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Some Basic Notes – Chemical Reactions
Pre AP Chemistry (Bartlett)
2017
I.
a)
Name: _______________________________________________ Period: ____ Date: _________
Some basics and review:
Subscripts in a formula identify how may atoms of each element are in the compound
Examples:
CO2
Atoms = 1 Carbon and 2 Oxygen
Mg3(PO4)2
Atoms = 3 Magnesium, 2 Phosphorus, and 8 Oxygen
Subscripts outside a parentheses are multiplied by the subscripts inside the parentheses (distributed)
b) Diatomic Molecules – “The Halogen H-O-N-ies” These when in free or ELEMENTAL state are written as
diatomic, ex. Nitrogen gas or just nitrogen is N2, especially in equations.
Diatomic means there are 2 atoms that are covalently bonded (sharing electrons equally). Seven elements exist in
nature as diatomic molecules:
1.
2.
3.
4.
5.
6.
7.
c)
F2
Cl2
Br2
I2
H2
O2
N2
Writing formulas for ionic compounds: composed of metals (lose valence electrons to form cations) and nonmetals (gain electrons to make 8 valence electrons and are called anions).
Must check charges on compounds that you write:
Ex: sodium phosphide = Na P = Na1+ P3+ = Na3 P
If name has a Roman numeral, this will tell the charge on the cation.
Ex: iron (II) phosphide = Fe P = Fe2+ P3+ = Fe3 P2
d) Writing formulas for covalent compounds: composed of non-metals (share valence electrons). The Greek
prefixes will tell you the number of atoms, you do not need to write or check charges.
Ex: trinitrogen pentoxide = N3 O5
CO = carbon monoxide
e)
Symbols: Physical states: (s) solid, (l) liquid, (g) gas, (aq) aqueous or in water solution, (cr) crystalline
 means “yields” or “produces” , + reacts with or between multiple reactants or products,
 indicates a gas is formed
used after a product indicates a solid formed
reaction in equilibrium
indicates a reversible reaction
……. shows that heat is supplied to the reaction
Pt …. is used to indicate a catalyst is supplied, in this case, platinum. (catalysts speed up reactions)
indicates equivalent resonance structures
Symbols above and/or below the yield sign: These are conditions of the reaction, ex. light, catalysts, heat, initiators,
or by-products.
uv
2H2O2(aq) 
2H2O(l)  O2(g)
II.
1.
2.
Writing and Balancing Equations
A chemical equation represents on paper what happens in a chemical reaction.
Chemical Reaction – the way atoms are joined together changes
Reactants (starting materials)  Products (substances produced
3. Law of Conservation of Mass – atoms are neither created nor destroyed only rearranged, number of starting atoms
must equal number of atoms in the products.
Same of each atom on both side of the equation
Example: C + O2  CO2
1 C and 2 O on both sides = Balanced Equation
Proving that it follows the LCM: Add the total mass (atomic masses) and compare the “before (reactants) with the
“after” (products). C = 12.01 + 2 (16.00) = [12.01 + 2(16.00)]
*all masses are in grams/mole
4.
Skeletal or unbalanced Equation:
Example: C + O2  CO
5.
Balancing a Skeletal Reaction
MUST make sure all formulas are written correctly (check the charges and
criss-cross if needed). Make sure elements have been checked to see if diatomic.
Add compounds or elements until there is an equal number of each elements on both sides
1.
Can only add more of what you already have
2.
Do NOT change the formula, can’t make CO into CO2 Notice the subscript only applies to the element
it is next to!
3.
4.
5.

Shows all reactants and products but does NOT show balanced quantities
C + O2  CO
C + O2  CO + CO
C + C + O2  CO + CO
2C + O2  2CO
Coefficient - Large Number left of a compound’s formula. Represents the number of molecules/atoms.
Ex. 3 H2 O 3 is the coefficient, means 3 molecules of water. 2 is the subscript, means 2 atoms of H.
Balancing Example
Do these Example Problems:

___ Ba + ___ Cl2 

___ Al2O3  ___ Al +

___ H2 + ___ N2  ___ NH3

___ KClO3  ___ KCl + ___ O2

___ CH4 + ___ O2  ___ CO2 + ___ H2O
___ BaCl2
___ O2
TIPS!
1) Write the skeleton equation ….. what is reacting and what is made…..
2) Check that formulas are correct (write in charges and check total charge).
3) Count the atoms of the elements in the reactants & the products. Make a list if it helps.
4) Change the coefficients to equalize the number of each element in reactants and
products. Work with one element at a time. Do H and especially O last!
[NEVER CHANGE A SUBSCRIPT!!!!!!!!!!!!!!!!!!!!!!!!!].
H2O is not the same as H2O2.
5) Write the coefficients in lowest ratio.
Ex 2 H2 + O2
III.
 2 H2O
not
4 H 2 + 2 O2
 4 H2O
Reaction Types
1. Synthesis/Combination: Two or more reactants combine to make one product
Example - C + O2  CO2
2.
Decomposition: One reactant breaks down into two or more simpler products.
Example - CO2  C + O2
3.
Single Replacement: One free element replaces another element in a compound. Use the Activity
series of metals to determine if the free metal will react with the bonded metal.
Metal Activity Series
See it on your Ion Sheet also.
 The activity series ranks the relative reactivity of metals.

Not all single replacement reactions will take place. The single element must be more reactive (stronger)
than the element it wants to replace in order for the reaction to happen.
[Rb K Na Li Ba Sr Ca] [Mg Al Mn Zn Cr Fe Cd] [Co Ni Sn Pb] [H2 Sb As Bi Cu Hg] [Ag Au Pt]
(most reactive)
(very unreactive)
(w/acids; react w/cold W)




(w/acids; react w/steam)
(w/acids; not w/ W) (w/O2 not w/most acids or W)
H may be displaced from acids (released as H2)
< or >
H2 is given off when a metal reacts with water:
metal + H(OH)  H2 + metal(OH)
Water/Acid: Single replacement reactions with WATER (H2O)
depends on metal reactivity & water temperature.
Examples:
a. Al + H2O(steam)  H2 + Al(OH)3
b. Cu + H2O  NO REACTION
Metal replaces metal:
See activity series
Example - K + NaCl  Na + KCl
K is more active than Na so they will react
Example - Fe + NaCl  no reaction
Na is more active than Fe so they will NOT react
Do these: Mg + 2HCl  _____________________
H2 + Al2O3  ____________________
Try these Example Problems:
Fe
Li
Cu
Na

+
+
+
+
Cu(SO4)  _____________________
Zn(CO3)
 _________________
H2O
 ___________________
H2O  _____________________
Nonmetal replaces nonmetal:
Ni + NaCl
 _____________________
Al + CuCl2  _____________________
Ca + H2(SO4)  ___________________
See Halogen group on PT
Ex – F2 + NaCl  Na + KCl F is more active than Na so they will react
Ex – Br2 + NaCl  no reaction Cl is more active than Br so they will NOT react
Try these examples: Br2 + AlF3  ______________________
F2 + Al I3  ______________________
4.
Double Replacement: Two compounds switch element partners.
Example:
MgCl2 + NaBr  NaCl + MgBr2
Always switch the 2 cations. To see if there is a visible reaction, check the solubility of the compounds. If
all are soluble in water, then no visible reaction even though there will be a dynamic equilibrium where the
reaction is reversible.
5.
Combustion: A hydrocarbon (compound composed of Hydrogen &
Carbon) reacts with oxygen (aka Burning). A combustion reaction takes
place when a fuel and an oxidizing agent, or oxidant, react, releasing
energy in the form of heat, and sometimes light.
A combustion reaction can be either.........
a) Complete – plenty of O2 and products will be CO2 and H2O, metabolism produces these 2
products but at a slower pace and with more steps. Usually burns with blue flame.
C10H8 + 12 O2 → 10 CO2 + 4 H2O (burning of naphthalene)
b) Incomplete – not a lot of O2 , products will be CO and H2O unless otherwise specified, charcoal
(Carbon) can be made by this process. Usually burns with yellow flame.
(methane) 2 CH4(g) + 3 O2(g) → 2 CO(g) + 4 H2O(l)
(propane) 2C3H 8 (g) + 7O2 (g)  2C(s) + 2CO(g) + 2CO2(g) + 8H2O (g)
Try this:
Complete combustion of propane:
________________ + _________________
IV.
 _______________ + _______________
Other types of reactions:
Redox: Oxidation is loss of electrons. Reduction is gain of electrons. LEO the lion goes GER
For example, in the extraction of iron from its ore.
Because both reduction and oxidation are going on side-byside, this is known as a redox reaction.
TRY: 2 S2O32−(aq) + I2(aq) → S4O62−(aq) + 2 I−(aq)
ANS= ________ is oxidized, _______ is reduced
Hydrolysis: The “splitting” of a large molecule into smaller parts by
addition of water.
Neutralization Reactions: Special types of double displacement
reactions that involve the reaction between an acid and base to form a salt and water.
acid + base → salt + water
Heat is usually given off in neutralization reactions.
Mg(OH)2 (s) + 2HCl (aq) → MgCl2 (aq) + 2H2O (l)
Precipitation Reactions: Aqueous reactions that involve the formation of a precipitate (solid).
soluble compound + soluble compound → insoluble compound
2KI (aq) + Pb(NO3)2 (aq) → 2KNO3 (aq) + PbI2 (s)
The physical state symbol (aq) says the reaction is taking place in a water solution. The physical state symbol (s) says the
lead (II) iodide is a solid - therefore insoluble in the solution.
V.
Word equations:
a chemical reaction expressed in words rather than chemical formulas.
would be expressed as hydrogen gas + oxygen gas → steam
2 H2(g) + O2(g) → 2 H2O(g)
Try: sodium chloride and lead (II) nitrate yields lead (II) chloride and sodium nitrate
________________________________________
VI.
→
________________________________________
Energy Changes
1. Endothermic reaction - energy of products is greater than energy of reactants.
2.
Exothermic reaction - chemical energy of reactants is greater than chemical energy of products. X-axis
can be in time or energy measured in kilojoules (kJ).
3.
Addition of a catalyst to the Energy Diagram
Activation Energy is reduced allowing the reaction to
proceed faster.
A– Activation Energy = Energy needed to start reaction
B – Energy Released
C – ΔH (change in Enthalpy) = Final Energy – Initial Energy
VII.
Reaction Rates
1. Collision Theory: How fast a reaction takes place depends on the types, energy, and frequency of
collisions, (and even the orientation of the molecules) between the particles.
2.
Factors that affect Reaction Rates
a) Temperature
1. T Increases ∴ Increases collisions  speeds up reaction
2. T Decreases ∴ Decreases collisions  slows reaction
Bacteria metabolize faster at room temperature than in the refrigerator.
b) Surface area
1. Increases ∴ more interaction  speeds up
Powdered starch reacts much faster with O2 than a lump of it.
c) Concentration
1. Increases ∴ Increases number of particles – increases collisions – speeds up
6.0 Molar HCl reacts faster than 1.0 Molar HCl
d) Catalysts
1. Lowers the activation energy of the reaction  speeds is up
2. The catalyst is neither a reactant nor a product, so it is written over the arrow.
e) Decrease Volume – increase collision since concentration has increased
f) Stirring – increases collisions
VIII.
Equilibrium and Le Chatelier’s principle
1. Any product that is a precipitate or gas will be removed from the solution and cause the reaction to go to
completion.
2. When a change does not go to completion, an equilibrium is established
3. Equilibrium: forward and reverse paths of a change take place at the same rate.
Equilibrium symbol =
4. Example:
a) Temperature
1. Decreasing heat will drive the reaction in the direction that will produce more heat (forward)
2. Increasing heat will drive the reaction in the direction that will use up the excess heat (reverse)
b) Pressure
1. Decreasing pressure drives the reaction in the direction that will produce the most moles of product (higher
sum of coefficients) (reverse)
2. Increasing pressure drives the reaction in the direction that will produce the least moles of product (lowest
sum of coefficients) (forward)
c) Concentration
1. Decreasing the concentration of any compound with drive the reaction in the direction that will produce
more of that compound.
a. Decreasing N2 or H2 – reverse
b. Decreasing NH3 – forward
2. Increasing the concentration of any compound will drive the reaction in the direction that will use up the
excess amount of that compound.
a. Increasing N2 or H2 – forward
b. Increasing NH3 - reverse
Drawing reactions: Check sizes on atoms. Check which atoms are bonded together. Check diatomic elements.
all formula for correctness. Draw a representation of what occurs.
Check
Complete and Net Ionic equations:
Most reactions in an aqueous medium involve ions. However, when a reaction takes place two kinds of ions usually emerge:
"spectator" ions and "player" ions. The ions that are present as free ions on both sides of the equation (they do not
change) are called spectator ions.
Balanced equation:
Ba(NO3)2 (aq) + CuSO4 (aq)
BaSO4(s) + Cu(NO3)2
(aq)
Note: A soluble compound (dissolves in water) is aqueous (aq) , on the solubility tables an S = soluble but
write (aq) in the equation.
For insoluble compounds an I = insoluble but write an (s) in the equation.
Complete ionic equations are written to represent the actual state of the compounds. Any ionic compound that is soluble
will separate into its free ions. These will move around in the solution and are able to carry an electric current. The soluble
compounds are written as distinct ions. All insoluble compounds OR liquids (H2O) Or gases (ex. CO2 ) are kept together.
Ba2+ + 2 NO3- + Cu2+ + SO42-
BaSO4 (s) + Cu2+ + 2 NO3-
Net ionic equations are equations that show only the soluble, strong electrolytes reacting (these are represented as ions) and
omit the spectator ions, which go through the reaction unchanged. The spectator ions (those that have the same form on both
sides of the equation are cancelled out.
Ba2+ (aq) + SO42-
(aq)
BaSO4 (aq)
Precipitate: A solid that forms when solutions of soluble chemicals are mixed and react. Can be seen as cloudy, “snow”, or
a powder at the bottom of the container. Any ionic compound that is insoluble will exist as a solid mass either as little
pieces floating or as a mass of sloid at the bottom on the container. This effectively removes those ions from the solution.
When you go to AP Chemistry/college you will need to know the following solubility rules, so it is good to memorize
the main ones! And the most common polyatomic ions too!
Also keep in mind that net ionic equations, which are the bare bones of the chemical reaction, usually take place in aqueous
environments. Here are those solubility rules:
1.
Most alkali metal compounds and NH4 1+ compounds are soluble.
2.
Cl-, Br-, I- compounds are soluble, except when they contain Ag+ , Hg2 2+ , or Pb2+.
3.
F- compounds are soluble, except when they contain group 2A metals.
4.
ClO3 1- ClO4 1- NO3 1- NO3 1- NH4 1+ NH4 1+ and CH3COO- compounds are soluble.
5.
compounds are soluble, except when they include Ca2+, Sr2+, Ba2+, Ag+, Pb2+, or Hg22+ .
6.
Most CO3 2- PO4 3- CrO4 2- C2 O4 2-
S2-, OH-, and O2- compounds are insoluble.
7.
Group 2A metal oxides are classified as strong bases even though they are not very soluble.
Check the solubility tables on the back of your ion sheet or look online.
Examples: Tell if these are soluble or not:
silver nitrate = _____________
barium phosphate = _________
silver chloride = _________ Cesium carbonate = ____________
strontium sulfate = _____________
The two solubility rules that you will use the most are numbers 1 and 4. You should memorize that all group 1A metal and
ammonium compounds are soluble. As soon as you see a compound with NH4 1+ Li, Na, K, Rb, Cs, or Fr, you should
know that it’s soluble. Also, all nitrates are soluble—look at the end of the compound. If it ends in NO3 1- , you know that it’s
soluble.
What’s the big deal with solubility? Well, if the ion is soluble, it won’t form a precipitate, and this means it doesn’t react and
should be left out of the net ionic equation.
The key is first to write the compound’s chemical formula and then determine if it’s soluble. If it is soluble, then ionize it—if
it isn’t, don’t ionize it; leave it as a “molecule”.
Check the CRC handbook (Table of Ionic compounds) to see what color any insoluble compounds are (precipitates).
Now try writing some net ionic equations, using the rules above.
**Example
Write the balanced equation for a mixture of solutions of silver nitrate and lithium bromide.
_________________ ( ) + __________________ ( )
→ ________________ ( )
+ _________________ ( )
Write the complete ionic equation for a mixture of solutions of silver nitrate and lithium bromide.
______________________________________________________________________________________________
Cancel out the spectator ions in the above equation and write the net ionic equation:
________________________________________________________________________________