Download unit 6 - writing and balancing chemical equations

Chemistry
2013-2014
Chemical Equations
& Reactions
Chemistry
Unit 6
Dec. 2-12, 2013
The Law of Conservation of Mass states that matter can be changed from one form into
another, mixtures can be separated or made, and pure substances can be decomposed, but the
total amount of mass remains constant. We can state this important law in another way. The total
mass of the universe is constant within measurable limits; whenever matter undergoes a change,
the total mass of the products of the change is, within measurable limits, the same as the total
mass of the reactants.
The formulation of this law near the end of the eighteenth century marked the beginning
of modern chemistry. By that time many elements had been isolated and identified, most notably
oxygen, nitrogen, and hydrogen. It was also known that, when a pure metal was heated in air, it
became what was then called a calx (which we now call an oxide) and that this change was
accompanied by an increase in mass. The reverse of this reaction was also known: Many calxes
on heating lost mass and returned to pure metals. Many imaginative explanations of these mass
changes were proposed. Antoine Lavoisier (1743-1794), a French nobleman later guillotined in
the revolution, was an amateur chemist with a remarkably analytical mind. He considered the
properties of metals and then carried out a series of experiments designed to allow him to
measure not just the mass of the metal and the calx but also the mass of the air surrounding the
reaction. His results showed that the mass gained by the metal in forming the calx was equal to
the mass lost by the surrounding air.
With this simple experiment, in which accurate measurement was critical to the correct
interpretation of the results, Lavoisier established the Law of Conservation of Mass, and
chemistry became an exact science, one based on careful measurement. For his pioneering work
in the establishment of that law and his analytical approach to experimentation, Lavoisier has
been called the father of modern chemistry.1
2
1
2
http://chem.wisc.edu/deptfiles/genchem/sstutorial/Text1/Tx14/tx14.html
http://sceti.library.upenn.edu/sceti/smith/scientist.cfm?PictureID=2305&ScientistID=172
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Before you can write or balance a chemical equation, you must first be able to count the
atoms that are present. A subscript multiplies only what is immediately in front of it. A
coefficient placed in front of a molecule or a formula unit multiplies every atom in the formula
by that number.
Example 1: How many oxygen atoms are there in a formula unit of ferric dichromate:
1Fe+3Cr2O7-2  Fe2(Cr2O7)3  (7)3 x 1 = 21 oxygen atoms
(b) in 5 formula units of ferric dichromate?
5Fe2(Cr2O7)3  5 x (7)3 = 105 oxygen atoms
Example 2: How many total atoms are there in 6 formula units of barium chloride?
6Ba+2Cl1-  6BaCl2  6Ba + 6x2Cl = 6 + 12 = 18 total atoms
Example 3: How many elements are present in 3 formula units of aluminum bicarbonate?
3Al+3HCO31-  3Al(HCO3)3  Al, H, C, O = 4 elements
A chemical equation is exactly what it says it is – an equality
between the reactants (which are substances written on the left
side of the equation) and the products (which are substances
written on the right side). An arrow pointing to the right
serves as the = sign and it is read "yields".
What is a chemical equation?3
In chemistry, we use symbols to represent the various
chemicals. Success in chemistry depends upon developing a
strong familiarity with these basic symbols. A chemical
equation is an expression of a chemical process.
For example:
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
In this equation, AgNO3 is mixed with NaCl. The equation
shows that the reactants (AgNO3 and NaCl) react through some process () to form the products
(AgCl and NaNO3). Since they undergo a chemical process, they are changed fundamentally.
Coefficients are used in all chemical equations to show the relative amounts of each substance
present. This amount can represent either the relative number of molecules, or the relative
number of moles (described below). If no coefficient is shown, a one (1) is assumed. A
coefficient ALWAYS goes in FRONT of a compound and is a multiplier for that compound.
3
http://www.shodor.org/UNChem/basic/stoic/index.html#whatisc
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On some occasions, a variety of information will be written above or below the arrows. This
information, such as a value for temperature, shows what conditions need to be present for a
reaction to occur. For example, in the graphic below, the notation above and below the arrows
shows that we need a chemical, Fe2O3, a temperature of 1000C, and a pressure of 500
atmospheres for this reaction to occur.
There are six physical state symbols that are used as subscripts immediately following substances
whose physical states are known or given. They are:
(s) – solid, which is used for solids or precipitates
(l) – liquid, which is used only for "true" liquids such as elements which are liquids at
room temperature such as mercury and bromine; also for water, and for molten
(melted) substances
(g) – gas, which is used for gases or vapors.
(aq) – aqueous solution, which means that the compound is dissolved in water, making a
solution
(cr) – crystal, which means that the substance is in its crystalline form
(pr) – precipitate, which is used only on the product side of the equation
HOW TO WRITE A BALANCED CHEMICAL EQUATION:
(1) Write a correct formula for each of the reactants. A plus sign means "added to" or
"reacts with". Put a plus sign between the reactants to separate them.
(2) Draw the yield arrow
(3) Write a correct formula for each of the products, putting a plus sign between them
also. On the products side, the plus sign means "as well as" or "in addition to"
(4) Balance the metals first, putting coefficients where necessary
(5) Balance the polyatomic ions next (IF THEY STAY TOGETHER-- that is, if the same
polyatomic ion shows up on each side of the equation) by adding a coefficient in front
of the entire compound (not in the middle of the compound)
(6) If the polyatomic comes apart OR if there is no polyatomic present, balance the nonmetals except hydrogen and oxygen next by adding coefficients where necessary
(7) Balance the hydrogen atoms and the oxygen atoms (which were not part of the
polyatomic ion) last
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The graphic below works to capture most of the concepts described above:
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Example 4: Write a balanced equation for the following reactions:
(a) A solid piece of zinc reacts with hydrochloric acid to produce a solution of zinc
chloride and hydrogen gas as the products
solid zinc is Zn(s)
reacts with is (+)
hydrochloric acid is HCl
to produce is 
solution zinc chloride is Zn+2Cl1-(aq)  ZnCl2(aq)
and is +
hydrogen gas is H2 (diatomic!)
Zn + HCl  ZnCl2 + H2
(b) Phosphoric acid reacts with a solution of sodium hydroxide to produce a solution of
sodium phosphate and water as the only products
phosphoric acid H3PO4
reacts with is +
to produce is 
+1
-1
solution sodium hydroxide Na OH (aq)  NaOH(aq)
solution sodium phosphate Na+1PO4-3(aq)  Na3PO4(aq)
and is +
water H2O
H3PO4 + NaOH(aq)  Na3PO4 + H2O
(c) Chlorine gas is bubbled through a solution of lithium iodide, and the products are
found to be a solution of lithium chloride and solid iodine.
(d) Solutions of silver nitrate and barium chloride are mixed, and the products are a
precipitate of silver chloride and a solution of barium nitrate
6A - BALANCING EQUATIONS WORKSHEET
Balance the following equations. Show all of your work. If it is balanced
already, write “balanced” beside the right side of the equation.
Remember – coefficients always go in front of the entire formula.
1. ___Ag3PO4 + ___KBr  ___K3PO4 + ___AgBr
2. ___H2SO3 + ___O2  ___H2SO4
3. ZnS + O2  ZnO + SO2
4. Na2SO4 + Fe(NO3)3  NaNO3 + Fe2(SO4)3
5. AgNO3 + CaCl2  AgCl + Ca(NO3)2
6. Al2O3 + HCl  AlCl3 + H2O
7. C12H22O11 + O2  CO2 + H2O
8. Al2(SO4)3 + Ba(NO3)2  Al(NO3)3 + BaSO4
9. Na2O2 + H2O  NaOH + O2
10. CuO + H3PO4  Cu3(PO4)2 + H2O
11. Al2(SO3)3 + HCl  AlCl3 + H2O + SO2
12. BaCl2 + (NH4)2CO3  BaCO3 + NH4Cl
13. Ca3(PO4)2 + H2SO4  H3PO4 + CaSO4
14. KClO3  KCl + KClO4
15. CaCl2 + Fe2(SO4)3  CaSO4 + FeCl3
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HOW TO PREDICT THE PRODUCTS OF A CHEMICAL REACTION:
To be able to predict the products of a reaction, you must first be able to recognize what type of
reaction it is. There are five specific types of reactions that you need to be able to recognize:
(1) SYNTHESIS – means "putting together"; characterized by having two pure elements as its
reactants and there will be only ONE product formed which is a compound between these
two elements; LIKE A MARRIAGE; A + B → AB
Example 5: Write a balanced equation for the reaction that will take place when potassium
reacts with chlorine gas.
(b) Write a balanced equation for the reaction that will take place when aluminum is
allowed to react with oxygen gas.
(c) If iron (III) nitride is the only product formed in this reaction, write a balanced
equation for the reaction.
(2) DECOMPOSITION – means "breaking apart"; characterized by having only ONE reactant
which simply comes apart into its elements (the products); there must be some type of energy
which causes this reaction to happen, and this energy is usually written OVER the yield arrow;
LIKE A DIVORCE; CD → D + C
Example 6: Write a balanced equation for the decomposition of lead (II) oxide by heat.
(b) Write a balanced equation for the decomposition (by electricity) of water
(c) This is also a decomposition reaction since it has only ONE reactant, but you would
not be able to predict the products, so I will give them to you. Write a balanced
equation for the decomposition (by heat) of potassium chlorate into potassium chloride
and oxygen gas.
(3) COMBUSTION – literally means "burning", but we will take it to mean that a hydrocarbon
reacts with oxygen. The products are always the same – carbon dioxide and water vapor, no
matter which pure hydrocarbon is burned. REMEMBER TO ALWAYS WRITE OXYGEN AS
A REACTANT IN A COMBUSTION REACTION.
Example 9: Write a balanced equation for the combustion of propane.
(b) Write a balanced equation for the burning of octane.
(c) What is the balanced equation when methyl alcohol (CH3OH) is burned?
(d) Cyclobutane reacts with oxygen
(4) DOUBLE REPLACEMENT – characterized by having two compounds as reactants and
two different compounds will be the products; LIKE DO-SI-DO AND CHANGE PARTNERS.
Example 8: Write a balanced equation for solutions of barium chlorate and silver nitrate
being mixed.
(b) Write a balanced equation for the reaction of potassium cyanide solution plus tin (II)
fluoride crystals
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(c) Calcium phosphate and aluminum bromide are the PRODUCTS of this reaction
(d) Hydrochloric acid is mixed with a solution of calcium hydroxide
(5) SINGLE REPLACEMENT – characterized by having an element and a compound as
reactants and the products will be another element and another compound.
There are really two types of single replacement reactions – one in which the positive ion of the
compound is replaced by the lone element and one in which the negative ion of the compound is
replaced by the lone element; LIKE A LOVE TRIANGLE; A + BC → AC + B; or, in other
cases, XY + Z → XZ + Y
Example 7: Write a balanced equation for when magnesium is added to a solution of tin (IV)
nitrate.
(b) Lithium is added to a solution of barium hydroxide
(c) Magnesium chloride and hydrogen gas are the PRODUCTS of this reaction
(d) Liquid bromine is poured into a solution of sodium iodide.
(e) Fluorine gas bubbles through a solution of strontium bromide
We have been assuming that if you can write an equation, it will occur, and this is not
necessarily true. It is difficult to predict whether equations will actually take place or not, so we
will assume that ALL REACTIONS EXCEPT SINGLE REPLACEMENT REACTIONS WILL
TAKE PLACE as written. Whether or not a single replacement reaction will actually take place
will be based on the ACTIVITY SERIES of metals shown below. It is not necessary to
memorize it because it will be provided for you on all quizzes and exams, but you must know
how to use it AND YOU MUST REMEMBER TO USE IT!
The Activity Series for Non-metals will be confined to the halogens (Group VIIA) and can be
read directly off the Periodic Chart: Cl > Br > I >At
The last five metals will only replace each other if more active than the metal in the compound
(higher up on the chart).
The middle set of metals will: (1) replace each other if more active
(2) will replace the hydrogen from acids.
The first six metals will: (1) replace each other if more active
(2) will replace the hydrogen from acids
(3) will replace the hydrogen from water
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Activity Series of Metals
Li
K
Ba
Ca
Na
Mg
Al
Mn
Zn
Cr
Fe
Cd
Co
Ni
Sn
Pb
Sb
Cu
Hg
Ag
Pt
Au
IF A REACTION DOES NOT TAKE PLACE, write the reactants correctly, draw the yield
arrow, and write “no reaction.”
Example 10: Write a correct balanced equation for each of the following:
(a) A piece of zinc is added to a solution of silver nitrate.
(b) Tin(II) chloride solution is poured over magnesium
(c) Iron(III) nitrate solution plus copper
(d) Copper(II) chloride solution plus aluminum
(e) Sodium chloride solution plus calcium
(f) Silver is dropped into hydrochloric acid
(g) Magnesium is added to phosphoric acid
(h) Sodium plus water
(i) Calcium fluoride plus iodine
(j) Water is poured over gold
(k) Sulfuric acid plus lithium
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6B - PREDICTING PRODUCTS AND BALANCING EQUATIONS
ALWAYS check to see what you have on reactant side: element + element
(synthesis), 1 cmpd (decomposition), element + cmpd (single
replacement), cmpd + cmpd (double replacement), any hydrocarbon
(CxHy) being burned or combusted (combustion). Check the notes and
book if you have forgotten what the products should be predicted!
1. a piece of silver will react chemically with oxygen when it is heated
2. some pieces of aluminum foil are dropped into a solution of tin (IV)
sulfate
3. liquid octane (C8H18) is burned completely
4. chlorine gas is bubbled through a solution of sodium bromide
5. pure barium is added to a solution of silver acetate
6. solid copper (I) nitrite is dropped into a solution of magnesium
perchlorate
7. a solution of lead (IV) nitrate is reacted with calcium metal
8. crystalline bismuth (III) oxide decomposes with heat
9. methane is COMPLETELY combusted
10. solid iron reacts with oxygen gas to produce only one product and that
product is a solid iron compound in which the oxidation number of iron is
the highest possible
11. methyl alcohol (CH3OH) is COMPLETELY combusted.
12. crystalline sodium nitride is heated until it decomposes
13. lithium phosphate solution is mixed with cesium iodide solution
14. acetic acid neutralizes a solution of strontium hydroxide
15. bubbles of hydrogen gas and a solution of strontium chloride are
produced when metallic strontium is dropped into hydrochloric acid
16. a sample of mercury (I) oxide (solid) is heated until it decomposes
17. solid tetraphosphorus decoxide reacts with water to produce
phosphoric acid as its only product
18. bubbling chlorine gas through a solution of potassium iodide results in
two new products, one of which is a solid element and the other is a
solution of a compound
19. iodine crystals react with chlorine gas to form solid iodine trichloride
as its only product
20. a solution of sodium hydroxide is neutralized with phosphoric acid
To determine which of the physical states should be placed after every reactant and every
product, you must know the following solubility rules for ionic compounds in water.
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SOLUBILITY RULES FOR IONIC COMPOUNDS:
(1) All compounds whose cations are Group IA or ammonium are soluble, no matter what
the anion is.
(2) All nitrates, chlorates, perchlorates, and acetates are soluble, no matter what the cation
is.
(3) All chlorides, bromides, and iodides are soluble except silver, lead, and mercury(I).
(4) All sulfates are soluble except silver, lead, mercury(I), calcium, strontium, and
barium.
(5) Everything else will be INSOLUBLE unless it begins with a Groups IA or
ammonium cation.
Example 11: Predict whether the following ionic compounds are soluble in water:
(a) iron(III) nitrate
(j) ammonium sulfide
(b) sodium iodide
(k) cadmium fluoride
(c) ammonium phosphate
(l) tin(II) iodide
(d) mercury(I) chlorate
(m) cesium phosphite
(e) barium sulfate
(n) strontium cyanide
(f) zinc nitrite
(o) silver nitrite
(g) calcium sulfate
(p) silver bromide
(h) lead(II) perchlorate
(q) lead(IV) sulfite
(i) potassium bromide
(r) copper(I) chloride
HOW TO WRITE NET IONIC EQUATIONS:
1. Write a molecular equation for the reaction given. It is not necessary to balance it.
2. Break apart ions (ionize) of all strong acids, strong bases, and soluble salts that are present as
reactants unless the context indicates that they are present as a solid crystal.
(b) Ionize all strong, strong bases and soluble salts that form as products.
There are only 7 strong acids, and it is easier to just memorize them and know that
everything else is weak. These are HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4.
The strong bases are the hydroxides of Group I and II A except hydrogen, magnesium,
and beryllium. All other hydroxides (including ammonium hydroxide) are weak.
Soluble salts are broken into their ions. (The soluble ionic salts are determined from
memorizing the solubility rules.)
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3. Do NOT ionize molecular compounds such as gases, water, or organic molecules. Leave them
together as a molecule.
4. Remove all spectator ions (species which occur unchanged in any way from the reactant side
to the product side of the equation). [Unchanged means oxidation number as well as physical
state.] Sometimes nothing is removed; sometimes everything is removed and there is no
reaction.
Example 12: Write net ionic equations for the following reactions:
(a) Sulfuric acid is poured into a solution of calcium hydroxide
(b) A solution of acetic acid is poured into ammonium hydroxide
(c) Solutions of silver nitrate and hydrochloric acid are mixed
(d) Phosphoric acid solution is mixed with barium hydroxide solution
(e) Lead (II) acetate solution is stirred into sodium chlorate solution
(f) Hydrochloric acid is poured over sodium carbonate crystals
(g) Acetic acid neutralizes a solution of potassium hydroxide.
(h) Solutions of ammonium chloride and sodium hydroxide are mixed
(i) Solid potassium chlorate (when heated) decomposes into potassium chloride and
oxygen gas
(j) Sodium chloride solution is mixed with mercury (II) acetate solution
(k) Sulfurous acid (H2SO3) is poured into calcium chloride solution
6C - NET IONIC EQUATIONS
Finish writing the equation to the right of the arrow. Then, write the
complete ionic equation in the space below (leaving room on the blank to
write the next step). Mark a slash mark (/) through the spectator ions.
Then, write the net ionic reaction in the blank. Write NR in the blank if no
reaction occurs. Balancing the net ionic equations is not necessary.
1. NaCl + AgNO3 
2. BaCl2 + K2SO4 
3. KNO3 + LiCl 
4. Pb(NO3)2 + KI 
5. CaBr2 + Na2CO3 
6. Zn(ClO3)2 + (NH4)2S 
7. HCl + CsOH 
8. AlCl3 + KOH 
9. AgNO3 + NaC2H3O2 
10. CH3COOH + NaOH 
11. K2SO4 + Pb(ClO4)2 
12. NH4NO3 + Hg(ClO3)2 
13. BaCl2(s) + Na2SO4(aq) 
14. AlBr3(aq) + NaOH(aq) 
15. CuCl2(aq) + K2S(aq) 
16. AgClO4(aq) + Na2CO3(aq) 
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17. MgI2(s) + LiOH(aq) 
18. ZnBr2(aq) + Na2O(aq) 
19. CaCl2(aq) + K2CO3(aq) 
20. BaI2(aq) + Na3PO4(aq) 
21. Na2SO4 + KNO3 
22. HC2H3O2(aq) + Ca(OH)2(aq) 
6E – RECOGNIZING & WRITING TYPES OF REACTIONS
Complete and balance each of the following. Write the NET IONIC
equation (with formulas!) in the blank space under the words. In the blank
on the left, tell the type of reaction (D for decomposition, S for synthesis,
SR for single replacement, DR for double replacement, or C for
combustion).
1. Benzene (C6H6) + Oxygen
2. Calcium + Bromine
3. Potassium Hydroxide + Nitric Acid
4. Barium + Copper(II) Sulfate
5. Water
6. Magnesium + Nitrogen
7. Calcium Sulfate + Lithium Fluoride
8. Methane (CH4) + Oxygen
9. Calcium + Water
10. Sodium Hydroxide + Phosphoric Acid
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