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Introduction to Chemical Reactions

In any chemical change mass is conserved.

The Law of Conservation of Matter says that matter can neither be created nor destroyed, but

can be changed in form. This is one of the most fundamental laws of chemistry and is credited

to Antoine Lavoisier who is often called the father of modern chemistry. In other words, the

total mass of the material(s) before the reaction is the same as the total mass of material(s) after

the reaction. To us that mean that only the arrangement of particles changes during a chemical

reaction. Generally, this fact has been confirmed by countless experiments. However, there is

another way to understand this concept. You can think in terms of the atoms themselves.

The Combustion of Methane

CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (g)

1 moleculeof methane

2 molecules of oxygen 1 molecule of

carbon dioxide2 molecules of water

For example: Take the synthesis reaction between hydrogen gas and oxygen gas. The product

of this reaction is water. The unbalanced chemical equation for this reaction is

Hydrogen combines with oxygen to form water

H2 + O2 H2O reactants products

But there is a problem here. If we count the element symbols, taking subscripts into account,

the numbers of each type of atoms are not the same on opposite sides of the arrows. We say

this reaction is "not balanced" because it does not show conservation of mass. More

importantly, it does not show the ratio in which hydrogen and oxygen actually react.

To correct the situation we need to balance the reaction. In order to do this, we place numbers

(coefficients) in front of the symbols. For example:

2 H2 + O2 2H2O

The Rules: To balance a reaction you must:

• know the correct formulas for the reactants and products

• use coefficients to make the same number of each type of atom on each side

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Reactions in Solutions

Most of the reactions we will see during our course fall into three broad categories:

§ precipitation reactions

§ reactions of acids

§ reduction-oxidation reactions ("redox")

1- Precipitation Reactions A precipitation reaction is a reaction in which soluble ions in separate solutions are mixed

together to form an insoluble compound that settles out of solution as a solid. That insoluble

compound is called a precipitate.

Solubility Rules for Ionic Compounds

• Solubility rules are useful summaries of information about which ionic compounds (or

combinations of ions) are soluble in water and which are not.

• They are also important tools for making predictions about whether certain ions will

react with one another to form a precipitate.

• In addition, they are useful for figuring out what ions might be involved when a

precipitation reaction has been observed (net ionic equation).

In this section you will use solubility rules to predict precipitation reactions and then write

equations to represent them.

"Solubility Rules":

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To write chemical equations for the reactions observed, we need to recall what is happening

during precipitation: ions exchange partners. At least ONE of the products is insoluble if a

precipitate forms.

These molecular reactions give the basic information about the processes such as the ratios in

which the substances react, but they don't represent accurately what is happening. Since all of

these compounds are electrolytes in solution, they actually exist as separated ions, not ion pairs

like NaCl. And if only one of the "products" is insoluble, the other actually never forms.

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Net Ionic Reaction To try to express better what is happening in the solution two other types of reactions may be

written. The first of these may be thought of as transitional. It is called the total ionic reaction

and shows everything in the mixture as it exists. For the first reaction we wrote it would look

like this:

Ag+ + NO3- + K+ + Cl- AgCl + K+ + NO3

- The transformation of this to the net ionic reaction is what we are really interested in. This

third form of a balanced chemical equation shows only the substances that actually do

something in the reaction--and in their correct forms. To get it, we take the total ionic reaction

and eliminate those things which do not change from one side of the arrow to the other. The

species NO3- and K+ are "spectators" and their elimination yields the net ionic reaction:

Ag+ + Cl- AgCl The "spectators", of course, are the soluble compounds, so knowing your solubility rules makes

the transition from molecular to net ionic much easier.

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Summary of Writing Balanced Ionic Equations The first step in writing a balanced equation is predicting the products of the reaction as

discussed above. Then the steps below are completed in sequence:

• Balance the Molecular Equation: In the “molecular” equation, nothing is broken up into

ions. Salt formulas are written so that the cation charges exactly balance out the anion

charges so that the salt is neutral. Then the equation is balanced for atoms.

• Balance the Total Ionic Equation: The first step in writing an ionic equation is to decide

what species should be broken up into ions. The rules below should help!

Break up into Ions Do NOT break up! Leave “as is”!

• Strong Acids. HCl, HBr, HI, HNO3, HClO4,

and H2SO4 are the most common examples;

assume other acids are weak.

• Strong Bases. NaOH, KOH, or Ba(OH)2 are the

most common examples; assume other bases are

weak.

• Soluble Salts. Salts of the alkali metals, salts

containing the NH4+ ion, the NO3

- ion, and some

other salts

• Weak Acids. Nearly all acids are

weak.

• Weak Bases. Nearly all bases are

weak.

• Insoluble Salts. Most salts are

insoluble.

• Non-electrolytes or Weak

Electrolytes. Examples include H2O,

gases, pure elements, hydrocarbons,

and alcohols.

• Balance the Net Ionic Equation: Identify all spectator ions: these are ions that are

identical on both sides of the balanced total ionic equation. Remove the spectator ions from

the equation. What remains is the net ionic equation. Finally, simplify the stoichiometric

coefficients if all of them are divisible by a common factor.

If all the ions are spectator ions so that nothing is left for your net ionic equation, no reaction

has taken place!

____________________________________________________________________________

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Exercises

For each of the following reactions, complete the chart. Be sure to balance all of your

equations.1. Mg(OH)2(s) + HCl(aq)

(a) Reaction type:

Formulas of Products Formed:

(b) Molecular Equation

(c) Total Ionic Equation

(d) Net Ionic Equation

Answer to 1.(d): Mg(OH)2(s) + 2 H+(aq) Mg2+(aq) + 2 H2O(l)

____________________________________________________________________________

2. AgNO3(aq) + K2Cr2O7(aq)

(a) Reaction type:

Formulas of Products Formed:

(b) Molecular Equation

(c) Total Ionic Equation

(d) Net Ionic Equation

Answer to 2.(d): 2 Ag+(aq) + Cr2O72-(aq) Ag2Cr2O7(s)

____________________________________________________________________________

3. NH3(aq) + HC2H3O2(aq)

(or CH3COOH)

(a) Reaction type:

Formulas of Products Formed:

(b) Molecular Equation

(c) Total Ionic Equation

(d) Net Ionic Equation

Answer to 3.(d): NH3 + HC2H3O2 NH4+(aq) + C2H3O2

-(aq)

____________________________________________________________________________

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4. NaOH(aq) + H2SO4(aq)

(a) Reaction type:

Formulas of Products Formed:

(b) Molecular Equation

(c) Total Ionic Equation

(d) Net Ionic Equation

Answer to 4.(d): OH-(aq) + H+(aq) H2O(l) (obtain this after all coefficients

have been divided by 2)

____________________________________________________________________________

5. H2S(aq) + Ba(OH)2(aq)

(a) Reaction type:

Formulas of Products Formed:

(b) Molecular Equation

(c) Total Ionic Equation

(d) Net Ionic Equation

Answer to 5.(d): H2S(aq) + Ba2+(aq) + 2 OH-(aq) BaS(s) + 2 H2O(l)

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II- Reaction of Acids

Unlike precipitation reactions, those involving acids may or may not have characteristic visual

cues. Of the four types of acid reactions, only two had distinctive physical changes:

• acids react with metal carbonates to release carbon dioxide gas (bubbles)

• acids react with many metals to release hydrogen gas (more bubbles)

Of these two, the second is actually a redox reaction. So what is going on in many acid

reactions if we can't see much?

All acids are electrolytes. You can generally tell a substance is an acid from its formula. Most

acid formulas begin with H.

So in acid solutions ions are present and electricity will flow through the mixtures. Perhaps the

simplest kind of acid reaction is with a base. Bases are also electrolytes and most common

bases are metal hydroxides like NaOH or KOH. Ammonia (NH3) is the only common base

which is not a hydroxide. When acids react with bases, the same kind of "ion switch" happens

as with a precipitation reaction. But instead of two ionic compounds forming, water is one of

the products.

acid + base salt + water

HCl + NaOH NaCl + H2O

Acids can also react with metal oxides but the visual evidence generally looks like simple

dissolving. The chemical process that occurs, however, is very much like the reaction of an

acid with a base. In fact, a metal oxide can be thought of as a "base without water".

acid + metal oxide salt + water

H2SO4 + MgO MgSO4 + H2O

We also saw that acids react with metal cabonates to release carbon dioxide gas. A metal

carbonate is just a metal oxide which has picked up carbon dioxide, and so:

acid + metal carbonate salt + water + carbon dioxide

2HCl + K2CO3 2KCl + H2O + CO2

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The fourth type of reaction does not really belong with this group. Acid reactions involve the

transfer of H+. But the reaction between an acid and a metal to produce hydrogen gas involves

the transfer of electrons. This indicates it is a redox process. The metal "displaces" the H from

the acid and takes its place, forming a salt and releasing hydrogen gas:

acid + metal salt + hydrogen

2 HCl + Mg MgCl2 + H2

One of the things we have learned in the lab is that there are dyes which give different colors

when exposed to acids, bases and "neutral" solutions. The colors which we saw are shown

below for the indicators used.

In general indicators are chosen for particular situations because they give visual cues to what

otherwise would be an unseen process. Some color changes are more easily observed than

others.

Molecular reactions for the two acid-base combinations observed in the lab can now be written:

Balanced equations such as these have additional utility.

Net ionic versions of these neutralization reactions simply show which substances are

electrolytes and which substances actually change during the reaction.

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Molecular and net ionic equations for the reaction of calcium oxide and calcium carbonate with

acid can also now be written, keeping in mind that soluble ionic compounds will be electrolytes

and insoluble ones are essentially non-electrolytes.

While most acids react with many metals to release hydrogen gas, nitric acid is exceptional for

its ability to react with some metals that other acids cannot attack. But the products of such a

reaction are very different.

For now, we'll keep to the simple displacement of hydrogen gas and write molecular and net

ionic reactions for that.

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III- Redox Reaction

In the reactions with acids that we have looked at in the lab the constant characteristic is the

transfer of H+ from the acid to some other substance (base, oxide, etc.). Reactions in which

electrons are transferred are referred to as reduction-oxidation reactions or simply "redox".

Visually these reactions vary considerably. Color changes, precipitates, gas formation,

temperature changes--all are possible singly and in combination.

The substance which loses electrons is said to be oxidized. The substance which gains those

electrons is said to be reduced.

How do you know when a reaction is redox? If it does not fit the models we have established

for precipitation and acid reactions, it is likely to be redox. The only sure way to determine

whether a reaction is redox or not is to check for a transfer of electrons.

Oxidation numbers are useful for exactly this kind of task.

Redox Term:

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Oxidation number:

• In the very simple process of magnesium combining with oxygen we can see the change

in oxidation number happening:

2 Mg + O2

2 MgO+ 2 -20 0

When the oxidation numbers change, the reaction is redox.

• Contrast that to a precipitation reaction:

MgCl2 + 2 NaOH Mg(OH)2 + 2 NaCl+2 -1 +1 -2 +1 +2 -2 +1 +1 -1

All of the oxidation numbers remain the same on both sides, so this is not a redox

process.

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The rules for oxidation state are in some ways arbitrary and unnatural, but here they are:

1. Any free (unattached) element with no charge has the oxidation state of zero. Diatomic gases such as O2 and H2 are also in this category.

2. All compounds have a net oxidation state of zero. The oxidation state of all of the atoms add up to zero.

3. Any ion has the oxidation state that is the charge of that ion. Polyatomic ions (radicals) have an oxidation state for the whole ion that is the charge on that ion. The ions of elements in Group I, II, and VII (halogens) and some other elements only have one likely oxidation state.

4. Oxygen in compound has an oxidation state of minus two, except for oxygen as peroxide, which is minus one.

5. Hydrogen in compound has an oxidation state of plus one, except for hydrogen as hydride, which is minus one.

6. In radicals or small covalent molecules, the element with the greatest electronegativity has its natural ion charge as its oxidation state.

KNOW THIS IS IT A REDOX REACTION?

A redox reaction will have at least one type of atom releasing electrons and another type of

atom accepting electrons. How can you most easily tell if a reaction is redox? Label every atom

on both the reactant and product side of the equation with its oxidation number. If there is a

change in oxidation number from one side of the equation to the other of the same species

of atom, it is a redox reaction. Each complete equation must have at least one atom species

losing electrons and at least one atom species gaining electrons. The loss and gain of electrons

will be reflected in the changes of oxidation number.

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Let’s take the following equation:

K2(Cr2O7) + KOH 2 K2(CrO4) + H2O

Potassium potassium potassium water

Dichromate hydroxide chromate

Is it a redox equation or not?

• All of the oxygens in compound have an oxidation state of -2.

• All of the hydrogens have an oxidation state of +1.

• Potassium is a group I element, so it should have an oxidation state of +1 in the

compounds. That seems to make sense because dichromate and chromate ions have a

charge of -2 and there are two potassium atoms in each compound.

• Hydroxide ion has a charge of -1 and it has one potassium.

• But what about the chromium atoms? We can do a little primitive math on the material

either from the starting point of the compound or the ion to find the oxidation state of

chromium in that compound.

• The entire compound must have a net oxidation state of zero, so the oxidation numbers

of two potassiums one chromium and four oxygens must equal to zero.

2 K + Cr + 4 O = 0

• We know the oxidation state of everything else but the chromium.

2(+1) + Cr + 4 (-2) = 0 and Cr = +6.

Or we could do it from the point of view of the chromate ion.

Cr + 4 O = -2

The oxygens are minus two each. Cr + 4 (-2) = -2

Either way Cr = +6.

Now the dichromate; 2 K + 2 Cr + 7 O = 0 and

2 (+1) + 2 Cr + 7 (-2) = 0.

Then 2 Cr = +12 and Cr = +6.

You can do the math for the dichromate ion to see for yourself that the chromium does not

change from one side of this equation to the other. As suspicious-appearing as the equation

might have seemed to you, it is not a redox reaction.

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MATERIAL OXIDATION STATES MATERIAL OXIDATION STATES 1. NaCl Na = +1 , Cl = -1 17. P4O6 O = -2, P = +3

2. KMnO4 K = +1, O = -2, Mn = +7 18. H3PO4 H = +1, O = -2, P = +5

3. diamond C = 0 19. Mg3N2 Mg = +2, N = -3

4. CO2 C = +4, O = -2 20. MgH2 Mg = +2, H = -1 (hydride)

5. CO C = +2, O = -2 21. NH3 H = +1, N = -3

6. KCN K = +1, C = +2, N=-3 22. N2H4 H = +1, N = -2

8. Fe2O3 O = -2, Fe = +3 23. (NH4)+ H = +1, N = -3

9. Fe3O4 O = -2, Fe = +8/3 24. N2 N = 0

10. (ClO4)- O = -2, Cl = +7 25. (NO3)- O = -2, N = +5

11. (ClO3)- O = -2, Cl = +5 26. (NO2)- O = -2, N = +3

12. (ClO2)- O = -2, Cl = +3 27. NO2 O = -2, N = +4

13. (ClO)- O = -2, Cl = +1 28. NO O = -2, N = +2

14. Cl- Cl = -1 29. N2O O = -2, N = +1

16. P2O5 O = -2, P = +5 30. Na2O2 Na = +1, O = -1 (peroxide!)

Problem:

What are the oxidation states of the transition metal in each of the following compounds?

a) KMnO4 b) Na2CrO4 c) CrO3 d) MnO2 e) Na2Fe2O4 f) Mn(CO)3

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ANSWERS TO REDOX EQUATIONS

31. 0 0 +1 -2 2( H0 H+1 +e-) Oxidation H2 + O2 H2O 2e- + O0 O-2 Reduction Balanced equation 2 H2 + O2 2 H2O

32. +2 -2 0 0 2e- + Hg+2 Hg0 Reduction HgO O2 + Hg O-2 O0 + 2e- Oxidation Balanced equation 2 HgO O2 + 2 Hg

33. +1 +5-2 +1 -1 0 6e- + Cl+5 Cl-1 Reduction KClO3 KCl + O2 3( O-2 O0 + 2e-) Oxidation Balanced equation 2 KClO3 2 KCl + 3 O2

34. +1 -1 0 +2 -1 0 2( e- + H+1 H0) Reduction H Cl + Zn ZnCl2 + H2 Zn0 Zn+2 + e- Oxidation Balanced equation 2 HCl + Zn ZnCl2 + H2

35. +1 +5 - 2 0 +2 +5 -2 0 2( e- + Ag+1 Ag0 Reduction Ag NO3 + Cu Cu(NO3)2 + Ag Cu0 Cu+2 + 2 e- Oxidation Balanced chemical equation 2 AgNO3 + Cu Cu(NO3)2 + 2 Ag 36. +8/3 -2 +2 -2 0 +4 -2 4 ( C+2 C+4 + 2 e-) Oxidation Fe3O4 + CO Fe + CO2 3 ( 8/3 e- + Fe+8/3 Fe0) Reduction Balanced chemical reaction 3 Fe3O4 + 4 CO 3 Fe + 4 CO2

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37. 0 +4 -2 +1 +6 -2 +2 +6 -2 +1 -2 Pb0 Pb+2 + 2e- Oxidation Pb + PbO2 + H2SO4 PbSO4 + H2O 2e- + Pb+4 Pb+2 Reduction Balanced equation Pb + PbO2 + 2 H2SO4 2 PbSO4 + 2 H2O - lead oxidizes and reduces.

38. -4 +1 0 +1 -2 +4 -2 C-4 C+4 +8e- Oxidation CH4 + O2 H2O + CO2 4( 2e- + O0 O-2 ) Reduction Balanced equation CH4 + 2 O2 2 H2O + CO2

39. -9/4 +1 0 +4 -2 +1 -2 25 ( 2e- + O0 O-2) Reduction C8H18 + O2 CO2 + H2O 2 ( C-9/4 C+4 + 25/4 e-) Oxidation Balanced equation 2 C8H18 + 25 O2 16 CO2 + 18 H2O

40. 0 +1 +5 -2 +2 -2 +2 +5 -2 +1 -2 2( 3e- + N+5 N+2) Reduction Cu + H(NO3) NO + Cu(NO3)2 + H2O 3( Cu0 Cu+2 + 2e-) Oxidation Balanced equation 3 Cu + 8 HNO3 4 H2O + 2 NO + 3 Cu(NO3)2

41. +1 +6 -2 +1 -1 0 +3 -1 +1 -2 +1 -1 3( Cl-1 Cl0 + 1e-) Oxidation K2(Cr2O7) + HCl Cl2 + CrCl3 + H2O + KCl 3 e- + Cr+6 Cr+3 Reduction Balanced equation K2(Cr2O7) + 14 HCl 3 Cl2 + 7 H2O + 2 CrCl3 + 2 KCl

42. +1 +7 -2 +1 +2 -3 +1 -2 +4 -2 +1 -2 +1 +1 -2 +4 -3 KMnO4 + K(CN) + H2O MnO2 + K(OH) + K(OCN) 3( C+2 C+4 +2e-) Oxidation 2(3e- + Mn+7 Mn+4) Reduction Balanced equation 2 KMnO4 + 3 KCN + H2O 2 MnO2 + 2 KOH + 3 K(OCN)

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43. +1 +7 -2 +1 +3 -2 +1 +6 -2 +2 +6 -2 +4-2 +1 -2 +1 +6-2

KMnO4 + H2C2O2 + H2SO4 MnSO4 + CO2 + H2O + K2SO4 5e- + Mn+7 Mn+2 Reduction 5( C+3 C+4 + e-) Oxidation Balanced 2 KMnO4 + 5 H2C2O4 + 3 H2SO4 2 MnSO4 + 10 CO2 + 8 H2O + K2SO4

Because there are so many different scenarios for electrons transfer, we focused on only a few

• metal displacement

• hydrogen displacement

• halogen displacement

• non-trivial redox

Actually that last category is made-up. It signifies reactions which are typically pretty complex

and which cannot be balanced by simple inspection as we do with all other types of reactions.

The first three categories are all single displacement reactions and they all follow a similar

"script". Each involves an element which pushes out another element in a compound and takes

its place. The elements which exchange places are similar in some way. For example, metals

may displace metals, but not non-metals. Substances which are positive when ions may

displace positive ions, but not negative ions. Etc.

Your data from the lab suggests that elements may in fact be ranked according to their ability to

displace others. Such rankings are known as activity series and the most common one includes

metals and hydrogen.

In the lab, metal displacement is seen as a darkening of the spot of solution placed on the metal

strip. Hydrogen displacement is observed as bubbling or fizzing since hydrogen is a gas. What

we want to do with the observations is rank, from most able to displace to least able to displace,

the metals and hydrogen.

Activity series:

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The written reactions all look alike. For example, copper is able to displace silver from

solution:

Although the halogens are not generally included in the common activity series, they form a

series of their own when compared with each other. The observations you made in the lab can

be confusing. You have to remember what color each halogen is in hexane and which halogen

you originally placed into each mixture. If the color of the hexane is what you would expect for

the halogen you added, then no reaction occurred.

Some reactions:

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As with the metal displacements, all of the halogen displacement reactions observed are written

in the same way: