Download (null): 110.ReactionsIntro

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04/07/14
Resources:
Cu + HNO3 (in a jar) demo
Homework for 04/08:
Bell Work 4/7:
Acid in Eye (safety reminder) << do if time
Bell Work 4/8:
GYOH
Cannon & TOE Tie-in
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STILL going to re-arrange energy stored in chemical bonds
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Energy can be released as in cannon …
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… by way of reaction C2H2 + O2 -> ???
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When do rxns happen? Why do rxns happen? Where does energy go
when rxns happen?
A. Reactions Rule / Reaction Rules
1. Rules: Reactions always Re-arrange (= move around but never
gain or lose!)
a. Re-arrange electrons
1) Electrons cannot be created or destroyed, only rearranged
2) Old bonds are broken and new bonds are made …
3) So, new substance must be made (chem change)
4) Ex: Zn + HCl  ZnCl2 + H2
a) label each chemical with bond type (metallic,
covalent, ionic, covalent)
b) Have to break e.g. metallic zinc bond so Cl can
steal an electron and form ionic bond
c) Zn loses e, Cl gains e, H stays the same
d) electrons are re-arranged to form new compounds
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b. Re-arrange atoms
1) Atoms cannot be created or destroyed, only re-arranged
2) Same number & identity before an after do MASS must be
conserved
3) Re-write Zn + HCl equation, then create “before” and
“after” atom tables. Use to balance equation
c. Re-arrange energy
1) Energy cannot be created or destroyed only transformed
or transferred
2) Options (refer to TOE)
a) Chem PE (in bonds) transformed to KE (thermal,
electrical, radiant, mechanical)
b) KE transformed into chem PE
3) Do reduced version of Zn & HCl: one zinc pellet in
test tube plus a few ml of HCl. While bubbling discuss
where energy is stored and where it goes
4) Return to reaction, have Ss feel test tube (warm!) &
decide if reaction followed Option 1 or 2
d. “Re-arrange” collisions
1) All reactions require a certain number of collisions
between molecules and atoms – cannot change the NUMBER
of collisions for a given reaction
2) RATE of collisions can be changed by 4 factors: temp,
concentration, surface area & catalyst
2. Following the Rules: Cu + HNO3
a. Express chemical reaction using a chemical equation
1) “Reactants  Products”
2) Names: Copper (solid) + Nitric Acid (aqueous)  Copper
Nitrate + Nitrogen Dioxide (gas) + Water (liquid)
3) DO DEMO HERE:
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a) 1 penny in bottom of 500-ml round flask / beral
pipette in 1-hole stopper
b) Draw 1 ml of nitric into pipette / put pipette into
flask
c) squeeze nitric onto pennies / swirl / observe gas &
heat (CAREFULLY hold flask and move among tables,
allow a couple Ss to feel bottom of flask)
d) NOTE: rxn will only produce 70 ml of NO2 with new
penny
e) Once flask is shown to Ss, take to sink, open
stopper and quickly pour in 100-200 ml water. Dilute
HNO3 pushes rxn to NO not NO2. Flask can then be
opened and poured into sink.
4) Formulas (leave off product coefficients for now)
Cu (s) + 4 HNO3 (aq) → Cu(NO3)2 (aq) + __ NO2 (g) + __
H2O (l)
b. Rule 1 — Re-arrange electrons (break old, make new bonds)
1) What indications suggest electrons re-arranged?
2) H-NO3 bond broken, H2O bond forms etc.
3) Rule 1 √
c. Rule 2 — Re-arrange atoms (same number & identity)
1) Count atoms on both sides
2) Figure out missing coefficients on product side
3) Solution:  Cu(NO3)2 (aq) + 2 NO2 (g) + 2 H2O (l)
4) Rule 2 √
d. Rule 3 — Re-arrange energy
1) Two options — which one happened based on observation
that flask got warm?
a) Surroundings lost KE and chemicals gained PE?
b) Chemicals lost PE and surroundings gained KE?
e. Rule 4: How could we accelerate reaction?
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1) higher temp
2) higher concentration
3) use powdered or shredded copper
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A Historical Sidelight:
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Ira Remsen on Copper and Nitric Acid
Ira Remsen (1846-1927) founded the chemistry department at Johns
Hopkins University, and founded one of the first centers for
chemical research in the United States; saccharin was discovered
in his research lab in 1879. Like many chemists, he had a vivid
"learning experience," which led to a heightened interest in
laboratory work:
While reading a textbook of chemistry I came upon the statement,
"nitric acid acts upon copper." I was getting tired of reading
such absurd stuff and I was determined to see what this meant.
Copper was more or less familiar to me, for copper cents were
then in use. I had seen a bottle marked nitric acid on a table
in the doctor's office where I was then "doing time." I did not
know its peculiarities, but the spirit of adventure was upon me.
Having nitric acid and copper, I had only to learn what the
words "act upon" meant. The statement "nitric acid acts upon
copper" would be something more than mere words. All was still.
In the interest of knowledge I was even willing to sacrifice one
of the few copper cents then in my possession. I put one of them
on the table, opened the bottle marked nitric acid, poured some
of the liquid on the copper and prepared to make an observation.
But what was this wonderful thing which I beheld? The cent was
already changed and it was no small change either. A green-blue
liquid foamed and fumed over the cent and over the table. The
air in the neighborhood of the performance became colored dark
red. A great colored cloud arose. This was disagreeable and
suffocating. How should I stop this? I tried to get rid of the
objectionable mess by picking it up and throwing it out of the
window. I learned another fact. Nitric acid not only acts upon
copper, but it acts upon fingers. The pain led to another
unpremeditated experiment. I drew my fingers across my trousers
and another fact was discovered. Nitric acid acts upon trousers.
Taking everything into consideration, that was the most
impressive experiment and relatively probably the most costly
experiment I have ever performed. . . . It was a revelation to
me. It resulted in a desire on my part to learn more about that
remarkable kind of action. Plainly, the only way to learn about
it was to see its results, to experiment, to work in a
laboratory.
from F. H. Getman, "The Life of Ira Remsen"; Journal of Chemical
Education: Easton, Pennsylvania, 1940; pp 9-10; quoted in
Richard W. Ramette, "Exocharmic Reactions" in Bassam Z.
Shakhashiri, Chemical Demonstrations: A Handbook for Teachers of
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Chemistry, Volume 1.
Press, 1983, p. xiv:
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Madison: The University of Wisconsin