Download 225 Unit 7, Lab 1 - Pope John Paul II High School

Unit 7, Lab 1
We continue to build on our model of matter as bonded atoms that combine in definite ratios to
include the rearrangement of these atoms to form new substances during chemical reactions.
Title:
Purpose/Question:
Procedure:
Day 1
1. Label, then record the mass of a Dixie cup.
2. Record the mass of three iron nails together and then place them into the Dixie cup.
3. Add about 50 mL of copper (II) sulfate solution to the dixie cup.
4. Observe the reaction; record your observations. Place the labeled Dixie cup in the place
designated by your teacher.
Day 2
6. Label a second, clean Dixie cup. Record its mass.
7. Remove the nails from the first cup with forceps. Rinse or scrape the precipitate (copper
metal) from the nails into your first labeled Dixie cup.
8. Once all the precipitate is off the nails, place the nails in the second clean, labeled Dixie
cup. Note the appearance of the nails.
9. Decant solution from the first Dixie cup. Rinse the precipitate with about 25 mL of
distilled water. Try to lose as little of the solid copper as you can when you decant. After a
2nd rinse with distilled water, rinse the copper with 25 mL of 1 M HCl. Rinse one last time
with distilled water. Then place the labeled beaker in the drying oven.
Day 3
11. Record the mass of the second Dixie cup and dry nails.
12. Mass the beaker + dry copper.
13. Clean up your lab table. Discard the cup with the nails. Place the cup with copper in a
place indicated by the instructor.
Data:
Organize your measurements in a data table.
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Analysis:
1. Calculate the mass of copper produced in the reaction.
2.
Calculate the mass of iron used during the reaction.
3.
Calculate the moles of copper involved in the reaction.
4.
Calculate the moles of iron involved in the reaction.
5.
Determine the ratio moles of copper.
moles of iron
Express this ratio as an integer. For example, a ratio of 1.33 can be expressed as
0.67 can be expressed as
4
;
3
2
, etc.
3
Conclusion:
1.
Why did the reaction stop? Which reactant was completely consumed or used up during
the reaction? Provide observations/evidence that support your answer.
2.
There are two possible chemical equations that could describe the reaction observed
because iron is a transitional metal. Write the balanced chemical equations below.
a.
b.
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3.
Based on the class data, which option is the reaction observed in lab? Explain.
4.
What is this type of reaction called?
5.
What would happen to the ratio of copper to iron if you had placed six nails in the
beaker instead of three? Why?
6.
What would happen to the ratio of iron to copper if you allowed the reaction go for less
time? (The nails were removed from the copper (II) chloride solution after only 15
minutes.)
7.
What is the accepted ratio of iron to copper in this reaction? ___________________
Suggest specific reasons why your experimental ratio measured in the Nail Lab differs
from the accepted value? (Your reasons must be logical given your ratio.)
8.
In the space given below, draw a diagram that illustrates what happens during the
Nail Lab. Be sure that the drawing obeys the law of conservation of mass. The
number and kind of atoms on the reactant side must equal the number and kind of
atoms on the product side.
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Name:______________________________________
Veritas:______________________________________
Unit 7, Worksheet 1—
Rearranging Atoms
Background
Describe what you already know about each of these ideas. Give an example in each of the
last 4 items.
Features of Our Current Model of Matter (include a diagram):
Conservation of Mass:
Chemical Formula:
Subscripts in formulas:
Coefficient (Hint: what is the function of a coefficient in math?):
Procedure:
1. Use your atom model kit to construct the reactant molecules for each chemical change
below. Then rearrange the atoms to form the product molecules. Add more reactant
molecules as needed to form complete product molecules with no left-overs.
2. Draw particle diagrams for each reactant molecule used and each product molecule
produced under the reaction.
3. Determine the number of each reactant molecule you needed in order to make the
product(s) with no leftovers (a complete reaction) and record each number as a coefficient in
front of its reactant formula.
4. Determine how many product molecules you would get from the complete reaction. Write
that number as a coefficient in front of each product formula.
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Rearranging Atoms
1.
_____H2
+
_____O2
→
_____ H2O
+
_____Cl2
→
_____ HCl
Diagram:
2.
_____H2
Diagram:
3. _____Na
+
_____O2
→
_____ Na2O
_____H2
→
_____ NH3
Diagram:
4. _____N2
Diagram:
5.
+
_____CH4
+
_____O2
→
_____CO2
+
_____O2
→
_____ NO2
+
_____H2O
Diagram:
6.
_____NO
Diagram:
7. _____Fe
+
_____Cl2
→
_____ FeCl3
Diagram:
8. ____CH3OH + _____O2 →
Diagram:
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_____CO2
+ _____H2O
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Analysis
1. In each the equation for each reaction, compare the total number of atoms you have before
the reaction (reactant atoms) to the total number after the reaction (product atoms).
2. At the beginning of the year we observed that mass is conserved in changes. How does your
answer to question 1 explain conservation of mass?
3. Look at the product molecule (ammonia) in reaction #4.
a. What does the coefficient tell us about this substance?
b. What do the subscripts on the nitrogen and hydrogen in NH3 tell us about the
composition of the ammonia molecule?
c. Note that the sum of the reactant coefficients does not equal the sum of the product
coefficients for reaction #4. Yet in reaction #2, the sums are equal. Explain why the
sums of coefficients do not necessarily have to equal one another in a reaction.
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Chemical Equations
by Anthony Carpi, Ph.D.
Chemical reactions happen all around us: when we light a match, start a car, eat
dinner, or walk the dog. A chemical reaction is the process by which substances bond
together (or break bonds) and, in doing so, either release or consume energy. A
chemical equation is the shorthand that scientists use to describe a chemical reaction.
Let's take the reaction of hydrogen with oxygen to form water as an example. If we
had a container of hydrogen gas and burned this in the presence of oxygen, the two
gases would react together, releasing energy, to form water. To write the chemical
equation for this reaction, we would place the substances reacting (the reactants) on
the left side of an equation with an arrow pointing to the substances being formed on
the right side of the equation (the products). Given this information, one might
guess that the equation for this reaction is written:
H+O
H2 O
The plus sign on the left side of the equation means that hydrogen (H) and oxygen (O)
are reacting. Unfortunately, there are two problems with this chemical equation.
First, both hydrogen and oxygen are found as diatomic molecules, H2 and O2,
respectively. Hydrogen gas, therefore, consists of H2 molecules; oxygen gas consists of
O2. Correcting our equation we get:
H2 + O2
H2 O
But we still have one problem. As written, this equation tells us that one
hydrogen molecule (with two H atoms) reacts with one oxygen molecule (two O atoms)
to form one water molecule (with two H atoms and one O atom). In other words, we
seem to have lost one O atom along the way! To write a chemical equation correctly,
the number of atoms on the left side of a chemical equation has to be
precisely balanced with the atoms on the right side of the equation. How does this
happen? In actuality, the O atom that we "lost" reacts with a second molecule of
hydrogen to form a second molecule of water. During the reaction, the H-H and O-O
bonds break and H-O bonds form in the water molecules.
The balanced equation is therefore written:
2H2 + O2
2H2O
In writing chemical equations, the number in front of the molecule's symbol (called a
coefficient) indicates the number of molecules participating in the reaction. If no
coefficient appears in front of a molecule, we interpret this as meaning one.
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In order to write a correct chemical equation, we must balance all of the atoms on
the left side of the reaction with the atoms on the right side. Let's look at another
example. If you use a gas stove to cook your dinner, chances are that your stove burns
natural gas, which is primarily methane. Methane (CH4) is a molecule that contains
four hydrogen atoms bonded to one carbon atom. When you light the stove, you are
supplying the activation energy to start the reaction of methane with oxygen in the
air. During this reaction, chemical bonds break and re-form and the products that are
produced are carbon dioxide and water vapor (and, of course, light and heat that you
see as the flame). The unbalanced chemical equation would be written:
CH4(methane) + O2(oxygen)
CO2(carbon dioxide) + H2O(water)
Look at the reaction atom by atom. On the left side of the equation we find one carbon
atom, and one on the right.
CH4 + O2
^ 1 carbon
CO2 + H2O
^ 1 carbon
Next we move to hydrogen: There are four hydrogen atoms on the left side of the
equation, but only two on the right.
CH4 + O2
CO2 + H2O
^ 4 hydrogen
^ 2 hydrogen
Therefore, we must balance the H atoms by adding the coefficient "2" in front of the
water molecule (you can only change coefficients in a chemical equation, not
subscripts). Adding this coefficient we get:
CH4 + O2
CO2 + 2H2O
^ 4 hydrogen
^ 4 hydrogen
What this equation now says is that two molecules of water are produced for every
one molecule of methane consumed. Moving on to the oxygen atoms, we find two on
the left side of the equation, but a total of four on the right side (two from the
CO2 molecule and one from each of two water molecules H2O).
CH4 + O2
CO2 + 2H2O
^2 oxygen ^4 oxygen total
To balance the chemical equation we must add the coefficient "2" in front of the
oxygen molecule on the left side of the equation, showing that two oxygen molecules
are consumed for every one methane molecule that burns.
CH4 + 2O2
CO2 + 2H2O
^4 oxygen ^4 oxygen total
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Dalton's law of definite proportions holds true for all chemical reactions. In essence,
this law states that a chemical reaction always proceeds according to the ratio defined
by the balanced chemical equation. Thus, you can interpret the balanced methane
equation above as reading, "one part methane reacts with two parts oxygen to
produce one part carbon dioxide and two parts water." This ratio always remains the
same. For example, if we start with two parts methane, then we will consume four
parts O2 and generate two parts CO2 and four parts H2O. If we start with excess of
any of the reactants (e.g., five parts oxygen when only one part methane is available),
the excess reactant will not be consumed:
CH4 + 5O2
CO2 + 2H2O + 3O2
Excess reactants will not be consumed.
In the example seen above, 3O2 had to be added to the right side of the equation to
balance it and show that the excess oxygen is not consumed during the reaction. In
this example, methane is called the limiting reactant.
Although we have discussed balancing equations in terms of numbers of
atoms and molecules, keep in mind that we never talk about a single atom (or
molecule) when we use chemical equations. This is because single atoms (and
molecules) are so tiny that they are difficult to isolate. Chemical equations are
discussed in relation to the number of moles of reactants and products used or
produced (see our The Mole module). Because the mole refers to a standard number of
atoms (or molecules), the term can simply be substituted into chemical equations.
Thus, the balanced methane equation above can also be interpreted as reading, "one
mole of methane reacts with two moles of oxygen to produce one mole of carbon
dioxide and two moles of water."
Conservation of Mass
The law of conservation of mass states that matter is neither lost nor gained in
traditional chemical reactions; it simply changes form. Thus, if we have a certain
number of atoms of an element on the left side of an equation, we have to have the
same number on the right side. This implies that mass is also conserved during a
chemical reaction. The water reaction, for example:
2H2
+
O2
2H2O
+
2 * 2.02g
+
32.00g
=
2 * 18.02g
The total mass of the reactants, 36.04g, is exactly equal to the total mass of
the products, 36.04g. This holds true for all balanced chemical equations.
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Name:______________________________________
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Unit 7, Worksheet 2—
Balancing Chemical Equations
Balance the following equations by inserting the proper coefficients. Use a separate sheet of
paper to draw particle diagrams, if necessary. For selected reactions, draw before and after
particle diagrams to show the particles involved in the reaction. Be sure to provide a key.
1.
____C + ____H2O → ____CO + ____H2
2.
____MgO → ____Mg + ____O2
3.
____Al + ____O2 → ____Al2O3
#3
4.
____Zn + ____H2SO4 → ____ZnSO4 + ____H2
5.
____Cl2 + ____KI → ____KCl + ____I2
6.
____CuCl → ____Cu + ____Cl2
7.
____Na + ____Cl2 → ____NaCl
8.
____Al + ____HCl → ____AlCl3 + ____H2
9.
____Fe2O3 → ____Fe + ____O2
10.
____P + ____O2 → ____P2O5
11.
____Mg + ____HCl → ____MgCl2 + ____H2
12.
____H2 + ____N2 → ____NH3
13.
____BaCl2 + ____H2SO4 → ____BaSO4 + ____HCl
14.
____CH4 + ____O2 → ____CO2 + ____H2O
Before
#8
Before
#14
Before
After
After
After
15. a) ____ZnCl2 + ____(NH4)2S → ____ZnS + ____ NH4Cl
b) Find the molar mass of these reactants.
c) How many moles of ZnCl2 would be in 25 g?
(NH4)2S have?
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How much mass would 0.55 moles of
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Write the formulas of the reactants and products, then balance the equations. (See Clues and
Hints below.)
1.
Nitric oxide (NO) reacts with ozone (O3) to produce nitrogen dioxide and oxygen gas.
2.
Iron burns in air to form a black solid, Fe3O4.
3.
Sodium metal reacts with chlorine gas to form sodium chloride.
4.
Acetylene, C2H2, burns in air to form carbon dioxide and water.
5.
Hydrogen peroxide (H2O2) easily decomposes into water and oxygen gas.
6.
Hydrazine (N2H4) and hydrogen peroxide are used together as rocket fuel. The products
are nitrogen gas and water.
7.
If potassium chlorate is strongly heated, it decomposes to yield oxygen gas and potassium
chloride.
8.
When sodium hydroxide is added to sulfuric acid (H2SO4), the products are water and
sodium sulfate.
9.
In the Haber process, hydrogen gas and nitrogen gas react to form ammonia, NH3.
CLUES and HINTS:
Ø Products usually follow words like produces, yields, forms
Ø Watch for our diatomic elements (H2,N2, etc…), which are often (but not always) gases
Ø Include ‘state subscripts’ behind each substance [ (s), (l), (g) ] when the state is given
Ø Remember air is a mixture of (primarily) two gases, O2 and N2. Which is most likely to
participate in a reaction?
Ø Elemental metals exist as single, unbonded atoms. (Ex: formula for copper metal is Cu)
Ø Watch for ionic vs. molecular compounds. Use nomenclature rules, and your ion chart and
periodic table to figure out the formulas for these.
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Name:______________________________________
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Unit 7, Worksheet 3—More Balancing Chemical Equations
Balance the following equations by inserting the proper coefficients. Use a separate sheet of paper to draw
particle diagrams, if necessary.
1.
____SO2 + ____O2 → ____SO3
2.
____CH4 + ____O2 → ____CO + ____H2O
3.
____P + ____Cl2 → ____PCl3
4.
____CO + ____O2 → ____CO2
5.
____CH4 + ____ O2 → ____CH3OH
6.
____Li + ____Br2 → ____LiBr
7.
____Al2O3 → ____Al + ____O2
8.
____Na + ____H2O → ____NaOH + ____H2
9.
____CO2 + ____H2O → ____C6H12O6 + ____O2
10.
____H2SO4 + ____NaCl → ____HCl + ____Na2SO4
11.
____H2 + ____SO2 → ____H2S + ____H2O
12.
____CaCO3 + ____SO2 + ____ O2 → ____CaSO4 + ____CO2
13.
____AgNO3 + ____CaCl2 → ____AgCl + ____Ca(NO3)2
14.
____HCl + ____Ba(OH)2 → ____BaCl2 + ____H2O
15.
____H3PO4 + ____NaOH → ____Na3PO4 + ____H2O
16.
____Pb(NO3)2 + ____KI → ____ PbI2 + ____KNO3
17.
____CuO + ____NH3 → ____N2 + ____Cu + ____H2O
18.
____C2H5OH + ____O2 → ____CO2 + ____H2O
19.
____C2H6 + ____O2 → ____CH3COOH + ____H2O
20.
____NO2 + ____H2O → ____HNO3 + ____NO
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1.
When a solution of hydrogen chloride is added to solid sodium bicarbonate (NaHCO3), the
products are carbon dioxide, water and aqueous sodium chloride.
2.
Steam (gaseous water) reacts with carbon at high temperatures to produces carbon
monoxide and hydrogen gases.
3.
Limestone, CaCO3, decomposes when heated to produce lime, CaO, and gaseous carbon
dioxide.
4.
Ethyl alcohol (a liquid), C2H6O, burns in air to produce carbon dioxide and gaseous water.
5.
Solid titanium(IV) chloride reacts with water, forming solid titanium(IV) oxide and
aqueous hydrogen chloride.
6.
At high temperatures, the gases chlorine and water react to produce hydrogen chloride and
oxygen gases.
7.
Steel wool (nearly pure Fe) burns in air to form the solid iron oxide, Fe2O3.
8.
During photosynthesis in plants, carbon dioxide and water are converted into glucose,
C6H12O6, and oxygen gas.
9.
Solutions of calcium hydroxide, Ca(OH)2and nitric acid, HNO3, react to produce water and
aqueous calcium nitrate, Ca(NO3)2.
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Reading: Lavoisier and the Law of Conservation of Mass
http://www.chemteam.info/Equations/Conserv-of-Mass.html
Introduction
The Law of Conservation of Mass (or Matter) in a chemical reaction can be stated thus:
In a chemical reaction, matter is neither created nor destroyed.
It was discovered by Antoine Laurent Lavoisier (1743-94) about 1785. However, philosophical
speculation and even some quantitative experimentation preceded him. In addition, he was
certainly not the first to accept this law as true or to teach it, but he is credited as its discoverer.
Pre-history leading up to Lavoisier
Anaxagoras, circa 450 B.C. said:
"Wrongly do the Greeks suppose that aught begins or ceases to be; for nothing comes into being or
is destroyed; but all is an aggregation or secretion of pre-existing things; so that all becoming
might more correctly be called becoming mixed, and all corruption, becoming separate."
Circa 1623, Francis Bacon wrote:
"Men should frequently call upon nature to render her account; that is, when they perceive that a
body which was before manifest to the sense has escaped and disappeared, they should not admit
or liquidate the account before it hs been shown to them where the body has gone to, and into
what it has been received."
Joseph Black (1728-1799) made extensive studies of the carbonates of the alkali and alkaline earth
metals and is considered the discoverer of carbon dioxide (which he called "fixed air"). In 1752,
he wrote the following, which will be explained below:
"A piece of perfect quicklime, made from two drams of chalk, and which weighed one gram and
eight grains, was reduced to a very fine powder, and thrown into a filtered mixture of an ounce of
a fixed alkaline salt and two ounces of water. After a slight digestion, the powder being well
washed and dried, weighed one dram and fifty-eight grains. It was similar in every trial to a fine
powder of ordinary chalk, and was therefore saturated with air which must have been furnished
by the alkali."
I want you to notice that the quicklime came from two drams of chalk and at the end he produced
one dram and 58 grains of chalk. Since one dram = 60 grains, we can see there is a difference of
only 2 grains. As best as I can tell, one grain is equal to a modern value of about 0.4 grams. Here
in modern terms, are the chemical reactions Black carried out:
He made lime (CaO) from chalk (CaCO3) by heating it:
CaCO3 ---> CaO + CO2
Then, he reacted the lime with an excess of fixed alkali (K2CO3) and got back chalk:
CaO + K2CO3 ---> CaCO3 + K2O
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K2O is potassium oxide (in modern terms) and in the water would react to produce KOH, which was
called caustic alkali.
Black was interested in showing that the weight change from chalk to lime was only due to the loss
of fixed air and he never went beyond that. In fact, right before the above quote is this:
"With respect to the second proposition, . . . ."
That second proposition is as follows:
"If quick-lime be no other than a calcarious earth deprived of its air, and whose attraction for fixed
air is stronger than that of alkalis, it follows that, by adding to it a sufficient quantity of alkali
saturated with air, the lime will recover the whole of its air, and be entirely restored to its original
weight and condition . . . "
I'm not a true historian of chemistry, but I don't think Black missed the "big picture" because he was
so focused on his own agenda. The spirit of careful, quantitative measurements in chemistry was,
in the mid-1700's, still fairly new. Black was a careful experimenter, but I believe he was too
early in the game, so to speak, to recognize the Law of Conservation of Mass. To the modern
eye, his work is clear evidence for the Law of Conservation of Mass, but Black just never got to
that point.
Henry Cavendish (1731 - 1810) was one of the great chemists of the eighteenth century (or any other
century for that matter). Among his many discoveries was the composition of water and the
recognition that atmospheric air was a mixture of nitrogen and oxygen in constant proportion. In
1784, he wrote the following:
"In Dr. Priestley's last volume of experiments is related an experiment of Mr. Warltire's, in which it
is said that, on firing a mixture of common and inflammable air by electricity in a close[d]
copper vessel holding about three pints, a loss of weight was always perceived, on an average
about two grains, though the vessel was stopped in such a manner that no air could escape by the
explosion . . . . [This experiment], if there was no mistake in it, would be very extraordinary and
curious; but it did not succeed with me . . . though the experiment was repeated several times
with different proportions of common and inflammable air, I could never perceive a loss of
weight of more than one-fifth of a grain, and commonly none at all."
Cavendish adds a footnote one sentence later saying: "Dr. Priestley, I am informed, has since found
the experiment not to succeed." remember also that one gran equals about 0.4 gram, so
Canvendish, in the above quote, was discussing a weight difference of about 0.08 grams.
Cavendish is famous even today for the careful, meticulous nature of his work, but he also missed
credit for announcing the Law of Conservation of Mass. I think it was because he was taken with
other things. For example, just two paragraphs after the above is written, Cavendish begings
discussing the fact that common air (in modern terms, the atmosphere) consistently has a
maximum reduction in volume of about one-fifth after reacting with inflammable air (in modern
terms, hydrogen gas).
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Today, we know that the atmosphere is about 79% nitrogen and almost 21% oxygen, with small
amounts of other gases (carbon dioxide, water, argon, etc.). In 1784, this was a very, very
important discovery.
However, notice how he says "extraordinary and curious" in the above quote. He must have had
some awareness of what we now call the Law of Conservation of Mass, but he never announced
it as a proven, scientific principle.
The work of Lavoisier
Lavoisier wrote in 1785:
"Nothing is created, either in the operations of art or in those of nature, and it may be considered as a
general principle that in every operation there exists an equal quantity of matter before and after
the operation; that the quality and quantity of the constituents is the same, and that what happens
is only changes, modifications. It is on this principle that is founded all the art of performing
chemical experiments; in all such must be assumed a true equality or equation between
constituents of the substances examined, and those resulting from their analysis."
At this point, he was well into his scientific career. It turns out he had assumed the validity of the
law and then assembled ample verification of it before making a formal announcement.
However, there is an important point related to Lavoisier and the law. As one historian in 1914
wrote:
"What Lavoisier did, was to assume this permanency of weight to apply to the substances with
which chemists dealt, and to be independent of the effect of heat, till then supposed by many to
be ponderable." Ponderable means to have weight.
In 1890, another historian wrote:
"Lavoisier established a radical different between on the one hand ponderable matter, . . . matter of
which the balance proved the invariability before, during, and after combustion; and on the other
hand, the igneous fluid, of which the introduction from an outside source, or the withdrawal
during combustion,, neither increased nor diminished the weight of substances; contrary to what
the partisans of phlogiston has thought."
Lavoisier was able to establish that heat played no role in adding or decreasing weight, as had been
claimed by the phlogiston theory. This is not the place to discuss phlogiston, except to say it was
a chemical theory that had lasted about 100 years and was decisively destroyed by the work of
Lavoisier. (Lavoisier's prime scientific rival, Joseph Priestley of England, accepted the
phlogiston theory.)
Lavoisier was able to assemble a number of experiments, all done in closed vessels, in which the
weight remained constant, within experimental error. This included tin or lead being reacted with
oxygen as well as the analysis of mercury calx (HgO). Over the years of his work, Lavoisier had
several large burning lenses (which focused the sun's rays), constructed and these were
instrumental in reaching the high temperatures need to cause the chemical reactions to take
place. (Lavoisier was also able to burn a diamond with a large lens and show that only CO2 was
produced.)
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Your teacher will probably never require you to know the history involved, but will probably test
this statement: In a chemical reaction, matter is neither created nor destroyed.
It is the Law of Conservation of Mass. Antoine Laurent Lavoisier is its discoverer.
Final comments on the science involved
The manner in which the Law of Conservation of Mass was discovered did not follow the usual
"scientific" way that is taught to students. Lavoisier DID NOT arrive at the law by induction, that
is generalizing from a large number of specific cases. There was simply not enough data for him
to do this.
What Lavoisier did was to ASSUME the validity of the law during the course of his work and then
let the verification come from the fact that deductions from the law always - within experimental
error - showed the deduction to be correct. Another way to say it is to say that, again within
experimental error, the results of a complete analysis of a substance ALWAYS add up to 100%
of the starting material.
What is interesting is the issue of experimental error. Suppose an experiment is performed in which
mass is lost or gained. This IS NOT taken as evidence of the failure of the law, but as a failure of
the experiment. At least at the beginning, a person like Lavoisier must have had a very strong,
almost unscientific, belief that he was right, no matter what the data showed or didn't show.
This happened in 1905 with Einstein and the special theory of relativity. The very first scientific
article which dealt with relativity after Einstein announced it was by a man named Walter
Kaufmann and it CONCLUSIVELY refuted Einstein, showing him to be incorrect. Einstein was
undeterred by this, stuck to his guns and was shown to be correct, to the point where everybody
knows who Einstein was and hardly anybody remember Kaufmann.
This belief in the correctnes of your conclusion also guided Robery Milikian in 1913, when he
determined the charge on the electron to a very high degree of accuracy. His laboratory
notebooks are littered with comments like "bad value" or "something wrong, don't use." How did
he KNOW a given experimental run produced a poor value? He must have had some idea of
where he wanted to go before going there and this affected his selection of data.
If you have a desire to go into a scientific field for your career, I urge you to learn about the history
of your chosen area. There are many lessons to learn from those who went before us about how
science is done.
Very, very last comment
There actually is a better law called the Law of Conservation of Mass-Energy. Conservation of mass
was amended due to the discovery of E = mc2 by Einstein. We also know that 100 kJ = about
10¯9 gram and in these modern times, that is very near to the detection limit of some of the better
mass spectroscopy instruments in the world. I have heard that this tiny mass loss (actually
conversion of mass to energy) in a chemical reaction has been detected, but I do not have a
journal reference for this.
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Unit 7, Lab 2
Title:
Purpose/Question:
Procedure:
Carry out the reactions using the approximate quantities of reagents suggested. Unless
otherwise stated, use test tubes. When heating reagents in test tubes, slant the test tube so
that the opening is pointed away from people. Heat the test tube at the surface of the
material and work down towards the bottom of the tube. Discard solutions down the drain,
wash and rinse your glassware. Discard solid waste in the waste cans on the lab tables. In
the data section you will balance the equation, write the word equation and record your
observations.
A. Combination reactions:
1a. Grasp a strip of magnesium ribbon in crucible tongs and ignite it in the burner flame. Hold
it over a watch glass (do not drop it in the watch glass until reaction complete). Do not look
directly at the flame!
1b. Add a few drops of distilled H2O to the ash in the watch glass. Stir with a stirring rod and
place a drop of the solution on red litmus paper. Red litmus turning blue is evidence for the
presence of a base.
3. Heat a piece of copper metal strongly in the Bunsen burner flame for about 30 s. Remove
the copper from the flame and note the change in appearance. Discard the product in the
solid waste can.
B. Decomposition reactions:
1. Place about 1 scoopful of solid sodium hydrogen carbonate NaHCO3 into a dry test tube.
Mass the test tube with the powder. Heat the sodium hydrogen carbonate in the test tube
strongly for 2 minutes. Observe any changes that occur during the heating. Toward the end
of the heating, light a wood splint and insert the flaming splint into the mouth of the test
tube. Note what happens to the splint. Once the tube has cooled, mass the tube and
contents again.
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C. Single replacement reactions:
1. Place a small strip of aluminum wire in a test tube with enough copper (II) chloride solution
to cover it. Set this test tube aside, then observe the surface of the metal at the conclusion
of the lab.
2. Place a couple of pieces of mossy zinc metal in a test tube approximately 1/4 full of
3M HCl. Place a stopper loosely in the tube. After a few minutes, light a wood splint and
insert the flaming splint into the mouth of the test tube. Hold the test tube in your hand to
feel if the temperature has changed.
D. Double replacement reactions:
1. Add 0.1M AgNO3 to a test tube to a depth of about 1 cm. Add a similar quantity of 0.1M
CaCl2 solution. Observe the reaction.
2. Place a scoopful of solid Na2CO3 in a test tube to a depth of about 1 cm. Add a dropperful
of 3M HCl. While the reaction is occurring, test with a flaming splint as in part B.
Check to see if the temperature of the mixture has changed.
E. Combustion reactions:
Place about 10 drops of isopropyl alcohol, C3H7OH, in a small evaporating dish. Ignite the
alcohol from the top of the liquid with a Bunsen burner. Hold a cold watch glass well above
the flame and observe the condensation of water on the bottom. The formation of the mist
will be fleeting; watch closely.
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Data and Analysis:
Equations written in words should include the following: number of moles of each substance, state of
matter of each substance, and full nomenclature (no symbols).
A. Combination reactions:
1a. Observations:
Mg +
O2 à MgO
Write equation in words:
1b. Observations:
MgO +
H2O à
Mg(OH)2 (aq)
Write in words:
2. Observations:
___Cu +
O2 -->
CuO
Write in words:
B. Decomposition reactions
1. Observations:
NaHCO3 à
Na2O + ______ H2O +
CO2 (g)
Write in words:
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C. Single replacement reactions
1. Observations:
AgNO3 (aq) +
Cu à
Ag +
Cu(NO3)2 (aq)
Write in words:
2. Observations
Zn +
HCl(aq) à
ZnCl2(aq) + H2(g)
Write in words:
D. Double replacement reactions
1. Observations
AgNO3(aq) +
CaCl2(aq) à AgCl(s) +
Ca(NO3)2(aq)
Write in words:
2. Observations
Na2CO3 +
HCl (aq) à NaCl(aq) +
H 2O +
CO2(g)
Write in words:
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E. Combustion reactions
1. Observations
C3H7OH(l) +
O2(g) à
CO2(g) +
H2O(g)
Write in words:
Post-Lab Questions
1. What are several examples of observable changes that are evidence that a chemical
reaction has taken place?
2. How did the flaming splint behave when it was inserted into the tube with CO2 (g)?
In what way was this different from the reaction of the H2(g) to the flaming splint?
3. In the reaction of magnesium with oxygen gas, a considerable amount of energy was
released. This is an example of an exothermic reaction. Write the balanced chemical
equation for this reaction below:
If energy was released, what can you conclude about the amount of energy stored in the
reactants compared to the amount of energy stored in the product? Explain.
What other examples of exothermic reactions did you observe?
Re-write the balanced equation for the reaction of Mg and O2, this time with the term “+
energy” on the appropriate side of the equation.
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4. You had to heat the NaHCO3 strongly in order for it to decompose. This is an example of an
endothermic reaction. Write the balanced chemical equation for this reaction below:
If you had to add energy to the system for the reaction to proceed, what does this tell you
about the amount of energy stored in the reactant compared to the amount of energy stored
in the products? Explain.
Write the balanced equation for the decomposition of NaHCO3, this time with the term “+
energy” on the appropriate side of the equation.
5. Develop a set of “rules” that define the following types of chemical reactions: combination
(synthesis), decomposition, single replacement, double replacement, and combustion.
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Unit 7, Worksheet 4—
Reaction Types and More Balancing
Balance the following equations. If the equation is already balanced, write “balanced” after it. In
the blank, identify the type of reaction (combination/synthesis, decomposition, single replacement,
double replacement, combustion)
1.
_______________
H2 +
2.
_______________
N2 +
3.
_______________
S8 +
4.
_______________
N2 +
5.
_______________
HgO à
6.
_______________
Zn +
HCl à
ZnCl2 +
7.
_______________
Na +
H2O à
NaOH +
8.
_______________
H3PO4 à
9.
_______________
C10H16 +
10.
_______________
Al(OH)3 +
11.
_______________
Fe +
12.
_______________
Fe2(SO4)3 +
13.
_______________
C 7H 6O 2 +
14.
_______________
FeS2 +
O2 à
15.
_______________
Al +
FeO à
16.
_______________
Fe2O3 +
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O2 à
H 2O
H2 à
O2 à
NH3
SO3
O2 à
N 2O
Hg +
O2
H2
H 4P 2O 7 +
Cl2 à
O2 à
H 2O
C+
H2SO4 à
H2
HCl
Al2(SO4)3 +
H 2O
Fe2O3
KOH à
O2 à
H2 à
K2SO4 +
CO2 +
Fe2O3 +
Al2O3 +
Fe +
Fe(OH)3
H 2O
SO2
Fe
H 2O
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Unit 7, Worksheet 5—
Writing Balanced Chemical Equations
Write balanced chemical equations for the following reactions.
1.
Ammonia (NH3) reacts with hydrogen chloride to form ammonium chloride.
2.
Calcium carbonate decomposes upon heating to form calcium oxide and carbon dioxide.
3.
Barium oxide reacts with water to form barium hydroxide.
4.
Acetaldehyde (CH3CHO) decomposes to form methane (CH4) and carbon monoxide.
5.
Zinc reacts with copper(II) nitrate to form zinc nitrate and copper.
6.
Calcium sulfite decomposes when heated to form calcium oxide and sulfur dioxide.
7.
Iron reacts with sulfuric acid (H2SO4) to form iron(II) sulfate and hydrogen gas.
8.
A nitrogen containing carbon compound, C2H6N2, decomposes to form ethane, C2H6, and
nitrogen gas.
9.
Phosgene, COCl2, is formed when carbon monoxide reacts with chlorine gas.
10. Manganese(II) iodide decomposes when exposed to light to form manganese and iodine.
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11. Dinitrogen pentoxide reacts with water to produce nitric acid (HNO3).
12. Magnesium reacts with titanium (IV) chloride to produce magnesium chloride and
titanium.
13. Carbon reacts with zinc oxide to produce zinc and carbon dioxide.
14. Bromine reacts with sodium iodide to form sodium bromide and iodine.
15. Phosphorus (P4) reacts with bromine to produce phosphorus tribromide.
16. Ethanol, C2H5OH, reacts with oxygen gas to produce carbon dioxide and water.
17. Calcium hydride reacts with water to produce calcium hydroxide and hydrogen gas.
18. Sulfuric acid, H2SO4, reacts with potassium hydroxide to produce potassium sulfate and
water.
19. Propane, C3H8, burns in air to produce carbon dioxide and water.
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Unit 7 — Free Notes Page
(Activity Series)
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Unit 7, Worksheet 6—
Activity Series Online Simulation
Directions:
1. Go to:
http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/redox/home.
html
2. Click on start and follow the directions for Activity 1. After filling in the data table and
listing the metals in order of reactivity, do activity 2, 3, and 4.
Activity 1
Zn(NO3)2
Zn
Mg(NO3)2
Cu(NO3)2
AgNO3
Mg
Ag
Cu
List the metals in order of most reactive to least reactive:
Write the balanced equation and create a particle drawing for three reactions.
1-
2-
3-
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Activity 2
Fe(NO3)2
Zn(NO3)2
Cu(NO3)2
Pb(NO3)2
Fe
Zn
Pb
Cu
List the metals in order of most reactive to least reactive:
Write the balanced equation and create a particle drawing for three reactions.
1-
2-
3-
Activity 3
Fe(NO3)2
Pb(NO3)2
Ni(NO3)2
Sn(NO3)2
Fe
Pb
Ni
Sn
List the metals in order of most reactive to least reactive:
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Write the balanced equation and create a particle drawing for three reactions.
1-
2-
3-
Activity 4
Document the amount of bubbling and how much metal is lost
Ag
Cu
Fe
Mg
Ni
Pb
HCl
Sn
Zn
List the metals in order from most reactive to least reactive in acid.
Conclusion:
1. Merge the four lists together to create a reactivity series of metals.
2. Do all single displacement reactions occur?
3. How would you be able to tell if an elemental metal could displace a metal cation from
solution?
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Unit 7, Worksheet 7—
Single Replacement Reaction Predictions with Activity Series
Write balanced chemical equations for the following reactions. If no reaction will occur, write
“No Reaction” after the arrow.
1. A copper wire is placed in a solution of calcium nitrate.
2. Sodium fluoride is added to iodine.
3. A strip of magnesium is added to a solution of iron III chloride.
4. Chlorine gas in bubbled through a solution of copper II bromide.
5. An aluminum sulfate solution is added to a beaker containing iron filings.
6. Potassium sulfate solution is added to a beaker containing iron filings
7. A magnesium chloride solution is added to calcium.
8. Sodium is placed into water.
9. Magnesium is placed into liquid water.
10. Magnesium is placed into steam.
11. Zinc is added to sulfuric acid (H2SO4).
12. Silver is added to hydrochloric acid (HCl)
13. Barium is added to cold water.
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14. Powdered aluminum metal is added to a zinc nitrate solution.
15. Sodium carbonate decomposes when it is heated.
16. Iron combines with oxygen.
17. Calcium hydroxide solution is mixed with a hydrofluoric acid solution.
18. Copper wire is added to a sodium sulfide solution.
19. Propane (C3H8) is burned.
20. Chlorine gas is bubbled through a solution of magnesium iodide.
21. Sodium chloride undergoes electrolysis.
22. Sulfur dioxide is added to water.
23. Sodium oxide is added to water.
24. Zinc metal is added to a hydrochloric acid solution.
25. Nitrogen and oxygen combine.
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Difference Between Exothermic and Endothermic
http://www.differencebetween.net/science/difference-between-exothermic-and-endothermic/
Exothermic and Endothermic
In chemistry we have learned about exothermic and endothermic reactions. But how it is applicable
in our daily lives is not known to many.
Firstly, an exothermic reaction is one in which heat is produced as one of the end
products. Examples of exothermic reactions from our daily life are combustion like the burning of
a candle, wood, and neutralization reactions. In an endothermic reaction, the opposite happens. In
this reaction, heat is absorbed. Or more exactly, heat is required to complete the reaction.
Photosynthesis in plants is a chemical endothermic reaction. In this process, the chloroplasts in the
leaves absorb the sunlight. Without sunlight or some other similar source of energy, this reaction
cannot be completed.
In exothermic reactions the enthalpy change is always negative while in endothermic reactions the
enthalpy change is always positive. This is due to the releasing and absorption of heat energy in the
reactions, respectively. The end products are stable in exothermic reactions. The end products of
endothermic reactions are less stable. This is due to the weak bonds formed.
‘Endo’ means to absorb and so in endothermic reactions, the energy is
absorbed from the external surrounding environment. So the
surroundings lose energy and as a result the end product has higher
energy level than the reactants. Due to this higher energy bonds, the
product is less stable. And most of the endothermic reactions are not
spontaneous. ‘Exo’ means to give off and so energy is liberated in
exothermic reactions. As a result, the surroundings get heated up. And
most exothermic reactions are spontaneous.
When we light a matchstick, it is an exothermic reaction. In this reaction, when we strike the stick,
stored energy is released as heat spontaneously. And the flame will have lower energy than the heat
produced. The energy being released is previously stored in the matchstick and thus do not require
any external energy for the reaction to occur.
When ice melts, it will be due to the heat around. The surrounding environment will have a higher
temperature than the ice and this heat energy is absorbed by the ice. The stability of the bonds is
reduced and as a result and the ice melts into liquid.
Some exothermic reactions in our lives are the digestion of food in our body, combustion reactions,
water condensations, bomb explosions, and adding an alkali metal to water. So now you must have
an idea of what exothermic and endothermic reactions are.
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Energy of Combustion
http://www.elmhurst.edu/~chm/vchembook/512energycombust.html
The combustion of all fossil fuels follows a very similar reaction:
Fossil Fuel (any hydrocarbon source) plus oxygen yields carbon dioxide and water and ENERGY.
The world and modern society are driven by the need to produce energy to make products
(manufacturing), to move around (transportation), to heat homes and buildings, and to create
light (electricity). At least 75% of these needs are met by the combustion of fossil fuels. Energy
is stored in chemical compounds in the bonds that bind atoms to each other.
CH4[g] + 2 O2[g] -> CO2[g] + 2 H2O[g] + ENERGY
A chemical reaction occurs by the rearrangement of atoms and molecules in the reactant
(starting) molecules and the end product molecules. Some bonds are broken while others are
reformed. The process of breaking and forming bonds results in a net energy needed or given off
for a reaction.
In the example above and to the left, the combustion reaction of methane and oxygen to form carbon
dioxide and water is shown broken into steps to show the entire energy "using" and "forming"
process. First it takes energy to break bonds, all four of the C-H bonds in methane must be
broken. The energy units are kilojoules, a positive sign means that the process is endothermic or
energy is required to break the bonds.
In a similar fashion, two diatomic oxygen molecules are broken apart which requires more energy.
Now all of the individual atoms in the reactant molecules have been broken apart.
On the right side of the diagram in a second step, the various atoms form new bonds in new
molecules of carbon dioxide and water. The formation of new bonds is an exothermic process
where heat is given off. Again the energy given off is totaled to form new bonds in carbon
dioxide and water molecules.
Finally, the overall reaction yields an excess of energy given off -802 kj. (the minus sign means that
this is an exothermic process). In more familiar units this is equivalent to 191 kilocalories per 16
grams of methane. This is a little more than the 150 calories in a can of Coke.
The excess of energy given off is mainly in the form of heat. Chemical energy stored in the bonds of
molecules is transformed into heat and light energy. Most chemical reactions are of this type and
thus are exothermic. Less energy is required to break old bonds than is given off in the process of
forming new bonds.
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Unit 7, Worksheet 8—
Representing Chemical Potential Energy in Change
For#1-2, write the balanced chemical equation, including the energy term on the correct side
of the equation. Then represent the energy storage and transfer using the bar graphs. Below
the bar graph diagram, sketch a standard chemical potential energy curve for the reaction.
1.
When you heated sodium hydrogen carbonate, you decomposed it into sodium oxide, water
vapor, and gaseous carbon dioxide.
2. When solid zinc was added to hydrochloric acid, the products were hydrogen gas and an
aqueous solution of zinc chloride. You could feel the test tube get hotter.
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3.
Isopropyl alcohol burned in air to produce carbon dioxide and water vapor.
4. In chemical cold packs, solid ammonium chloride dissolves in water forming aqueous
ammonium and chloride ions. As a result of this solvation reaction, the pack feel cold on
your injured ankle.
5. In chemical hot packs, solid sodium acetate crystallizes from a supersaturated solution of
sodium acetate. The pack feels warm to the touch for 30 minutes or longer.
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Unit 7— More Practice Problems
Conservation of Mass and Atoms
The total number of atoms of each kind must remain the same from the beginning to the end
of a reaction.
1.
Aqueous iron (III) nitrate reacts with aqueous sodium hydroxide in a double
replacement reaction
Balanced Equation:
_____________________________________________________________________
Number of iron atoms
________
Number of nitrogen atoms
________
Number of oxygen atoms
________
Number of sodium atoms
________
Number of hydrogen atoms
________
Total number of atoms on each side of the equation ________
2.
In the spaces given below, draw particle diagrams that show all reactants and all
products in the reaction described in problem 1.
è
Reactants
Products
3. Solid aluminum reacts with aqueous zinc sulfate in a single replacement reaction
Balanced Equation:
____________________________________________________________________
Number of aluminum atoms
________
Number of zinc atoms
________
Number of sulfur atoms
________
Number of oxygen atoms
________
Total number of atoms on each side of the equation __________
4.
In the spaces given below, draw particle diagrams that show all reactants and all
products in the reaction described in problem 3.
à
Reactants
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Classifying Reactions
Reactions can be classified into different types. Insert the proper coefficients to balance each
chemical equation. Use the codes given below to classify each reaction.
S
D
SR
DR
C
Synthesis
Decomposition
Single Replacement
Double Replacement
Combustion
A + B = AB
AB à A + B
A + BC à B + AC
AB + CD à AD + CB
Fuel + O2 à products
CODE
______ 5.
_____Al2O3 (s) → _____Al(s) + _____O2 (g)
______ 6.
_____C2H5OH (l)
______ 7.
_____Ca(s) +
______ 8.
_____H3PO4(aq)
______ 9.
_____Cu(s)
+
_____I2 (s) à _____ CuI(s)
______ 10.
_____Na(s)
+
_____NiCl2 (aq)
______ 11.
_____Sr(OH)2(aq)
______ 12.
_____HCl
+
_____ O2 (g) à _____CO2 (g)
_____O2(g)
+ _____H2O (g)
à _____ CaO(s)
+ ______NaOH(aq) → ______Na3PO4(aq) + ______H2O(l)
+
à
_____FeCl3(aq)
+ _____Ba(OH)2
→
_____ NaCl(aq)
à
+
_____Ni(s)
_____SrCl2(aq)
_____ BaCl2
+
+
_____Fe(OH)3 (s)
_____ H2
Predicting Products and Writing Balanced Chemical Equations
Be sure to include states in the equation.
13. Solid zinc metal is placed into an aqueous solution of silver nitrate.
______________________________________________________________________________________
14. Heptane (C7H16) burns readily in the presence of air.
______________________________________________________________________________________
15. Aqueous calcium chloride is mixed with sulfuric acid.
______________________________________________________________________________________
16. Phosphorus trichloride decomposes when heated strongly in a test tube.
______________________________________________________________________________________
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Energy Changes During Reactions
All compounds contain a certain amount of stored chemical potential energy (Ech). Some
chemical reactions release energy when reactants rearrange into products. This means
that the total amount of energy stored in the reactant molecules was greater than the total
amount of energy stored in the product molecules after rearrangement. These reactions are
called exothermic (or exergonic) reactions. Other reactions require an input of energy before
reactants rearrange into products. The input of energy is necessary because the products
contain more total energy than did the reactants. Without an input of energy, it is not
possible for reactants to rearrange into products. These reactions that require an input of
energy are called endothermic (or endergonic) reactions. Energy changes during a chemical
reaction can be tracked and summarized on the energy bar chart.
27. The elegant flaming dessert called cherry jubilee is a favorite in many upscale restaurants.
A liquor that contains ethyl alcohol is poured over the dessert just before it is served. One
spark from a match provides the activation energy to initiate the reaction. Write a
balanced chemical reaction for the reaction that occurs when the liquid ethyl alcohol
(C2H5OH) on top of the cherry jubilee burns in air. Add the + energy term to the side of
the equation where it belongs.
____________________________________________________________________________
Is this reaction endothermic or exothermic?____________________________________
Draw the energy bar chart for this reaction.
Describe the energy transfers that occur: _____________________________________
____________________________________________________________________________
____________________________________________________________________________
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29. Solid barium hydroxide octahydrate reacts with an aqueous solution of ammonium
thiocyanate forming aqueous barium thiocyanate, aqueous ammonia (NH3) and liquid
water. The flask becomes cold enough that the liquid water freezes into solid ice. Write
the balanced chemical equation for this reaction. Add the + energy term to the side of the
equation where it belongs.
_____________________________________________________________________________
Is this reaction endothermic or exothermic? ___________________________________
Draw the energy bar chart for this reaction.
Describe the energy transfers that occur: _____________________________________
____________________________________________________________________________
Even More Writing Equations Practice
Write and balance the follow reactions. Predict products, if necessary. Be sure to include states
in the equation. If there is no chemical reaction, write “no reaction.”
1. Zinc and lead (II) nitrate react to form zinc nitrate and lead.
2. Aluminum bromide and chlorine gas react to form aluminum chloride and bromine gas.
3.
Sodium phosphate and calcium chloride react to form calcium phosphate and sodium
chloride.
4. Potassium metal and chlorine gas combine to form potassium chloride.
5. Aluminum and hydrochloric acid (HCl) react to form aluminum chloride and hydrogen
gas.
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6. Calcium hydroxide and phosphoric acid (H3PO4) react to form calcium phosphate and
water.
7. Copper and sulfuric acid (H2SO4) react to form copper (II) sulfate and water and sulfur
dioxide.
8. Hydrogen gas and nitrogen monoxide react to form water and nitrogen gas.
9. The reaction of ammonia (NH3) with iodine to form nitrogen triiodide (NI3) and
hydrogen gas.
10. The combustion of propane (C3H8) to form carbon dioxide and water.
11. The reaction of copper (II) oxide with hydrogen to form copper metal and water.
12. The reaction of iron metal with oxygen to form iron (III) oxide.
13. The reaction of AlBr3 with Mg(OH)2
14. The decomposition of hydrogen peroxide to form water and oxygen.
15. When lithium hydroxide pellets are added to a solution of sulfuric acid, lithium sulfate
and water are formed.
16. When dirty water is boiled for purification purposes, the temperature is brought up to
1000 C for 15 minutes.
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17. If a copper coil is placed into a solution of silver nitrate, silver crystals form on the
surface of the copper. Additionally, highly soluble copper (I) nitrate is generated.
18. When crystalline C6H12O6 is burned in oxygen, carbon dioxide and water vapor are
formed.
19. When a chunk of palladium metal is ground into a very fine powder and heated to drive
off any atmospheric moisture, the resulting powder is an excellent catalyst for chemical
reactions.
Even More Classifying Reactions
Identify the type of reaction as: single replacement, double replacement, combustion,
combination/synthesis, or decomposition
1) Na3PO4 + 3 KOH à 3 NaOH + K3PO4
_________________________
2) MgCl2 + Li2CO3 à MgCO3 + 2 LiCl
_________________________
3) C6H12 + 9 O2 à 6 CO2 + 6 H2O
_________________________
4) Pb + FeSO4 à PbSO4 + Fe
_________________________
5) CaCO3 à CaO + CO2
_________________________
6) P4 + 3 O2 à 2 P2O3
_________________________
7) 2 RbNO3 + BeF2 à Be(NO3)2 + 2 RbF
__________________________
8) 2 AgNO3 + Cu à Cu(NO3)2 + 2 Ag
__________________________
9) C3H6O + 4 O2 à 3 CO2 + 3 H2O
_________________________
10) 2 C5H5 + Fe à Fe(C5H5)2
_________________________
11) SeCl6 + O2 à SeO2 + 3Cl2
_________________________
12) 2 MgI2 + Mn(SO3)2 à 2 MgSO3 + MnI4
_________________________
13) O3 à O. + O2
_________________________
14) 2 NO2 à 2 O2 + N2
_________________________
Modeling Chemistry
TN Modeling Curriculum Committee
Pope John Paul II High School
268
Even More Balancing and Classifying Reactions
Balance the following equations. Identify the type of reaction as: single replacement, double
replacement, combustion, combination/synthesis, or decomposition
1) ____ NaBr + ____ Ca(OH)2 à ___ CaBr2 + ____ NaOH
Type of reaction: _____________________________
2) ____ NH3+ ____ H2SO4 à ____ (NH4)2SO4
Type of reaction: _____________________________
3) ____ C5H9O + ____ O2 à ____ CO2 + ____ H2O
Type of reaction: _____________________________
4) ____ Pb + ____ H3PO4 à ____ H2 + ____ Pb3(PO4)2
Type of reaction: _____________________________
5) ____ Li3N + ____ NH4NO3 à ___ LiNO3 + ___ (NH4)3N
Type of reaction: _____________________________
6) ____ HBr + ___ Al(OH)3 à ___ H2O + ___ AlBr3
Type of reaction: _____________________________
Yes, that is right…Even More Balancing Chemical Equations
Balance the following equations. Identify the type of reaction as: single replacement, double
replacement, combustion, combination/synthesis, or decomposition
1) ____ C6H6 + ____ O2 à ____ H2O + ____ CO2
2) ____ NaI + ____ Pb(SO4)2 à ____ PbI4 + ____ Na2SO4
3) ____ NH3 + ____ O2 à____ NO + ____ H2O
4) ____ Fe(OH)3 à ____ Fe2O3 + ____ H2O
5) ____ HNO3 + ____ Mg(OH)2 à ____H2O + ____ Mg(NO3)2
6) ____ H3PO4 + ____ NaBr à ____ HBr + ____ Na3PO4
Modeling Chemistry
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269
7) ____ C + ____ H2 à ____ C3H8
8) ____ CaO + ____ MnI4 à ____ MnO2 + ____ CaI2
9) ____ Fe2O3 + ____ H2O à ____ Fe(OH)3
10) ____ C2H2 + ____ H2 à ____ C2H6
11) ____ VF5 + ____ HI à ____ V2I10 + ___ HF
12) ____ OsO4 + ____ PtCl4 à ____ PtO2 + ____ OsCl8
13) ____ CF4 + ____ Br2 à ___ CBr4 + ____ F2
14) ____ Hg2I2 + ____ O2 à ____ Hg2O + ____ I2
15) ____ Y(NO3)2 + ____ GaPO4 à ____ YPO4 + ____ Ga(NO3)2
More Predicting Products and Balacing!
Using the activity series, predict the products of the following equations. Balance the equations.
1) ___ Ag + ___CuSO4 à
2) ___ LiNO3 + ___Ag à
3) ___ Zn + ___Au(NO2)2 à
4) ___Ag + ___KNO3 à
5) ___Zn + ___AgNO3 à
5) ___Cl2 + ___KI à
6) ___ Cu + ___FeSO4 à
7) ____Al + ____H2SO4 à
8) ____ Al + ____CuCl2 à
Modeling Chemistry
TN Modeling Curriculum Committee
Pope John Paul II High School
270