Unit 7, Lab 1 - JPII - Faith Leads Us Beyond Ourselvesjp2hs.org/wp-content/uploads/2014/08/09_Unit-7-Course-Pack.pdf · Unit 7, Lab 1 We continue to build ... 2. Record the mass of

Modeling Chemistry TN Modeling Curriculum Committee Pope John Paul II High School

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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

€

43

;

0.67 can be expressed as

€

23

, etc.

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

Diagram: 2. _____H2 + _____Cl2 → _____ HCl

Diagram: 3. _____Na + _____O2 → _____ Na2O

Diagram: 4. _____N2 + _____H2 → _____ NH3

Diagram: 5. _____CH4 + _____O2 → _____CO2 + _____H2O

Diagram: 6. _____NO + _____O2 → _____ NO2

Diagram: 7. _____Fe + _____Cl2 → _____ FeCl3

Diagram: 8. ____CH3OH + _____O2 → _____CO2 + _____H2O

Diagram:


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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 H2O

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 H2O

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 CO2 + H2O ^ 1 carbon ^ 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:______________________________________

Veritas:______________________________________

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 Before After 4. ____Zn + ____H2SO4 → ____ZnSO4 + ____H2 5. ____Cl2 + ____KI → ____KCl + ____I2 6. ____CuCl → ____Cu + ____Cl2 7. ____Na + ____Cl2 → ____NaCl 8. ____Al + ____HCl → ____AlCl3 + ____H2 #8 Before After 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 #14 Before 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? How much mass would 0.55 moles of

(NH4)2S have?


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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:______________________________________

Veritas:______________________________________

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) + H2O + 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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Name:______________________________________

Veritas:______________________________________

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 + O2 à H2O

2. _______________ N2 + H2 à NH3

3. _______________ S8 + O2 à SO3

4. _______________ N2 + O2 à N2O

5. _______________ HgO à Hg + O2

6. _______________ Zn + HCl à ZnCl2 + H2

7. _______________ Na + H2O à NaOH + H2

8. _______________ H3PO4 à H4P2O7 + H2O

9. _______________ C10H16 + Cl2 à C + HCl

10. _______________ Al(OH)3 + H2SO4 à Al2(SO4)3 + H2O

11. _______________ Fe + O2 à Fe2O3

12. _______________ Fe2(SO4)3 + KOH à K2SO4 + Fe(OH)3

13. _______________ C7H6O2 + O2 à CO2 + H2O

14. _______________ FeS2 + O2 à Fe2O3 + SO2

15. _______________ Al + FeO à Al2O3 + Fe

16. _______________ Fe2O3 + H2 à Fe + H2O


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Name:______________________________________

Veritas:______________________________________

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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Name:______________________________________

Veritas:______________________________________

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 Mg(NO3)2 Cu(NO3)2 AgNO3

Zn

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 Sn Zn HCl

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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Name:______________________________________

Veritas:______________________________________

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.Â 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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Name:______________________________________

Veritas:______________________________________

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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Name:______________________________________

Veritas:______________________________________

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 Products


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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 Synthesis A + B = AB D Decomposition AB à A + B SR Single Replacement A + BC à B + AC DR Double Replacement AB + CD à AD + CB C Combustion Fuel + O2 à products CODE ______ 5. _____Al2O3 (s) → _____Al(s) + _____O2 (g)

______ 6. _____C2H5OH (l) + _____ O2 (g) à _____CO2 (g) + _____H2O (g)

______ 7. _____Ca(s) + _____O2(g) à _____ CaO(s)

______ 8. _____H3PO4(aq) + ______NaOH(aq) → ______Na3PO4(aq) + ______H2O(l)

______ 9. _____Cu(s) + _____I2 (s) à _____ CuI(s)

______ 10. _____Na(s) + _____NiCl2 (aq) à _____ NaCl(aq) + _____Ni(s)

______ 11. _____Sr(OH)2(aq) + _____FeCl3(aq) à _____SrCl2(aq) + _____Fe(OH)3 (s)

______ 12. _____HCl + _____Ba(OH)2 → _____ BaCl2 + _____ 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 _________________________


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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


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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

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