Download Class Notes

Class Notes
Unit 5 – Chemical Reactions (Ch. 11)
What do we already know about chemical reactions?
A chemical reaction is another way to talk about a chemical change. There are four
general indicators that a chemical change has taken place: (1) the formation of a solid
precipitate upon the mixing of 2 solutions; (2) the production of gas bubbles (not due to a
physical change such as boiling); (3) a color change (not due to mixing or dilution); and
(4) a temperature change (not due to external heating or cooling). If one of these four
indicators is not present it doesn’t mean there wasn’t a chemical change. They are just
general clues to look for.
Dalton said that a chemical reaction occurs when atoms are separated, joined, or
rearranged and that atoms of one element aren’t ever changed into atoms of another
element.
The components that you begin with in a reaction are the reactants. The components
formed through the chemical change are the products. In a chemical reaction, one or more
reactants change into one or more products. There is always an energetic driving force
that causes the bonds of reactants to break and create new bonds (to form the products).
Every chemical change involves a change in energy as well.
A chemical equation is a quick, shorthand notation used to convey as much as possible
about what happens in a chemical reaction. Reactants are shown on the left side of the
equation and products are shown on the right side. The reactants and products are
separated by an →. This arrow is can be interpreted as “yields”, “produces”, or other
words describing the type of chemical change.
Reactants → Products
If there are more than one reactants or products, they are separated by a + sign. For
example, in the reaction between iron and oxygen to produce rust (iron (III) oxide) the
chemical equation is:
Fe (s) + O2 (g) → Fe2O3 (s)
You can also express the state of matter of the reactants and products in a chemical
equation, as seen above. There are 4 symbols to show the state of matter of the
substances: (s) for solids, (l) for liquids, (g) for gases, and (aq) for aqueous solutions. An
aqueous solution involves a substance that is usually a solid at room temperature (like
ionic compounds) dissolved in water. These are shown as subscripts after the compound
and are another way to express in shorthand as much as possible about the reaction.
Balancing Equations
The law of conservation of mass (aka: the law of conservation of matter) says that during
a chemical reaction or a physical change mass is conserved; mass is neither created nor
destroyed. This implies that the atoms that were there in the reactants (before the
chemical change) must be there in the products (after the chemical change) just
rearranges somehow. The subscripts in chemical formulas tell the number of atoms of
each element involved in the compound.
Looking at the chemical equation for the formation of rust (Fe (s) + O2 (g) → Fe2O3 (s)) is
the law of conservation of mass obeyed? No, there is 1 atom of iron and 2 atoms of
oxygen that react to produce 2 atoms of iron and 3 atoms of oxygen. This chemical
equation implies that matter was created. This is impossible.
The chemical equation shown in this manner is known as a “skeletal equation” or an
“unbalanced equation.” It shows only the formulas of the compounds involved in the
reaction but it does not show the ratio of each compound (how many of each is reacted or
produced). To determine the real amounts of the components involved in the reaction,
you must “balance the equation.” The number of each component is shown through its
coefficient (the number in front of the compound). In balancing an equation, coefficients
are put in front of the chemical formulas to make the number and types of atoms in the
reactants equal to the number and types of atoms in the products. You can never change
the subscripts in the compound’s formula; you can only change the coefficient. Also, the
coefficient affects the entire compound (it is distributed/multiplied to each element in the
compound).
One method to balancing equations is to first list the elements involved for both the
reactants and the products (the elements will always be the same). It is good to list them
in the same order for each list so not to confuse yourself later. Then, write the initial
number of atoms of each type. Next, add coefficients to the compounds to try and make
the number and types of atoms in the reactants equal the number and types of atoms in
the products. This is an iterative process and coefficients can be changed multiple times if
need be before the equation is balanced.
A few general tips for choosing which element to balance when:
1) Balance H and O last. These are very common elements and often in most
compounds in the reaction. Start with less common elements, like those that
appear in only one of the reactants or products.
2) If there is an even number of an element on one side of the reaction and an odd
number on the other side, the odd number must be made even to balance the
reaction. Since only whole number coefficients are used, the even will never be
able to be made odd. Locate the odd element and multiply by a coefficient to
make it even.
3) Leave the elements in simple compounds (like O2, I2, etc.) until the end to balance.
It will be easy to add coefficients to these to make that element be whatever you
need it to be without affecting other elements.
Ex: Fe (s) + O2 (g) → Fe2O3 (s)
Fe 1
Fe 2
O 2
O 3
Fe (s) + O2 (g) → 2Fe2O3 (s)
Fe 1
Fe 2 | 4
O 2
O 3|6
4Fe (s) + O2 (g) → 2Fe2O3 (s)
Fe 1 | 4
Fe 2 | 4
O 2
O 3|6
4Fe (s) + 3O2 (g) → 2Fe2O3 (s)
Fe 1 | 4
Fe 2 | 4
O 2|6
O 3|6
Ex: Methane (CH4) gas and oxygen react to produce carbon dioxide and water.
CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l)
Catalysts
Ex: A solution of 3% hydrogen peroxide (H2O2) decomposes to produce water and
oxygen gas.
H2O2 (aq) → H2O (l) + O2 (g)
Does hydrogen peroxide decompose on its own? Do you have to be careful of your bottle
of hydrogen peroxide at home turning into water and oxygen gas? No, not very quickly at
least. There are a lot of factors that can affect whether a reaction will occur spontaneously
or not. These include: bond strength, temperature, motion of compounds, and
concentration of the reactants. Some reactions require some sort of a push to get them
started; something that lowers the required energy necessary to start a reaction. The
minimum energy required to cause the reaction to proceed is known as the “activation
energy.” The substances that lower the activation energy so that the reaction can proceed
at a faster rate are called catalysts.
A catalyst is a substance that speeds up the reaction but is not involved in the reaction. It
is neither a reactant nor a product. It is represented in the chemical equation above the
arrow. In general, items that are placed above the arrow represent conditions that must be
met before the reaction can take place at a reasonable rate. One of the most common
catalysts is heat. This is shown as a triangle (∆) above the arrow. Some other catalysts are
light (hυ) or metals/metallic salts (Pt, MnO2, KI, etc).
Types of Chemical Reactions
There are five general types of reactions. Most chemical changes fit into these categories.
Occasionally, some can apply to two. Recognizing what type of reaction you’re dealing
with allows you to be able to predict what products will form.
1. Combination Reactions (aka: synthesis reaction)
A combination or synthesis reaction is a chemical change where two or more substances
react to form a single new substance. It often involves two or more elements combining
but it can also involve compounds.
When you have a metal reacting with a nonmetal, there is one possible product which is a
single ionic compound (the metal cation and the nonmetal anion).
Ex: 2Mg(s) + O2(g) → 2 MgO(s)
Ex: 2Na(s) + Cl2(g) → 2 NaCl(s)
When you have a transition metal reacting with a nonmetal, there are as many products
possible as charges that the transition metal can have. The product is a mixture of the
possible compounds.
Ex: Fe(s) + S(s) → FeS(s)
(produces iron (II) sulfide)
2Fe(s) + 3S(s) → Fe2S3(s)
(produces iron (III) sulfide)
When two nonmetals react, more than one product is possible based on the ratio of the
nonmetal reactants involved. The product is a mixture of the possible compounds.
Ex: S(s) + O2(g) → SO2(g)
2S(s) + 3O2(g) → 2SO3(g)
Ex: Write and balance the eqn. for beryllium reacting with oxygen
2Be + O2 → 2BeO
Ex: Write and balance the eqn. for the formation of aluminum chloride.
2Al + 3Cl2 → 2AlCl3
Most combination reactions release large amounts of energy, usually in the form of heat.
The driving force of the reaction is to form more stable bonds. Becoming more stable
means becoming lower energy, means energy is released when the product compounds
are formed.
2. Decomposition Reactions
A decomposition reaction is the reverse of a combination reaction. It is a chemical
change in which a single reactant compound breaks down into two or more simpler
products.
Ex: 2H2O → 2H2 + O2
Ex: 2HI → H2 + I2
Energy is required for these types of reaction, usually in the form of heat, electricity, or
light. Why? Decomposition involves breaking stable compound bonds which requires
energy to do.
When you have a binary compound, it decomposes into the 2 component elements. When
there is a compound with more than 2 elements decomposing, it is difficult to always
predict what the products will be.
Ex: H2CO3 → H2O + CO2
3. Single-Replacement Reactions
A single-replacement reaction is a chemical change in which one element replaces a
second element of the same type in a compound. Both the reactants and the products
consist of one element and one compound. The compound can be thought of as two
people dancing and the element cuts in and switches places with the element of the same
type (metal and metal, nonmetal and nonmetal) in the compound.
Whether the reaction will occur or not depends on the activity of the element doing the
replacing. Metals are the most common replacing element in this type of reaction,
although halogens also replace other halogens. Table 11.2 shows the activity series of
metals. The activity of an element describes how reactive it is. For a single-replacement
reaction to happen, the element trying to switch into the compound must have a higher
activity than the element of the same type in the compound.
Activity series of metal (from most active to least active):
Li, K, Ca, Na, Mg, Al, Zn, Fe, Pb, H, Cu, Hg, Ag.
Metals not seen in the list can be considered to be inactive.
A reactive metal will replace any metal with a lower activity. For example, Fe will
replace Cu in a compound but not Al.
Halogens can also displace other halogens. The activity of halogens decreases as you go
down the group (most to least active: F, Cl, Br, I). Br is more active than I so
Br2 + NaI → NaBr + I2 occurs but if you reacted Br2 + NaCl, there would be no reaction
because Br is less active than Cl.
Ex: Zn + H2SO4 → ZnSO4 + H2
Ex: Cl2 + 2NaBr → 2NaCl + Br2
Ex: Cu + Pb(NO3)2 → No reaction
4. Double-Replacement Reaction (aka: Double-Displacement Reactions)
A double-replacement reaction is a chemical change involving an exchange of the
positive ions (the cations) between two compounds. The reactants are always two “ionic”
compounds and are usually in the form of aqueous solutions.
In order for a double-replacement reaction to occur, one of the products must be either
(1) a solid precipitate
(2) a gas (CO2, O2, etc)
(3) a small molecular compound (H2O)
The formation of one of these products is the driving force of the reaction and if it is not
possible to produce one of these three things, there will be no reaction.
Ex: CaBr2(aq) + 2AgNO3(aq) → 2AgBr(s) + Ca(NO3)2(aq)
Ex: FeS(s) + 2HCl(aq) → H2S(g) + FeCl2(aq)
Ex: Ca(OH)2(aq) + 2HCl(aq) → CaCl2(aq) + 2H2O(l)
5. Combustion Reactions
A combustion reaction is a chemical change where an element or compound reacts with
oxygen. Oxygen (O2) is always one of the reactants. The other reactant is known as the
“fuel.” All combustion reactions produce a large amount of energy, usually in as heat,
light, or both.
Often, fuels are hydrocarbons (a compound mainly composed of H and C).
methane - CH4; propane – C3H8; butane – C4H10; octane (gasoline) – C8H18
The complete combustion of a hydrocarbon always produces CO2 and H2O. If the
oxygen is limited, pure C (soot) and CO (carbon monoxide gas) may also be produces).
If the oxygen is removed or depleted, the reaction will end. For example, covering a
candle (removing its source of oxygen) makes the flame extinguish.
Ex: Car Engine Combustion
2C8H18(l) + 25O2(g) → 16CO2(g) + 18H2O(g)
Combustion reactions occur with other types of fuel (fuel does not have to be a
hydrocarbon).
Ex: 2Mg(s) + O2(g) → 2MgO(s)
This is an example of a reaction that can be classified as two types. It is both a
combustion reaction (because one of the reactants is oxygen) and a combination reaction
(because two reactants are combining to form only one product).
Summary of the 5 Types of Reactions
Reaction Type
Reactants
Procedure to Predict
Products
Combination
(aka: Synthesis)
2 or more
reactants;
(usually 2
elements)
The elements combine
together to form an ionic
compound (think of
charges and criss-cross)
or molecule
Decomposition
Combustion
SingleReplacement
DoubleReplacement
1 reactant
only
Fuel + O2
1 element +
1 ionic
compound
2 ionic
compounds
(usually
aqueous)
The reactant breaks up
into its components
(usually into the elements
making it up)
Products
Other
Comments
1 product only
energy usually
produced
2 or more
products
Remember some
elements cannot be
alone (diatomics)
energy/catalyst
usually required
If fuel is a hydrocarbon,
products are always
CO2 + H2O
CO2 + H2O
If fuel is an element, treat
as a combination reaction
1 product from
combination of
elements
The element alone
switches with the
element of the same type
in the compound
1 different
element +
1 different
compound
Split up the ionic
compounds into their
ions (each will have 1 +
and 1 – ion)
Switch the + OR the –
ions to form new
compounds
Usually 2 ionic
compounds
(but can be 1
ionic compound
+ a gas or water)
OR
energy usually
produced (heat and
light)
Must check
activity series to
see if reaction will
occur (element
switching in must
have greater
activity that the
one in the
compound or else
no reaction)
Must check
solubility rules one of the products
must be a solid,
gas, or liquid water
or else no reaction
Reactions in Aqueous Solution
Most of the world is water based. (70% of the Earth’s surface is water. 66% of the human
body is water.) Also, when reacting ionic compounds, they are more easily mixed in an
aqueous form than in their natural solid state. Therefore, it makes sense that many
reactions take place in aqueous environments.
When an ionic compound dissolves in water to make an aqueous solution, the crystalline
solid dissociates/separates into its cations and anions. Double replacement reactions
involve aqueous solutions of ionic compounds. These types of reactions are very
prevalent. In order for a double-replacement reaction to occur, one of the products bust be
either a solid precipitate, a gas, or water. These reactions that form a solid precipitate as
one of the products are specifically referred to as “precipitation reactions.”
A precipitate is a solid compound that falls out of solution because there is something
about it that makes it insoluble in water. General solubility rules can be determined based
on which double replacement reactions produce precipitates.
Solubility Rules:
This general set of rules for ionic compounds allows you to predict whether a product
will be soluble (will stay in the aqueous solution) or insoluble (form a solid and
precipitate out). Even though the rules may identify a salt that will precipitate, these are
just general guidelines and the actual precipitation depends on other factors as well such
as the concentration of ions in the mixture. Ions that are not listed below can be assumed
to form insoluble compounds.
Insoluble Compounds (compounds that form precipitates) include:
- Carbonates (CO3-2) (except when with an alkali metal or ammonium)
- Phosphates (PO4-3) (except when with an alkali metal or ammonium)
- Chromates (CrO4-2) (except when with an alkali metal or ammonium)
- Sulfides (S-2) (except when with an alkali metal, an alkali earth metal, or
ammonium)
- Hydroxides (OH-1) (except when with an alkali metal, an alkali earth metal, or
ammonium)
Soluble Compounds (compounds that stay in aqueous solution) include:
- Nitrates (NO3-1)
- Chlorates (ClO3-1)
- Acetates (C2H3O2-1)
- salts with alkali metals (Li+1, Na+1, K+1, etc)
- salts with ammonium (NH4+1)
- Sulfates (SO4-2) (except in salts with Pb+2, Ag+1, Hg+2, Ba+2, Sr+2, and Ca+2)
- Chlorides (Cl-1), Bromides (Br-1), and Iodides (I-1) (except in salts with Ag+1,
Pb+2, and Hg+2)
Reactions in aqueous environments can be represented by ionic equations. Ionic
equations are more descriptive versions of a reaction that show anything that is an
aqueous solution dissociated into ions – which is actually how aqueous ionic compounds
are. Complete ionic equation shows all aqueous ionic compounds as dissociated ions as
well as any other compounds in the reaction.
Ex:
AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)
Complete Ionic: Ag+1(aq) + NO3-1(aq) + Na+1(aq) + Cl-1(aq) → AgCl(s) + Na+1(aq) + NO3-1(aq)
The silver chloride produced is insoluble in water and precipitates as a solid compound
out of solution. That is why it remains as a compound in the complete ionic equation.
Only ionic compounds in aqueous solution are dissociated into their component ions in
ionic equations.
Some ions remain the same before and after the reaction. These are known as “spectator
ions” because they do not participate in the overall reaction – the main chemical change.
In describing the main reaction that takes place in a reaction, the spectator ions are left
out. Just like in math, anything that is the same on both sides is cancelled. This leaves
you with the “net ionic equation.” This is an equation for a reaction in solution that shows
only those particles that are directly involved in the chemical change. Net ionic equations
differentiate between ions that react to form a precipitate, a gas, or water, and ions that
simply remain in aqueous solution.
In the example above, the NO3-1(aq) and the Na+1(aq) are both spectator ions. Therefore the
net ionic equation for the reaction above is: Ag+1(aq) + Cl-1(aq) → AgCl(s).
In forming an ionic compound, the charges must be neutralized (equal to 0). Net ionic
equations are balanced not only by the number and types of atoms that they involve but
also by their charges. The net ionic charge on both sides of the reaction is always 0. The
compound formed must be neutral, therefore the reactants must sum to be neutrally
charged as well.
Before writing ionic equations (whether complete or net) make sure the equation is
balanced. These balancing coefficients carry through to the ionic equations as well.
Additionally, when writing the ions, you must say the total number of ions present in the
compound. For example, a reaction with 2H3PO4(aq) as a compound would be represented
by: 6H+1(aq) + 2PO4-3(aq) in the complete ionic equation.
Ex:
H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l)
Complete: 2H+1(aq) + SO4-2(aq) + 2Na+1(aq) + 2OH-1(aq) → 2Na+1(aq) + SO4-2(aq) + 2H2O(l)
Net:
2H+1(aq) + 2OH-1(aq) → 2H2O(l)