Transcript Chemical Equations

CHEMICAL REACTIONS
Reactants: Zn + I2
Product: Zn I2
1
Introduction
– Chemical reactions occur when bonds (between the
electrons of atoms) are formed or broken
– Chemical reactions involve
• changes in the chemical composition of matter
(the making of new materials with new properties)
• energy changes
– Symbols represent elements
– Formulas describe compounds
– Chemical equations describe a chemical reaction
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Chemical
Equations (eqns.)
What is a chemical equation?
How do you balance a chemical
equation?
How do you identify the type of
chemical equation?
Chemical Equations
A chemical equation is written as an
expression similar to a mathematic
equation that can be compared to a
recipe that a chemist follows in order
to produce desired results.
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Chemical Equations
Their Job: Depict the kind of reactants
and products and their relative
amounts in a reaction.
4 Al (s) + 3 O2 (g) ---> 2 Al2O3 (s)
The numbers in the front are called
stoichiometric coefficients
The letters (s), (g), and (l) are the
physical states of compounds.
Chemical Equations
Because of the principle of the
conservation of matter
(matter can not be created or
destroyed)
must be
balanced.
an equation
It must have the same number of atoms
of the same kind on both sides.
Law of Conservation of Energy MUST
ALSO BE FOLLOWED!
• Energy changes are written in
(endo-/ exothermic reactions)
Lavoisier, 1788
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Chemical Equations


All chemical equations have reactants
and products.
We express a chemical equation as
follows:
Reactants  Products
The arrow is equivalent to an “=“ math.
When we describe the equation we use
the word “yields” or “produces” instead of
equals


Example
C + O2  CO2
This reads “carbon plus oxygen react to
yield carbon dioxide”
Balancing a Chemical
Equation
A chemical equation is balanced when
the ions or atoms found on the
reactant side of the equation equals
that found on the product side.
The arrow can be considered the
balance point.
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Symbols Used in Equations
• Solid (s)
• Liquid (l)
• Gas (g)
• Aqueous solution (aq)
H2SO4
or
Pt
• Catalyst
• Escaping gas ()
• Depositing solid (↓)
• Change of temperature/ heat energy ( or +
3kJ or – 3kJ)
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Chemical Reactions and Chemical
Equations
Reactants
Products
Signs of Driving forces:
a) Color change
b) Formation of a solid/precipitate
c) Evolution of a gas
d) Evolution or absorption of heat
 In all cases, a more energetically stable (lower
energy product) is created
Balancing Equations
– When balancing a chemical reaction you may add
coefficients in front of the compounds to balance the
reaction, but you may not change the subscripts.
• Changing the subscripts changes the compound.
Subscripts are determined by the valence electrons
(charges for ionic or sharing for covalent)
– Think back to naming compounds/ determining
formulas. NaCl exists, because Na is + and Cl is -, but
NaCl2 does NOT exist since you would not have a
neutral compound! You can’t just add a number to a
formula to balance an equation.
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Subscripts vs. Coefficients
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• The subscripts
tell you how
many atoms of a
particular
element are in a
compound. The
coefficient tells
you about the
quantity, or
number, of
molecules of the
compound.
Chemical Equations
4 Al(s) + 3 O2(g)
---> 2 Al2O3(s)
This equation means
4 Al atoms + 3 O2 molecules
---produces--->
2 molecules of Al2O3
AND/OR
4 moles of Al + 3 moles of O2
---produces--->
2 moles of Al2O3
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Steps to Balancing Equations
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There are four basic steps to balancing a chemical equation.
1. Write the correct formula for the reactants and the products.
DO NOT TRY TO BALANCE IT YET! You must write the
correct formulas first.
**And most importantly, once you write them correctly DO NOT
CHANGE THE FORMULAS!
2. Find the number of atoms for each element on the left side.
Compare those against the number of the atoms of the same
element on the right side.
3. Determine where to place coefficients in front of formulas so
that the left side has the same number of atoms as the right
side for EACH element in order to balance the equation.
4. Check your answer to see if:
– The numbers of atoms on both sides of the equation are
now balanced.
– The coefficients are in the lowest possible whole number
ratios. (reduced)
Some Suggestions to Help You
Some helpful hints for balancing equations:
• Take one element at a time, working left to right
except for H and O. Metals, then nonmetals are a
good way, too. Save H for next to last, and O
until last.
• IF everything balances except for O, and there is no
way to balance O with a whole number, double all
the coefficients and try again. (Because O is
diatomic as an element)
• (Shortcut) Polyatomic ions that appear on both
sides of the equation should be balanced as
independent units
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Balancing Equations
___
2 H2(g) + ___ O2(g) ---> ___
2 H2O(l)
What Happened to the Other
Oxygen Atom?????
This equation is not balanced!
Two hydrogen atoms from a hydrogen
molecule (H2) combine with one of the
oxygen atoms from an oxygen molecule
(O2) to form H2O. Then, the remaining
oxygen atom combines with two more
hydrogen atoms (from another H2 molecule)
to make a second H2O molecule.
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Balance this equation!
Na + Cl2
Na- 1
Cl- 2
NaCl
Na-1
Cl-1
**note that the number of sodiums balance
but the chlorine does not. We will have to
use coefficients in order to balance this
equation.
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Inserting subscripts
Na + Cl2
Na- 1
Cl- 2
2 NaCl
Na- 1 2
Cl- 1 2
** Now the chlorine balances but the sodium
does not! So we go back and balance the
sodium.
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Finally balanced!
2Na + Cl2
Na- 1 2
Cl- 2
2 NaCl
Na-1 2
Cl-1 2
**Since the number of each element on the
reactant side and the product side of the
equation are equal, the equation is balanced.
Balancing
Equations
2 Al(s) + ___
3 Br2(l) ---> ___ Al2Br6(s)
___
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Balancing Equations
Sodium phosphate + iron (III) oxide 
sodium oxide + iron (III) phosphate
2 Na3PO4 +
3 Na2O +
Fe2O3 ---->
2 FePO4
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Balancing Equation Practice
1. CuCl3 + Li2S  Cu2S3 + LiCl
2. NiNO3 + KCl  NiCl + KNO3
3. FeCl3 + Na2O  Fe2O3 + NaCl
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Answers:
• 1. 2CuCl3 + 3Li2S  Cu2S3 + 6LiCl
• 2. NiNO3 + KCl  NiCl + KNO3
(already balanced)
• 3. 2FeCl3 + 3Na2O  Fe2O3 + 6NaCl
Types of Chemical
Reactions (rxns.)
Types of Reactions
•
•
Reactions are classified by their products.
There are five types of chemical reactions we
will talk about:
1.
2.
3.
4.
5.
•
Synthesis reactions
Decomposition reactions
Single displacement reactions
Double displacement reactions
Combustion reactions
You need to be able to identify the type of
reaction and predict the product(s)
Steps to Writing Reactions
•
Some steps for doing reactions
1.
2.
3.
Identify the type of reaction
Predict the product(s) using the type of
reaction as a model
Balance it
Don’t forget about the diatomic elements!
(BrINClHOF) For example, Oxygen is O2 as an
element.
In a compound, it can’t be a diatomic element
because it’s not an element anymore, it’s a
compound!
Synthesis or Combination
reactions
Synthesis (meaning to make) or combination
reactions are typified by their single product.
If you have a reaction in which at least 2
elements or compounds are reacted and
produce a single product, the reaction is a
synthesis reaction.
1. Synthesis reactions
•
•
Synthesis reactions are sometimes called
combination or addition reactions.
reactant + reactant  1 product
Basically: A + B  AB
•
•
•
Example: 2H2 + O2  2H2O
Example: C + O2  CO2
Note: Single Product! This is your clue that this is a synthesis or combination
reaction.
Synthesis Reactions
•
Here is another example of a synthesis reaction
Practice
•
•
•
•
Predict the products. Write and balance the
following synthesis reaction equations.
Sodium metal reacts with chlorine gas
2 Na(s) + Cl2(g)  2 NaCl(s)
Solid Magnesium reacts with fluorine gas
Mg(s) + F2(g)  MgF2(s) Balanced
Aluminum metal reacts with fluorine gas
2 Al(s) + 3 F2(g) 
2 AlF3(s)
2. Decomposition
Reactions
Decomposition reactions are really just the
opposite of a synthesis reaction.
Remember, if you can make a substance,
you should be able to break it back apart
into its components.
A good way to remember decomposition
reactions to to remember what happens
when something decomposes. It falls apart!
Decomposition Reactions
•
•
•
•
•
•
Decomposition reactions occur when a
compound breaks up into the elements or in a
few to simpler compounds
1 Reactant  Product + Product
Basically: AB  A + B
Example: 2 H2O  2H2 + O2
Example: 2 HgO  2Hg + O2
Note: Single Reactant! The single reactant is your clue that this is a
decomposition reaction.
Decomposition Reactions
•
Another view of a decomposition
reaction:
Decomposition
Exceptions
•
Carbonates and chlorates are special case
decomposition reactions that do not go to
the elements.
• Carbonates (CO32-) decompose to carbon
dioxide and a metal oxide
•
Example: CaCO3  CO2 + CaO
• Chlorates (ClO3-) decompose to oxygen gas
and a metal chloride
•
Example: 2 Al(ClO3)3  2 AlCl3 + 9 O2
• There are other special cases, but we will not
explore those in this class
Practice
•
•
•
Predict the products. Then, write and
balance the following decomposition
reaction equations:
Solid Lead (IV) oxide decomposes
PbO2(s)  Pb(s) + O2(g)
Aluminum nitride decomposes
2 AlN(s)  2 Al(s) + N2(g)
Practice
Identify the type of reaction for each of the
following synthesis or decomposition
reactions, and write the balanced equation:
N2(g) + O2(g)  2 NO (g)
BaCO3(s)  BaO(s) + CO2 (g)
Co(s)+ S(s)  Co2S3 (s)
NH3(g) + H2CO3(aq)  (NH4)2CO3(s)
I2 (s)
NI3(s)  N2 (g) +
3. Single Replacement Reactions
Single replacement reactions occur when one
chemical takes the place of another in a
reaction.
In the typical single replacement reaction, an
element trades places with one of the ions
in a compound.
Single Replacement Reactions
•
•
•
Single Replacement Reactions occur when
one element replaces another in a compound.
A metal can replace a metal (+) OR
a nonmetal can replace a nonmetal (-).
element + compound product + product
A + BC  AC + B (if A is a metal) OR
A + BC  BA + C (if A is a nonmetal)
(remember the cation always goes first!)
When H2O splits into ions, it splits into
H+ and OH- (not H+ and O-2 !!)
Single Replacement Reactions
•
Another view:
The Activity Series
► Not
all single replacement reactions will
occur.
► This depends upon the location of the
elements present in the activity series
► Elements above MAY replace elements
below; elements below MAY NOT replace
elements above them on the series
You will be given a copy of this!!!!
Single Replacement Reactions
Write and balance the following single
replacement reaction equation:
• Zinc metal reacts with aqueous
hydrochloric acid
Zn(s) + 2 HCl(aq)  ZnCl2 + H2(g)
Note: Zinc replaces the hydrogen ion in the
reaction
•
[If ZnCl2 + H2(g)  Zn(s) + HCl(aq) the reaction WOULD
NOT OCCUR because Hydrogen is below zinc on the
activity series]
Single Replacement Reactions
•
Sodium chloride solid reacts with fluorine gas
2 NaCl(s) + F2(g)  2 NaF(s) + Cl2(g)
Note that fluorine replaces chlorine in the compound
•
Aluminum metal reacts with aqueous copper
(II) nitrate
2 Al(s)+ 3 Cu(NO3)2(aq)3 Cu(s) +2 Al(NO3)3(aq)
4. Double Replacement Reactions
Double replacement reactions are identified
by two ions trading places and forming
new compounds.
Double Replacement Reactions
•
•
•
•
Double Replacement Reactions occur when a
metal replaces a metal in a compound and a
nonmetal replaces a nonmetal in a compound
Compound + compound  product + product
AB + CD  AD + CB
Notice that one ion from
compound AB replaces one
ion from compound CD.
Double Replacement Reactions
•
•
•
Think about it like “foil”ing in algebra, first and last
ions go together + inside ions go together
Example:
AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq)
Another example:
K2SO4(aq) + Ba(NO3)2(aq) 2 KNO3(aq) + BaSO4(s)
Solubility

For a double replacement reaction to have
occurred, a solid (precipitate) MUST be
formed OR water must be formed in an
acid-base reaction.
 There are rules to determine which of the
materials formed is the solid
 If no solid is formed, there is said to be no
reaction.
Figure 8.4: The forming of solid AgCl.
Solubility Tables


Solubility tables help determine which materials are soluble in
water and which are not
In general, Solubility Rules can be summarized as follows
1. All compounds containing alkali metal cations and the ammonium ion
2.
3.
4.
5.
6.
are soluble.
All compounds containing NO3-, ClO4-, ClO3-, and C2H3O2- anions are
soluble.
All chlorides, bromides, and iodides are soluble except those containing
Ag+, Pb2+, or Hg22+.
All sulfates are soluble except those containing Hg22+, Pb2+, Sr2+, Ca2+, or
Ba2+.
All hydroxides are insoluble except compounds of the alkali metals,
Ca2+, Sr2+, and Ba2+.
All compounds containing PO43-, S2-, CO32-, and SO32- ions are insoluble
except those that also contain alkali metals or NH4+.
You will be given a copy of this!!!!
Practice
•
Predict the products. Balance the equation
1.
HCl(aq) + AgNO3(aq) 
2.
Pb(NO3)2(aq) + BaCl2(aq) 
3.
FeCl3(aq) + NaOH(aq) 
4.
H2SO4(aq) + NaOH(aq) 
HNO3(aq) + AgCl(s)
PbCl2(s) + Ba(NO3)2(aq)
Fe(OH)3(s) + NaCl(aq)
H2O(l) + Na2SO4(aq)
Combustion Reactions
Combustion reactions are the ones
that burn (or explode!). There are
two types of combustion
reactions—complete or incomplete
reactions.
These reactions are identified by
their products. They either
produce carbon monoxide and
water or carbon dioxide and water.
Complete Combustion
Reactions
These reactions burn “efficiently”
which means they produce
carbon dioxide and water.
These reactions typically burn
cleanly and leave very little
residue behind.
5. Combustion
Reactions
• Combustion reactions
occur when a
hydrocarbon reacts with
oxygen gas.
• This is also called
burning!!! In order to
burn something you need
the 3 things in the “fire
triangle”:
1) A Fuel (hydrocarbon)
2) Oxygen to burn it with
3) Something to ignite
the reaction (spark)
Combustion Reactions
• In general:
CxHy + O2  CO2 + H2O
• Products in combustion
are ALWAYS carbon
dioxide and water.
(although incomplete
burning does cause some
by-products like carbon
monoxide)
• Combustion is used to
heat homes and run
automobiles (octane, as
in gasoline, is C8H18)
Complete Combustion
Reactions
CH4 + O2  CO2 + H2O
They may also be written:
CH4  CO2 + H2O
(O2 is usually written above the
arrow.)
Clue: CO2 (carbon dioxide) in the
product along with water
Combustion
Reactions
Edgar Allen Poe’s
drooping eyes and
mouth are potential
signs of CO
poisoning.
Combustion
• Example
•
C5H12 + 8 O2  5 CO2 + 6 H2O
• Write the products and balance
the following combustion
reaction:
10 CO2 + 22
11 H2O
• 2 C10H22 + 31O2 20
Incomplete Combustion
Reactions
Incomplete combustion reactions occur
when something does not burn
efficiently. This can cause a lot of
harm if the gases produced cannot
escape. Carbon monoxide,an
odorless and colorless gas, is
dangerous. People poisoned by this
gas usually become sleepy and can
die due to exposure.
Incomplete Combustion
Reactions
CH4 + O2  CO + H2O
These reactions may also be written by:
CH4  CO + H2O
(the O2 is usually written over the arrow.)
Clue: CO (Carbon monoxide as a
product.)
Mixed Practice
•
1.
2.
3.
4.
5.
State the type, predict the products, and
balance the following reactions:
BaCl2 + H2SO4  BaSO4 + HCl
C6H12 + O2  CO2 + H2O
Zn + CuSO4  ZnSO4 + Cu
Cs + Br2  CsBr
FeCO3  FeO + CO2
Predicting Products of
Reactions
Completing reactions requires
knowledge of the different reaction
types (sometimes called mechanisms).
You must first identify the reaction type
by the reactants. The only type of
reaction that cannot be predicted this
way is the combustion reaction since
the products are very similar.
First Step:

Identify reaction type
Example:
Al + O2 
Clue: 2 elements – Synthesis or
combination reaction
Second Step:

Write the net ionic equation for the
reactants
Al + O2  becomes
Al3+ + O2- 
Step 3

Using clues, complete reaction taking
care to write each formula correctly by
checking charges and “criss-crossing”
if necessary.
Al + O2  Al3+O2Al + O2  Al2O3
Predicting Products of
Reactions (cont.)
For Single Replacement reactions,
check activity series to make sure the
reaction goes.
 Once you write the molecular
equation, you should check for
reactants and products that are soluble
or insoluble. (Double Replacement
only)

Reactions in Aqueous
Solutions


a.k.a. Net Ionic Equations
Molecular Equations: shows complete formulas
for reactants and products
– Does not show what happens on the molecular level

Total (or Complete) Ionic Equations: All
substances that are strong electrolytes (are soluble
and dissociate) are written as their ions.
– Some ions participate in the reaction
– Some ions do NOT participate in the reaction-called
spectator ions.

Net Ionic Equations: show only the ions that
participate in the reaction
Writing Total Ionic Equations





Once you write the molecular equation
(synthesis, decomposition, etc.), you should
check for reactants and products that are
soluble or insoluble.
We usually assume the reaction is in water
We can use a solubility table to tell us what
compounds dissolve in water.
If the compound is soluble (does dissolve in
water), then splits the compound into its
component ions
If the compound is insoluble (does NOT dissolve
in water), then it remains as a compound
Writing Total Ionic Equations
Molecular Equation:
K2CrO4 + Pb(NO3)2 
PbCrO4 + 2 KNO3
Soluble
Insoluble
Soluble
Soluble
Total Ionic Equation:
2 K+ + CrO4 -2 + Pb+2 + 2 NO3- 
PbCrO4 (s) + 2 K+ + 2 NO3-
Net Ionic Equations
These are the same as total ionic
equations, but you should cancel out ions
that appear on BOTH sides of the equation
Total Ionic Equation:
2 K+ + CrO4 -2 + Pb+2 + 2 NO3- 
PbCrO4 (s) + 2 K+ + 2 NO3
(Spectator ions)
Net Ionic Equation:
CrO4 -2 + Pb+2  PbCrO4 (s)
Net Ionic Equations

Try this one! Write the molecular, total ionic, and net
ionic equations for this reaction: Silver nitrate reacts
with Lead (II) Chloride in hot water.
AgNO3 + PbCl2 
Molecular:
2 AgNO3 + PbCl2  2 AgCl + Pb(NO3)2
Total Ionic:
2 Ag+ + 2 NO3- + Pb+2 + 2 Cl-  2 AgCl (s) + Pb+2 + 2 NO3Net Ionic:
Ag+ + Cl-  AgCl (s)
Acid-Base Reactions

Acid:
– produces hydrogen ions (H+) in solution
(Arrhenius)
– proton donor

(Lewis)
Base:
– produces hydroxide ions (OH-) in solution
(Arrhenius)
– proton acceptor

(Lewis)
The reaction ALWAYS forms water and an
ionic compound (sometimes aqueous).
Acid-Base Reactions
 Example
HNO3 (aq) + KOH (aq) 
Molecular:
HNO3 (aq) + KOH (aq)  H2O (l) + KNO3 (aq)
Total Ionic:
H+ (aq) + NO3- (aq) + K+ (aq) + OH- (aq) 
H2O (l) + K+ (aq) + NO3- (aq)
Net Ionic:
H+ (aq) + OH- (aq)  H2O (l)
Oxidation-Reduction Reactions
 a.k.a.
Redox Equations
 Between a metal and a nonmetal forming
an ionic compound


Electron transfer occurs
Oxidation numbers: assigning an excess or
deficiency in electrons for each element (the
charge based on the compound).
Rules for Oxidation Number (ox. #)
Determination
1.
The sum of the oxidation numbers add up
to the charge
a. all elements have an ox. # of 0
b. ions of elements, ox. # is the charge (Cl-)
c. the sum of the ox. # of a complex ion equals
the charge (CO3-2 )
2.
H is 1+ when combined with a nonmetal
and 1- with a metal
H3PO4
H= 1+ PO4 = 3-
CaH2
Ca = 2+ H= 1-
Rules for Oxidation Numbers (cont.)
1.
2.
3.
F is always 1-; Cl, Br, I are 1- except when
combined with each other or O
O is 2- except when combined with F (F2O)
Group I is 1+ and Group II is 2+ in their
compounds
Recognizing Redox Rxns.
oxidation

1+
2 HCl (aq) + Mg (s)  MgCl2 (aq) + H20 (q)
0

Net:
2+
reduction
2 H+ (aq) + Mg2+ (aq)  Mg 2+ (aq) + H2 0 (q)


Loss of electron = oxidation
Gain of electron = reduction

“LEO the lion goes GER"
Half Reactions

Separate the individual oxidation and reduction
reactions.
 Look at electron movement
 Half rxn.:
Mg0 (aq)  Mg 2+ + 2 e2e- + 2 H+  H20

Net:
2 H+ (aq) + Mg0 (aq)  Mg 2+ (aq) + H2 0 (q)

Oxidizing agent: the one reduced (H+)
 Reducing agent: the one oxidized (Mg0)
Recognition of Redox rxns.



Oxidation # changes
Reactions with oxygen
Reaction of any element (forms a new
compound)
Balancing



Balance by mass
Balance by charge
Balance net ionic equation
Example Problem:
Fe (s) + Cl2 (aq)  Fe3+ (aq) + Cl- (aq)
1.
Balance by mass
Fe (s) + Cl2 (aq)  Fe3+ (aq) + 2 Cl- (aq)
2.
Write half reaction
Fe0 (s)  Fe 3+ + 3 e2e- + Cl2  2 Cl-
3. Balance by charge (want # of e- to
cancel)
(Fe0 (s)  Fe 3+ + 3e-) 2
+ (2e- + Cl2  2 Cl )
3
2 Fe0 (s) + 6e- + 3Cl2  2 Fe 3+ + 6e- + 6 ClFinal eqn.:
2 Fe (s) + 3Cl2  2 Fe 3+ + 6 Cl-
 Again
on types of reactions/ summary/
new ways of looking at it
The Myth!
MYTH: Little Mikey
(“Mikey Likes It”) from
the LIFE cereal
commercials in the
late 70s died when he
ate pop rocks then
drank a coke, causing
the pop rocks to
explode inside his
stomach.
The Myth!
Not surprisingly, this one
is completely false.
The Evidence:
– Mikey is alive and well. His
real name is John Gilchrist
and he’s an advertising
manager for a New York
Radio Station.
– There isn’t enough
carbonation in pop rocks to
release more than a tiny
amount of CO2 – much less
than in a coke. If the myth
were true, coke alone
should be able to explode
your stomach.
Types of Chemical Reactions
We can divide chemical reactions into five basic
types.
First Type:
– Synthesis Reaction:
– Looks like this:
A+B
AB
– In this reaction, substance A and substance B
“combine” to form a new compound.
– Substances A and B can be elements in atomic form,
elements in molecular form, or compounds.
Examples of Synthesis Reactions
2Na + S
Na2S
– This one is an example of two elements in
atomic form (Na and S) combining to form a
compound (sodium sulfide).
2H2 + O2
2H2O
– In this example, A and B are two elements in
molecular form (hydrogen and oxygen
molecules), and the product is water, which is
simply the chemical combination of hydrogen
and oxygen.
Examples of Synthesis Reactions
2Fe + O2
2FeO
– In this example, substance “A” is an element in
atomic form (Fe), and substance “B” is an element in
molecular form (O2). The result is a direct chemical
combination of the two elements (FeO, iron oxide,
which is “rust”).
CuO + H2O
Cu(OH)2
– This is an example where both substances going into
the reaction are molecules. The result is what you get
when you “add” all of the atoms in the reaction
together.
Chemical Reactions
Second Type:
– Decomposition Reaction:
– Looks like: AB
A+B
– Basically, the opposite of a synthesis reaction.
– A and B (i.e. the products of the reaction) can
be elements (atomic or molecular form), or
compounds.
– Example: 2HgO
2Hg + O2
Chemical Reactions
Third Type:
– Single Replacement
– Looks like this:
AB + C
AC + B
– As usual, A and B can be elements or
compounds, but C is typically an element.
– Example: 2H2O + 2Na
2NaOH + H2
Here a sodium atom replaces a hydrogen atom in
water.
The reaction is a bit easier to see if we write water
as hydrogen hydroxide and keep the freed
hydrogen as an atom:
HOH + Na
NaOH + H
Chemical Reactions
Fourth Type:
– Double Replacement:
– Looks like: AB + CD
AD + BC
– A, B, C, and D can be elements or
compounds (they are usually an element or a
polyatomic ion).
– Basically, the atoms (or polyatomic ions) in
the reactions “trade partners.”
Chemical Reactions
Examples:
– Most of our “practice” equations are examples
of this type of reaction.
– 2H3PO4 + 3Mg(OH)2
Mg3(PO4)2 + 6H2O
– H2SO4 + Mg(OH)2
MgSO4 + 2H2O
Chemical Reactions
Important Type of Double Replacement:
Acid-Base Neutralization Reactions
An acid: lots of free hydrogen ions (H+).
A base: lots of free hydroxide ions (OH-).
When you combine the two, the hydrogen ions in
the acid combine with the hydroxide ions in the
base to make water.
The substances that “donated” the hydrogen
and hydroxide ions also chemically combine to
make a special class of compounds called salts.
Acid-Base Neutralization
Chemically the reaction looks like this:
Acid + Base
Salt + Water
A classic example:
HCl + NaOH
NaCl + H2O
Hydrochloric Acid
Water
Sodium Hydroxide (Lye)
Sodium Chloride
(Table Salt)
Acid-Base Neutralization
Here’s the equation again:
– HCl + NaOH
NaCl + H2O
– Chemically, this is a double replacement
reaction:
The H traded its Cl for an OH
The Na traded its OH for a Cl.
Chemical Reactions
Hydrocarbon – A class of compounds
formed by combining hydrogen and
carbon. Often go by their common names.
Examples:
– CH4 is methane
– C2H6 is ethane
– C3H8 is propane
– C4H10 is butane
– C8H18 is octane
Chemical Reactions
Fifth Type:
– Combustion Reaction
– A hydrocarbon reaction with oxygen.
– “Complete” combusion: the products are
always CO2 and H2O (and a lot of energy in
the form of heat and light).
– “Incomplete” combusion: the products are
CO2, H2O, CO, and a “simpler” hydrocarbon
(one with fewer carbons).
Combustion Reactions
Complete reactions occur when there is
plenty of oxygen available for the reaction.
– Typically hard to ignite.
– Burn very hot.
Incomplete reactions occur when there
isn’t enough oxygen available for a
“complete” reaction.
– Easier to ignite.
– Cooler burning.
Combustion
Internal combustion engines (i.e. car
engines) typically don’t mix enough
oxygen in with the hydrocarbon (gasoline)
to support a complete reaction.
Result: Leftover hydrocarbons (toxic), and
the production of carbon monoxide (very
toxic). Both are major contributors to
pollution.
Combustion
Would letting more “air into the mix” (for a more
complete reaction) solve air pollution caused by
cars?
No!
– “Hard to ignite” means the car is more prone to
flooding and stalling.
– “Hotter burning” has several detrimental effects:
Hard on the engine’s oil.
Creates various nitrogen-oxygen compounds that can form
powerful acids (e.g. nitric acid), and other very toxic
pollutants.