Chapter 10 Chemical Reactivity
Now that we understand how molecules are made, let’s move on to the real fun part of
chemistry, how do the molecules react with each other
Use a different order? 10-4 Decomposition, 10-1 Combination, 10-2 Aqueous ions, 105 hydrates, 10-3 acids and bases?
10-1 Combination reactions
Key Concept:
A combination reaction occurs when two reactants combine to form a single
product.
Here are a couple of different combination reactions where we have two
reactants and one product:
4Al(s) + 3O2(g) 62Al2O3 (s)
4Fe(s) + 3O2(g) 6 2Fe2O3(s)
Both of these reactions are examples of combustion reactions, where one
substance (a fuel) is burned in the presence of some oxidizer (Often O2)
Anybody recognize the second reaction? Iron rusting?
We have to coat irons surface is some way to prevent iron from rusting and the
metal from flaking away as Fe2O3. Anybody know why we don’t do the same for
aluminum? The Fe2O3 flakes away from the Fe surface, exposing allowing the
Fe underneath to also rust. Al2O3 does not flake away, In fact I tightly adheres to
the Al underneath so the process stops!
In the above cases the product is a metal and a nonmetal, so it is an ionic
compound. What do you remember about the physical properties of ionic
compounds? Solids with high melting points
We can do a similar combination combustion reaction with nonmetals:
C(s) + O2(g) 6CO2(g)
S(s) + O2(g) 6SO2(g)
Here our products are covalent compounds. Note that these covalent
compounds are gases instead of solids. We have a lot more variability in the
physical states of covalent compounds they can be solids or liquids are gases at
RT.
So far all of our reactants have been elements. We can also have combination
reactions with molecules:
Na2O(s) + CO2(g) 6 Na2CO3(s)
NH3(g) + HCl(g) 6NH4Cl(s)
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10-2 Nomenclature of Polyatomic Ions
So our reactions have involved only solids, liquids or gases. Many reactions
involve chemicals dissolved in water or the aqueous (aq) state. We need to
discuss what the aqueous state really is before we look at any reaction that
occurs in water.
What happens when compounds dissolve in water
First of all, there is a big difference in what happens to ionic compounds and
covalent compounds when they dissolve in water. Remember the demonstration
I did back at the beginning of chapter 6, when I used light bulb to demonstrate
that Ionic compounds dissolved in water carry current, but covalent compounds
don’t? Lets explain that on a molecular level.
First question. Will all covalent compounds dissolve in water? If not, which one
do and which ones don’t?
(Polar do, non-polar don’t)
Why do polar compounds dissolve? Both water and polar covalents have
a dipole, and the dipole of one molecule interacts with the dipole of the
other , so the molecules spread out to maximize their interactions Figure
10.4. ( We will study this more in chapter 16)
What ions do is different. Remember that an ionic solid is held together by
alternating + and - charges. When exposed to water the water dipole interacts
with these charges. The negative ends of the dipoles interact with the + ions,
and the positive ends of the dipoles interact with the - ions as shown in figure
10.5
This explains why many ionic solids that are so strong can completely dissolve in
water at RT. And why the water can now carry a current.
This is actually a chemical reaction:
The Na+ (aq) represents the Na+ ion surrounded by water molecules. This
complex is called a solvated ion. Ditto for Cl-(aq)
You can also write the reaction as:
But I prefer the first way because it emphasizes the true chemical species in
solution.
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Again we can ask the same question we did with the covalent compounds. Does
this happen for all ionic compounds? And again the answer is no. Some ionic
compounds have such a strong attraction between the ions that they won’t break
apart in water and, instead, stay as a solid.
This is a very important point. The ionic compounds that dissolve to give you a
clear solution can only do this when water interacts so strongly with the ions that
it can literally tear the ionic solid apart into millions and millions of aqueous ions
interacting with the water molecules around them.
On the other hand there are lots of ionic compounds that have such a strong
ionic lattice that water cannot break the compound into individual ions. If an
ionic compound doesn’t turn into aqueous ions, then it remains as a solid and
does not dissolve.
Covalent ions
We have now talked about covalent compounds and ionic compounds, What
about covalent ions? When we were doing Lewis formula’s for covalent
compounds, I occasionally threw in some covalent ions like NH4+ or SO42-. What
happens to these molecules in water?
Key Concept:
When covalent ions interact with water ALL COVALENT BONDS REMAIN
INTACT. The covalent ions does NOT decompose into individual atomic ions
This is shown graphically in figure 10.6 with MgSO4
There are lots of these polyatomic ions in chemistry, and you need to know the
names and formula’s for many of them.
Table 10.1
Key Concept:
You must memorize the formula and name of the following polyatomic ions:
hydroxide, cyanide, acetate, nitrite, nitrate, chlorate, perchlorate, peroxide,
carbonate, sulfite, sulfate, oxalate, phosphate, ammonium and mercury(I)
The book goes into the naming system that is used to come up with some of
these names. As far as I am concerned, just memorize the names and don’t
worry about where they came from. (Unless this is something you want to know)
Practice Problems:
What is the correct name for NH4CH3COO? (Ammonium acetate)
What is the correct formula for sodium sulfite? (Na2SO3)
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Clicker questions:
Name and formulas of some salts made with complex ions
10-3 Bases and Acids
Bases
Key Concept:
One definition for a base is that it is a compound that yields hydroxide ions (OH-)
when dissolved in water
Many (but not all) metal oxides make bases when they undergo a combination
reaction with water, Thus:
Key Concept:
Metal oxides are generally basic when dissolve in water.
Here are some examples:
Na2O(s) + H2O(l) 62NaOH(aq)
BaO(s) + H2O(l) 6Ba(OH)2(aq)
-orNa2O(s) + H2O(l) 62Na+(aq) + 2 OH- (aq)
BaO(s) + H2O(l) 6Ba2+(aq) + 2 OH- (aq)
The above reaction works best with the reactive metals on the left hand of the
periodic table. The non-reactive metals to the right on the periodic table don’t do
so well.
Al2O3(s) + H2O(l)6 No reactions
Acids
Key Concept:
One definition of an acid is that it is a compound that produces hydrogen ions
(H+) when dissolved in water
If you think about it, what is the H+ ion? It is a single proton with no electrons
around it. A free proton like this cannot exist in water, so the H+ ion is actually a
shortcut method of writing a family of related water complexes like H3O+(aq),
H5O2+(aq) and H9O4+(aq). The hydronium ion, H3O+(aq) is the main species in
this mixture, so the Finicky chemist will use H3O+(aq), instead of H+(aq) to refer to
this acid species
Key Concept:
Many nonmetal oxides undergo a combination reaction with water to form acids.
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Examples include:
CO2(g) + H2O(l) 6H2CO3(aq)
SO3(g) + H2O(l) 6H2SO
There are many compounds with H in them that are not acids, things like H2(g) or
CH4(g) ore CH3CH2OH. How do you know when you have an acid, and when
you have a plain old covalent compound?
Key Concept:
Usually when your write the chemical formula of an acid, you write the acid H as
the first atom in the molecular formula. If there are other non-acidic hydrogens they are
given later in the formula.
Examples:
HCN, HCl, HC2H3O2 are hydrocyanic, hydrochloric and acetic acid respectively.
CH3CH2CH3 and CH3OH are not acids
Some compounds have more than one acidic proton. These compounds are
called polyprotic acids
Examples:
H2SO4 - Sulfuric Acid, a diprotic acid
H3PO4 - Phosphoric acids, a triprotic acid
Oxyacids Table 10.3
Any acid that contains an oxygen in the formula is called an oxyacid.
The names of the oxyacids are tied to the names of the ions that you obtain
when an acid deprotonates and goes into it’s anion form.
Key Concept:
If the anion name ends in -ite, the name of the corresponding acid is -ous acid
If the anion name ends in -ate, the name of the corresponding acid is -ic acid
Examples:
HNO3 (NO3- anion nitrate hence nitric acid)
HNO2 (NO2- anion, nitrite, hence nitrous acid)
Clicker questions:
H2SO3
H2SO4
Perchloric acid
Binary Acids Table 10.5
Binary acids consist of two atoms, one of which has to be the acidic hydrogen.
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Key Concept:
To name a binary acid you add the prefix hydro- to the anion and tack on the
suffix -ic acid to the end of the anion name.
Example:
HCl(aq) Hydrochloric acid
Clicker questions
HI
hydrosulfuric acid
One final note about th physical form of an acid. To be an acid you have to be in
a state that let’s you ionize. Thus HCl(g) or HNO3(g) are technically not acids. It
is only when they are dissolved in a solvent to become HCl(aq) or HNO3 (aq) that
these compound become true acids.
Some of the common properties of acids and bases are given in table 10.6
10-4 Decomposition Reactions
Back to chemical reactions
A decomposition reaction is just the opposite of a combination reaction
Key Concept:
In a decomposition reaction a single reactant breaks down into two or more
products.
Some examples of decomposition reactions are:
The airbag on your car uses the decomposition reaction
2NaN3(s) 6 2Na(s) + 3N2(g) to inflate your airbag
IF you want a little more force, and a little more ‘hair-trigger’ on the reaction, you
can use some much less stable compounds in the following reactions
Pb(N3)2(s) 6Pb(s) + 3 N2(g)
Hg(N3)2(s) 6Hg(l) + 3N2(g)
to set off a blasting cap!
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10-5 Hydrates
Many salts can combine with water to form interesting compounds called
hydrates.
Key Concept:
A hydrate is a salt that has incorporated a specific number of water molecules
into it’s solid crystal structure to make a new molecular structure.
For instance CuSO4(s) is a while powder
but it can react with 5 water molecules to form CuSO4@5H2O, a bright blue solid
Figure 10.14 No jpg file? Maybe do a demo?
Often the hydrated compound can be returned to the anhydrous (without water)
form by gentle heating.
Note how I wrote the ‘waters of hydration’ in the formula using a dot
convention. We do this because these waters are not that strongly bound to the
structure can be added or removed fairly easily .
10-6 Single replacement (Substitution) Reactions
Key Concept:
In a single replacement or substitution reaction an element in one compound is
replaced with a different element
Examples:
2Mg(l) + TiCl4(g) 6 2MgCl2(s) + Ti(s)
Fe(s) + H2SO4(aq) 6 FeSO4(aq) + H2(g)
We will see that many metals undergo a single replacement reaction with acids
to produce H2 gas.
10-7 Relative Activities of Metals
We can use the Single replacement reaction (previous section) to rank metals in
terms of relative reactivity
For instance Cu is considered more active than Ag because it can displace Ag in
the following reaction:
Cu(s) + 2AgNO3(aq) 6 Cu(NO3)2(aq) + 2Ag(s)
Figure 10.17 or demo
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By comparing one metal with another we can come up with the following ranking
of metals by activity.
Table 10.8
It would be nice if that activity table matched exactly with the periodic table, but it
doesn’t.
Key Concept:
The general trend however is that the more reactive metals are on the left and
the less reactive metals are on the right. Also in a single group the more reactive
metals tend to be lower on the periodic table.
If you are interested in metallurgy this activity ranking explains how smelting
works. See the text if you are interested.
10-8 Relative Activities of Halogens
We can try to rank the activities of nometals based on single displacement
reactions in a similar manner, but the reactivities of the nonmetals are too
diverse to make a nice table.
However, if we focus on just the halides, then you can get something to work.
For instance Br2 is considered to be more active than I2 because it will displace I2
in the reaction :
Br2(l) + 2NaI(aq) 6 2NaBr(aq) + I2(s)
(I won’t try to demonstrate this because both Br2 and I2 give off some pretty nasty
gases!)
The ranking you get is F2>Cl2> Br2> I2
Key Concept:
In general the nonmetals in a single group become more reactive as you go up
the periodic table
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10-9 Double-Replacement Reactions
Key Concept:
In a double-replacement reaction the cations and anions of two compounds are
exchanged
We will look at two major types of double-replacement reactions. In this section
we will examine precipitation reactions, and in the next section we will look at
acid-base reactions
Let’s try to react NaCl(aq) with LiNO3 (aq)
In a double-replacement reaction you exchange the + of on ionic compound with
the + of the other. This would give you the products LiCl(aq) and NaNO3(aq)
The reaction equation would be:
NaCl(aq) + LiNO3(aq) 6 LiCl (aq) + NaNO3 (aq)
Since all reactants and products ionize completely in water, a more appropriate
way to write this reaction would be:
Na+(aq) + Cl-(aq) + Li+(aq) + NO3-(aq) 6Li+(aq) + Cl-(aq) + Na+(aq) + NO3-(aq)
The clever chemist quickly realizes that NOTHING HAS HAPPENED! We have
all the same ions in the solution that we started with. Nothing has changed, so
this is actually a NO REACTION
So both
NaCl(aq) + LiNO3(aq) 6 N.R.
Na+(aq) + Cl-(aq) + Li+(aq) + NO3-(aq) 6 N.R.
Key Concept:
So simply switching cations around will not give you a reaction if both reactants
and products are soluble ionic compounds!
To have a chemical reaction some kind of chemistry must occur. The chemistry
that we will look at now is the chemistry of precipitation.
Back in section 10-2 I told you that some, but not all, ionic compounds dissociate
into individual ions in aqueous solution. Other ionic compounds have such a
strong bond between the anions and the cations that water cannot break this
interaction down, and the compound remains a solid.
What happens if we mix two solutions together, one that has the dissolved
aqueous cation and another that has a dissolved aqueous anion?
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What happens is that due to the random motions of the ions in the mixed
solutions a cation will get close enough to an anion that they can interact and
form the strong ionic bond and a tiny solid will start to form. As more and more
ions randomly bump into this tiny seed, the solid gets larger and larger, until is
get so large that the molecules drop out of solution and form a solid on the
bottom of the container. This process is called precipitation.
Key Concept:
Precipitation occurs when aqueous anions and cations coalesce for form a solid
that cannot remain in solution.
This is how we get our double-replacement reaction to actually do some
chemistry and produce a product. We start with two ionic compounds that are
completely dissolved into their respective aqueous ions. We then allow the
aqueous ions to mix together to form a new solid ionic precipitate.
Precipitation reactions
Our next task, then, is to identify which ions are soluble and which ones like to
form precipitates so we can identify when we have a precipitation reaction
It is not possible to predict the solubility (or insolubility) of all compounds, but
there are 6 simple rules you can use that will work about 95% of the time.
Key task:
Memorize and learn how to used Table 10.9 to predict the solubility of ionic
compounds.
Table 10.9
Note, as the table says, apply in order
Practice:
Predict the solubility of
(NH4)2SO4(s)
Soluble - rule 1
CaCO3
Insoluble - rule 5
Al2O3
Insoluble - Rule 5
Pb(NO3)2
Soluble - rule 2
Clicker question - Solubility of a couple of salts
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Now let’s integrate solubility with precipitation reactions
Practice problems:
Use the solubility to predict the products of the following reactions
BaCl2(aq) + Na2SO4(aq) 6
Products:
BaSO4 and NaCl
Solubility:
BaSO4 (s) exception to rule 6 ; NaCl (aq) rule 1
Final balanced reaction:
BaCl2(aq) + Na2SO4(aq) 6BaSO4 (s) + 2NaCl(aq)
LiOH(aq) + Pb(NO3)2(aq) 6
Products:
Pb(OH)2 and LiNO3
Solubility:
Pb(OH)2 (s) rule 5 and LiNO3(aq) rule 1
Final Balanced Reaction:
2LiOH(aq) + Pb(NO3)2(aq) 6 Pb(OH)2(s) +2 LiNO3 (aq)
NaCH3COO(aq) + CaBr2 (aq) 6
Products:
NaBr and Ca(CH3COO)2
Solubility:
NaBr (aq) rule 1; Ca(CH3COO)2 (aq) rule 2
N.R.
Clicker Question: one or two more solubility reactions
Writing Reaction Equations
So far we have been writing what are called Molecular equations because we
wrote all of our compounds as complete molecules. To show what is going on in
solution it is more appropriate to write complete ionic equations
Key concept:
In writing a complete ionic equation all soluble aqueous ionic compounds broken
down their component ions. However all covalent compounds, and all solids, liquids,
and gases remain unchanged.
Practice problems:
Let’s go back to the previous reactions and write the reactions as complete ionic
equations
BaCl2(aq) + Na2SO4(aq) 6BaSO4 (s) + 2NaCl(aq)
Ba2+ (aq) + 2Cl-(aq) + 2 Na+(aq) + SO42-(aq) 6BaSO4(s) + 2Na+(aq) + 2Cl-(aq)
Notice that the Cl2 in BaCl2 breaks into 2 independent Cl- ions
and the 2NaCl becomes 2Na+ and 2 Cl2LiOH(aq) + Pb(NO3)2(aq) 6 Pb(OH)2(s) +2 LiNO3 (aq)
2Li+ (aq) + 2OH-(aq) + Pb2+(aq) + 2NO3-(aq) 6Pb(OH)2(s) + 2Li+(aq) + 2 Cl-(aq)
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Clicker question
Write a single molecular equation and 3 out of 4 incorrect complete ionic
equations and have them pick out the correct equation
Notice in the above equations were pretty long and cumbersome to write and
included some ions that were unchanged throughout the reaction. We call these
ions spectator ions because all they do is watch, they don’t do any real
chemistry.
There is a third way to write a chemical equation that drops these spectator ions
out of the written equation because they don’t take part in in chemistry. We call
this third type of equation a net ionic equation
Key Concept:
A net ionic equation is written just like a complete ionic equation, however all
spectator ions are removed from both sides of the equation
Practice:
Continuing on with our examples
Molecular: BaCl2(aq) + Na2SO4(aq) 6BaSO4 (s) + 2NaCl(aq)
Complete Ionic:
Ba2+ (aq) + 2Cl-(aq) + 2 Na+(aq) + SO42-(aq) 6BaSO4(s) + 2Na+(aq) + 2Cl-(aq)
Net Ionic: Ba2+(aq) + SO42-(aq) 6BaSO4(s)
Molecular: 2LiOH(aq) + Pb(NO3)2(aq) 6 Pb(OH)2(s) +2 LiNO3 (aq)
Complete Ionic:
2Li+ (aq) + 2OH-(aq) + Pb2+(aq) 2NO3-(aq) 6Pb(OH)2(s) + 2Li+(aq) + 2 Cl-(aq)
Net Ionic: 2OH-(aq) + Pb2+(aq) 6Pb(OH)2(s)
Notice how much shorter the net ionic equations are to write
10-10 Acid-Base Reactions
Our second example of a double replacement reaction is the acid-base reaction.
HCl(aq) + NaOH(aq) 6 H2O(l) + NaCl (aq)
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) 6 H2O(l) + Na+(aq) + Cl-(aq)
H+ (aq) + OH-(aq) 6 H2O(l)
Here the aqueous ions H+ and OH- combine to form the covalent water molecule
In this reaction a strong, corrosive acid is combined with a strong, corrosive,
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base to make neutral water with a little dissolved salt. We call this a
neutralizing reaction, and, when the amount of acid is exactly equal to the
amount of base we can say that the acid and base have neutralized each other
because the acid and basic characters of the two solutions have been nullified.
You have heard the term ‘salt’.
Key Concept:
On definition of a salt is that it is the ionic compound that results (along with
water) from the reaction of an acid and a base.
Practice problems:
Predict the product and write the net ionic equation for the following reactions:
KOH(aq) + H2SO4(aq)
2KOH(aq) + H2SO4(aq) 62H2O(l) + K2SO4(aq)
2OH-(aq) + 2H+62H2O(l)
HC2H3O2(aq) + Ba(OH)2(aq)
HC2H3O2(aq) + Ba(OH)2(aq) 6 2H2O(l) + Ba(C2H3O2)(aq)
Note Ba(OH)2 slightly soluble rule 5. Ba(C2H3O2)(aq) soluble rule 2
Clicker Questions
Another acid base reaction
10-11 Oxidation-Reduction Reactions
Key Concept:
Oxidation-Reduction reactions are reactions in which electrons are transferred
from one reactant to the other
Many of our reaction between metals and nonmetals are oxidation reduction
reactions. For instance:
2Na(s) + S(s) 6Na2S(s)
Can you see the electrons? I can’t! But if you go to the electron configurations:
2Na• + •:S:•
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2Na+
+ ::S::2- (Solid broken into ions for clarity)
[Ne]3s1 [Ne]3s23p4
[Ne]
[Ne]3s23p6 =[Ar]
Now you can see eletrons leaving one atom to go to another.
Key Concepts:
1.) When an atom gives up its electron, it gains + charge, and we call this
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oxidation
2.) When an atom gains electrons it loses + charge and we call this reduction
3.) Electrons, like all mass, must be conserved, so in any oxidation reduction
chemical reaction the number of electrons lost by some of the reactant atoms must
equal the number of electrons gained by other reactant atoms
Many of the reactions we have looked at so far are oxidation-reduction reactions
One final bit of somewhat funky and confusing nomenclature:
In the above reaction:
2Na(s) + S(s) 6Na2S(s)
Na was being oxidized because it lost electrons so its charge increased
S was being reduced because it was gaining electrons so its charge decreases
Na is also called a reducing agent because it allowed the reduction of S to occur
S is called an oxidizing agent because the oxidation of Na could not occur
without it
Key Concepts:
In an oxidation-reduction reaction:
The material that loses electrons, gain + charge, is being oxidized, and is
considered the reducing agent in the reaction
The material that gains electron....loses + charge....is being reduced....is
considered to be the oxidizing reagent in the reaction.
Practice problems:
First identify the atom or molecules being oxidized and reduced in the following
reactions, then identify the oxidizing and reducing agents
2Na(s) + Cl2(g) 62NaCl(s)
Na is oxidized and is the reducing agent
Cl is being reduced and is the oxidizing agent
Clicker question:
First identify the atom or molecules being oxidized and reduced in the following
reactions, then identify the oxidizing and reducing agents
Zn(s) + HgCl2(aq) 6ZnCl2(aq) + Hg(l)
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Two finals notes before we leave this chapter.
1.Classifying reactions as decomposition, single replacement and double
replacement is based on simply looking at the numbers of reactant molecules vs
the number of product molecules. This is good for bookkeeping, but does not
really tell you much about the underlying chemistry. Later in the chapter we
started talking about precipitation reactions, acid-base reactions and oxidation
reduction reactions. This viewpoint clearly focuses on the underlying chemistry
of the reaction, and I find to be the better way to classify reactions.
2. This chapter very briefly introduced oxidation-reduction reactions. Your text
doesn’t return to this important class of reactions until chapter 24 which we
wouldn’t get to until the end of next semester (if we were lucky). Since I don’t
want you to miss out on this important type of reaction, we will now skip ahead to
chapter 24!