CrO4,
or 0.080 M K3PO4?
If we know any two quantities involved in the
molarity equation, we can solve for the
third.
 Calculate
the number of moles of HNO3 in
2.O L of a 0.200 M HNO3 solution.
How do we dilute solutions?
Moles solute in conc soln = moles dolute in dil solution
 We
want to prepare 250.0 mL (convert to L)
of 0.100 M CuSO4 solution by diluting a stock
solution containing 1.00 M CuSO4. How many
mL of concentration solution do we need?
 Do
Now:
 How
many milliliters of 3.0 M H2SO4 are
needed to make 450 mL of 0.10 M H2SO4?
 What
volume of 2.50 M lead (II) nitrate
solution contains 0.0500 mol of Pb2+ ions?
Two types of units exist:

Laboratory Units:

Chemical Units:
How do you think these differ?
Always convert laboratory units into chemical
units first.


Grams  moles using molar mass
Volume/molarity  moles using M= mol/L
Use Stoich. Coefficients to move between products and reactants.
Convert lab units back into required units

 Why
do we use
titrations?
 How
are titrations
helpful in
determining
concentration?
Suppose we know the molarity of an NaOH
solution and we want to find the molarity of
a given HCl solution.






What do we already know?
What do we want to know?
How do we get there?
How do we know when they are neutralized?
Define Equivalence Point.
After the titration is complete, how can we now
solve for the molarity?
 What
mass of NaCl is needed to precipitate
the silver ions from a 20.0 mL of 0.100 M
AgNO3 solution?
 How
many mL of 0.120 M HCl are needed to
completely neutralize 50.0 mL of 0.101 M
Ba(OH)2 solution?
 If
42.7 mL of 0.208 M HCl solution is needed
to neutralize a solution of Ca(OH)2, how
many grams of Ca(OH)2 must be in the
solution?

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