Transcript Chemical Equations

Chemical Equations
and Reactions
I. Chemical Equations
II. Chemical Equations
A. The equation must represent facts.
B. The equation must contain the correct
formulas for the reactants (on the left of
the arrow) and the products (on the right of
the arrow).
C. The law of conservation of mass and
energy must be satisfied. Therefore the
same number of atoms of each element
must appear on each side of a correct
chemical equation.
Symbol
Explanation of symbol
+

↔
(s)
↓
separates 2 or more reactants or products
“yield”, separates reactants from products.
indicates a reversible reaction
solid state. Placed after the formula of a substance
Alternative to (s) but used ONLY for a solid
PRODUCT, not reactants
indicates a liquid reactant or product
indicates an aqueous solution (where some solute
has been dissolved in water)
indicates a gaseous reactant or product
alternative to (g), but used ONLY for a gaseous
PRODUCT
indicates that heat is supplied to the reaction
(l)
(aq)
(g)
↑
Δ
A formula written above or below the  sign
indicates that it is used as a catalyst (something that
speeds up the reaction)
C. Diatomic and Polyatomic Molecules:
Element
Formula
State
Hydrogen
H2
gas
Nitrogen
N2
gas
Oxygen
O2
gas
Fluorine
F2
gas
Chlorine
Cl2
gas
Bromine
Br2
liquid
Iodine
I2
solid
Sulfur
S8
solid
Phosphorus P4
solid
III. Writing and Balancing
Equations
Example
Write a balanced equation for the following reaction:
Na + Cl2  NaCl
first write an atom inventory for the total
number of atoms of each element on each
side of the equation.
Na + Cl2  NaCl
Reactants
Products
# Na
# Na
# Cl
# Cl
Atom Inventory or Counting
Atoms
• you must be able to count atoms in order to balance
an equation. There are two ways to designate
numbers in a formula:
• subscripts – small numbers within a formula of a
compound. Tells the number of atoms in that
compound
• MgCl2 – 1 atom of Mg and 2 atoms of chlorine
• Sn3N2 – 3 atoms of tin and 2 atoms of nitrogen
• Coefficient – the large number in front of the formula
of a compound. Tells the number of molecules (in a
molecular compound) or formula units (in an ionic
compound) or atoms of an element.
• Remember that atoms cannot be created
or destroyed; we must balance an
equation using coefficients. Never
change a subscript to balance an
equation!!
Algebraic Method
Write the skeleton equation
Write a,b,c… on each substance
List the atoms present
Determine equalities for each atom
Assign a value 6 for A
Use A to determine other values
Reduce, if possible
Plug #s in as coefficients
Practice
__H2O  __H2 + __O2
__Pb(NO3)2 + __Na  __NaNO3 + __Pb
__C4H10 + __O2  __CO2 + __H2O
Examples
Zn
C2H6
+ H2O

Zn(OH)2
+ O2  CO2
+
+ H2O
Na2SO4 + Ba(NO3)2  BaSO4 +
NaNO3
H2
IV. Types of Chemical Reactions
A. Combination or Synthesis
• where 2 or more simple substances
(elements or compounds) combine to form
ONE complex substance
• 8Fe + S8  8FeS
• 2Sr + O2  2SrO
A. Combination or Synthesis
Li + P4 
N2 + Al 
Cl2 + Ca 
Na +
N2 
Special Combination or
Synthesis Reactions (Pre-AP
Only)
the metals that has a variable charge:
If one of these metals reacts with fluorine,
oxygen, or nitrogen (F, 0, N), these
nonmetals will
pull the metal to its HIGHEST charge or
oxidation number. Otherwise, when these
metals react in a
combination reaction, use their LOWEST
charge or oxidation number when forming
a new compound
Fe + O2 
Pb + N2 
Sn + S8 
Cu + P4 
Fe + Br2 
Cu + F2 
Practice:
B. Decomposition
a complex substance (compound)
decomposes into 2 or more simple
substances. Heat or electricity is usually
required.
Ex:
2NaCl  2 Na + Cl2
8MgS  8Mg + 2S8
Special decomposition reactions
to know (Pre-AP only):
• 2KClO3  2KCl + 3O2 • all metal chlorates decompose into
metal chloride + O2
• CaCO3  CaO + CO2
• metal carbonates decompose into a
metal oxide + CO2
• 2KOH  K2O + H2O
• metal hydroxides decompose into a
metal oxide + H2O
C. Combustion Reactions
• where oxygen reacts with another
substance, usually a hydrocarbon,
resulting in the release of energy, usually
heat or light.
• CH4 + 202  CO2 + 2H20
• Hydrocarbons always produce carbon
dioxide and water
Common Hydrocarbons
Name
Methane
Ethane
Propane
Butane
Pentane
Hexane
Heptane
Octane
Nonane
Decane
Molecular Formula
CH4
C2H6
C3H8
C4H10
C5H12
C6H14
C7H16
C8H18
C9H20
C10H22
Examples
C3H8 + O2 
C 2 H 2 + O2 
Ca + O2 
D. Single-Replacement
• occurs when one element displaces
another element in a compound. You
must check the “Activity Series of Metals”
to see if the “lone” element is active or
“strong” enough to displace the element in
the compound
Li
K
Ba
Ca
Na
Mg
Al
Zn
Fe
Ni
Sn
Pb
H
Cu
Hg
Ag
Au
Activity Series of Metals
Practice:
Li + KCl 
Sn + ZnCl2 
Sn + HCl 
Ni + HOH 
the halogens
F2
Cl2
Br2
I2
Decreasing strength
E. Double-Replacement
reactions
• occur when the cations (positive ions)
“switch” places. You do NOT need the
“activity series of metals” list in these
reactions. When you switch places, be
sure to correctly write the formula of the
new compound!!!!!
• Ex:
• 2 NaCl + Mg0  MgCl2 + Na20
Practice
CuS04 + Al(OH)3 
Ca3(P04)2 + ZnCr04 
Rules for Predicting Double
Replacement Reactions:
• 1. Predict the products of the double-replacement reaction and
indicate the solubility of both of the products
• Use the “Solubility Rules” handout (at end of notes) to determine the
solubility.
• If the compound is soluble that means that it will remain as ions in
the solution, if it is insoluble then the compound precipitated out of
the reaction (it became the precipitate or solid).
• 2. If at least one INSOLUBLE product is formed (which means a
precipitate will form) the reaction will occur!
• 3. If only SOUBLE products are formed then the reaction will NOT
occur (because no precipitate is formed)!
• 4. If water is produced the reaction will occur!
• 5. If the reaction occurs and one of the compounds formed is
soluble then that compound is written as ions and not as a
compound.
SOLUBILITY RULES
1. All salts whose cation is in Group 1 or is NH4+ are soluble – no
matter what the anion is.
2. All nitrates and nitrites are soluble.
3. All acetates are soluble.
4. All chlorates and perchlorates are soluble – no matter what the
cation is.
5. All chlorides are soluble except silver, lead (II), mercury
6. All bromides are soluble except silver, lead (II), mercury
7. All iodides are soluble except silver, lead (II), mercury
8. All flourides are INSOLUBLE.
9. All sulfates are soluble except silver, lead (II), mercury, calcium,
strontium and barium
10. All sulfides are INSOLUBLE except Group 1 and 2.
11. All phosphates, phosphites, carbonates, chromates, and
dichromates are INSOLUBLE unless the cation is in Group 1 or is
NH4+.
12. All hydroxides are INSOLUBLE except Group 1, calcium,.
Strontium, barium and ammonium
Net Ionic Equations
(Pre-AP Only)
• F. Net Ionic Equations (Pre-AP Only) – shows
only the compounds and ions that undergo a
chemical change in a double replacement
reaction
• Example: Na2S + Cd(NO3)2  Na+ + NO3 +
CdS(s)
• Step 1: Convert the chemical equation to an
overall ionic equation. All reactants are shown
as ions. For the products, all soluble ionic
compounds are shown as dissociated ions and
the precipitates are shown as solids.
• Na+ + S 2 + Cd 2+ + NO3 Na+ + NO3 +
CdS(s)
• Step 2: All spectator ions (ions that do not
take part in a chemical reaction and are
found as ions both before and after the
reaction) are removed from the equation.
• S2 + Cd2+  CdS(s)
Examples
BaCO3 + CuSO4  BaSO4(s) + CuCO3 (s)
K3PO4 + NaOH  no reaction (no ppt)
Na2S + Cd(NO3)2  Na+ + NO3 + CdS(s)
Types of Reactions Summary
Combination (synthesis)
A + B  AB
Decomposition
AB  A + B
Combustion
CxHy + O2  CO2 + H2O
Single Replacement
A + BC  AC + B
Double Replacement
AB + CD  AD + CB
Practice Predicting Products
l. AlCl3 
2. C2H4 + 02 
3. Zn + AgNO3 
4. H20 
Practice Predicting Products
5. Al + P4 
6. NaI +
MgS 
7. Cl2 + NaBr 
8. C6H1206 + O2 
Practice Predicting Products
1. AlCl3 + Na2CO3 
2. Ni + MgSO4 
3. Cl2 + K 
4. C5H12 + 02 
Practice Predicting Products (PreAP)
1. sodium metal is placed into water
2. methane gas is burned in the presence
of oxygen
3. potassium bromide solution is mixed with
chlorine gas
4. a solution of aluminum dichromate is
mixed with a solution of lithium oxalate
V. Oxidation Reduction
Reactions (Pre-AP only)
What is REDOX?
• Oxidation-Reduction (Redox) – involves a
transfer of electrons
• One specie is losing electrons
OIL – Oxidation is losing
Mg0  Mg+2 + 2 e-
• One specie is gaining
electrons
RIG – Reduction is
gaining
Mg+2 + 2 e-  Mg0
• The Species that is oxidized is the
reducing agent
• The Species that is reduced is the
oxidizing agent.
•
Mg0
+
O20
Mg+2O-2
Ox
Red ag
Ox ag
Red
REDOX reactions MUST:
• 1 Have a species that is oxidized and one
reduced – YOU cannot have one without
the other
• The number of electrons gained and lost
MUST be the SAME
• The number of atoms of each element
must be the same on both sides of the
equation
Balancing REDOX Reactions
• 1) Assign oxidation numbers to each atom
in the equation
• 2) Determine the substances oxidized,
reduced, oxidizing agent, reducing agent,
3) Write balanced half-reactions for the oxidation
and reduction reactions.
• Mg0
+
O 20
Mg+2O-2
Mg0  Mg+2 + 2 eO20 + 4 e-  2O-2

• 4) Multiply each equation so that the number of
electrons lost equals the number of electrons
gained.
• Mg0
+
O20
Mg+2O-2
• 2[Mg0  Mg+2 + 2 e-]= 2Mg0  2Mg+2 + 4 eO20 + 4 e-  2O-2
• 5) Add the two half-reactions. Place the
coefficients into the original equation.
• Mg0
+
O 20
Mg+2O-2
2Mg0
 2Mg+2 + 4 eO20 + 4 e 2O-2
2Mg0 + O20  2Mg+2O-2
• 6) Adjust other ions if necessary. Check all
atoms for conservation. Check hydrogen’s
and oxygen’s last.
2Mg + O2

2Mg O