1. Chemical Equations and Reactions Mrs. Partridge Loveland HS
Chemistry
Chemical Equations and Reactions
Mrs. Partridge
Loveland HS

2. Representing Chemical Changes
A chemical reaction is the changing of substances to other substances by the breaking of bonds in reactants and the formation of bonds in products.
Chemical reactions are represented by chemical equations which show what changes take place and the relative amounts of the various elements and compounds that take part in these changes.

3. Evidence for Chemical Reactions
A gas is produced
A precipitate (solid) is formed when 2 solutions are combined
A permanent color change occurs
Heat energy change is noted

4. The Logistics of a Chemical Equation
Every thing to the left of the arrow is considered a reactant. These are the starting substances.
Everything to the right of the arrow is considered a product. These are the substances that are formed as a result of the chemical reaction.
Coefficients go before the chemical formula and tell how many molecules, atoms or formula units of that species react or are formed.

6. The Logistics of a Chemical Equation
A double arrow indicates that the reaction is reversible.
A small triangle Δ over the reaction arrow means that heat was added to make the reaction occur.

7. The Logistics of a Chemical Equation
The small letters in parenthesis indicate the state of matter of the reactants and products
H2O(l )  2H2(g) + O2(g)
(s) = solid
(l) = liquid
(g) = gas
(aq) = aqueous  dissolved in water

8. The Logistics of a Chemical Equation
Arrows can also be used to indicate state of matter on the product side.
Up arrow = gas
Down arrow = precipitate

9. Reading a Chemical Equation
Read the equation by naming the state and the name of the reactants first. The arrow is read as “yields” or “produces”. The products are then read with the state and name.

10. Reading a Chemical Equation
EXAMPLE:
2Al (s) + Fe2O3 (s)  2Fe (s) + Al2O3 (s)
Solid aluminum reacts with solid iron (III) oxide to yield solid iron and solid aluminum oxide

11. WHEN GOING FROM THE WORD EQUATION TO THE CHEMICAL EQUATION, MAKE SURE TO PAY ATTENTION TO DIATOMIC MOLECULES!

12. Balancing Equations
The purpose of balancing equations is to ensure that the law of conservation of mass is maintained.
You must show an equal number of atoms for each element on both sides of the equation.
Remember, atoms can not be created or destroyed in a chemical reaction, they are ONLY rearranged!

13. Balancing Equations
To balance an equation, NEVER touch the subscripts. Add coefficients in front of chemical formulas until there are equal number of each type of atom on the reactant and the product side.

14. Balancing by Inspection
Step One: Write down what type of elements are present in the chemical equation
Step Two: Count the number of elements on the reactant side and on the product side
Step Three: Add coefficients to get the same number of elements on both sides of the equation

15. Balancing by Inspection
HINTS:
Balance compounds first
Treat polyatomics as one unit if they travel (stay) together
Leave any element that stands alone for LAST
Carbon and hydrogen are usually best left for last

16. Balancing by Algebra
Step One: Put a lower case letter in front of every chemical formula
Step Two: Write an equality that relates the reactant side to the product side for each type of element
Step Three: Assign the lower case letter with the highest coefficient the value of one
Step Four: Solve algebraically to determine the remaining letters
Step Five: The final answers must be WHOLE NUMBERS

17. Types of Chemical Reactions
Synthesis or Composition Reaction
Decomposition Reaction
Single Replacement Reaction
Double Replacement Reaction
Combustion Reaction

18. Synthesis (Composition) Reactions
Two or more substances combine to form a
new compound.
A + X  AX
Reaction of elements with oxygen and sulfur
Reactions of metals with Halogens
Synthesis Reactions with Oxides
There are others not covered here!

19. Decomposition Reactions
A single compound undergoes a reaction that
produces two or more simpler substances
AX  A + X
Decomposition of:
Binary compounds H2O(l )  2H2(g) + O2(g)
Metal carbonates CaCO3(s)  CaO(s) + CO2(g)
Metal hydroxides Ca(OH)2(s)  CaO(s) + H2O(g)
Metal chlorates 2KClO3(s)  2KCl(s) + 3O2(g)
Oxyacids H2CO3(aq)  CO2(g) + H2O(l )

20. Single Replacement Reactions
A + BX  AX + B
BX + Y  BY + X
Replacement of:
Metals by another metal
Hydrogen in water by a metal
Hydrogen in an acid by a metal
Halogens by more active halogens

21. The Activity Series of the Metals
Lithium
Potassium
Calcium
Sodium
Magnesium
Aluminum
Zinc
Chromium
Iron
Nickel
Lead
Hydrogen
Bismuth
Copper
Mercury
Silver
Platinum
Gold
Metals can replace other metals
provided that they are above the
metal that they are trying to
replace.
Metals above hydrogen can
replace hydrogen in acids.
Metals from sodium upward can
replace hydrogen in water

22. The Activity Series of the Halogens
Fluorine
Chlorine
Bromine
Iodine
Halogens can replace other
halogens in compounds, provided
that they are above the halogen
that they are trying to replace.
2NaCl(s) + F2(g) 
2NaF(s) + Cl2(g)
???
MgCl2(s) + Br2(g) 
No Reaction
???

23. Double Replacement Reactions
The ions of two compounds exchange places in an
aqueous solution to form two new compounds.
AX + BY  AY + BX
One of the compounds formed is usually a
precipitate, an insoluble gas that bubbles out of
solution, or a molecular compound, usually water.

24. Combustion Reactions
A substance combines with oxygen, releasing a large
amount of energy in the form of light and heat.
Reactive elements combine with oxygen
P4(s) + 5O2(g)  P4O10(s)
(This is also a synthesis reaction)
The burning of natural gas, wood, gasoline
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)

25. Net Ionic Equations
Double replacement reactions occur when 2 ionic compounds react by exchanging cations to form 2 different compounds.

26. Net Ionic Equations
When 2 solutions of ionic compounds are mixed the formation of a precipitate, gas or water may be one of the products.
One of the products is only slightly soluble and precipitates form solution
Na2S(aq) + Cd(NO3)2  CdS(s) + 2NaNO3 (aq)

27. Net Ionic Equations
One of the products is a gas that escapes out of the mixture.
2NaCN(aq)+H2SO4(aq)2HCN(g)+Na2SO4(aq)
One of the products is a molecular compound such as H2O.
Ca(OH)2(aq)+ 2HCl  CaCl2(aq) +2H2O(L)

28. To predict the formation of a precipitate we use SOLUBILITY RULES!
Net Ionic Equations
To predict the formation of a precipitate we use SOLUBILITY RULES!
**See appendix in book.

29. Net Ionic Equations
The world is water based. More than 70% of the earth’s surface is covered with water. Consequently, many important chemical reactions take place in water  aqueous solutions.

30. Net Ionic Equations AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq)
Most ionic compounds dissociate – separate into anions and cations when they dissolve in water.
A complete IONIC EQUATION shows dissolved ionic compounds as their free ions.
Ag+(aq)+ NO3-(aq) + Na+(aq) + Cl-(aq)
 AgCl(s) + Na+(aq) + NO3-(aq)

31. Net Ionic Equations
This equation can be simplified by eliminating the IONS that don’t participate in the reaction, spectator ions, from both sides of the equation.
Ag+(aq)+NO3-(aq)+Na+(aq)+Cl-(aq)AgCl(s) + Na+(aq)+NO3-(aq)
Rewritten, this equation is called the NET IONIC EQUATION, indicating only those particles that actually take part in the reaction.
Ag+(aq) + Cl-(aq)  AgCl(s)
Net ionic equations must have balanced ionic charges.