1. Chemical Reaction and Equations
Evidence of a Reaction
Chemical Equations
Balancing Chemical Equations

2. 8.1 Writing and Balancing Chemical Reactions and Equations
1. List four observations that suggest that a chemical reaction has taken place.
2. List three requirements for a correctly written chemical equation.
3. Write a word equation and a formula equation for a given chemical reaction.
4. Write a chemical equation by balancing a formula equation by inspection.

3. Evidence of a Chemical Reaction
The release of
light or
heat .

4. More Light and Heat Evidence

5. Evidence of a Chemical Reaction
The production of
a gas
(bubbles) .

6. More Gas Production Evidence
2 H2
2H2
HCl H2 Zn
H2O

7. Evidence of a Chemical Reaction
The formation of
a precipitate
(solid)
when two aqueous solutions are
mixed together.
Pb(NO3)2 KI
PbI2

8. When gases form a Precipitate
NH4Cl(s)

9. Evidence of a Chemical Reaction
PbI2
A color change is often evidence of a
chemical reaction.

10. Characteristics of Chemical Equations
The equation represents known facts.
The equation must contain correct formulas!
Reactant + Reactant  Product + Product
The equation must obey The Law of Conservation of Mass.
Mass of Reactants = Mass of Products

11. Word Equations
A chemical equation in which the reactants and products are represented by words (names).
Methane + Oxygen  Carbon dioxide + Water
The  means yields or produces
The + means “and”

12. Formula Equation
Reactants and products are described by using their symbols and/or formulas.
CH4 + O2  CO2 + H2O
A formula equation does not obey “The Law of Conservation of Mass.”

13. Chemical Equations
1. Use symbols and formulas to represent the reactants and products.
2. Adjust the coefficients in front of each reactant or product to identify the number of each reactant or product.
1 CH O2  1 CO H2O

14. Chemical Equations
3. Use special symbols in parenthesis to indicate the state of the reactants or products.
(s) = solid or (cr) Crystal
(l) = liquid
(g) = gas
(aq) = aqueous (dissolved in water)

15. Chemical Equations
4. Use symbols above the yield arrow to represent catalysts or special conditions.
heat Δ atm pressure MnO2

16. Reversible Reactions and Equations
Some chemical reactions occur in both directions.
N2 (g) H2 (g) NH3 (g)

17. Balanced Chemical Equations
The coefficients represent the relative amounts of reactants and products.
1 CH O2  1 CO H2O
1 molecule 2 molecules 1 molecule 2 molecules
1 mole moles mole moles

18. The Importance of a Balanced Chemical Equation cont.
Coefficients can be used to determine the relative masses of reactants and products
1 CH O2  1 CO H2O
1 mole moles mole moles
1(16.05g) + 2(32.00g) = 1(44.01g) (18.02g)
80.05 g = g
Law of Conservation of Mass 

19. A Chemical Reaction Evidence = ?
Reactants
CH4
2 O2

20. Light from a Big Bang!
Reactants
Products
CH4
CO2
2 H2O
2 O2
ENERGY

21. Balancing Chemical Equations
1.Write the Word (Name) Equation.
Methane + Oxygen  Carbon dioxide + Water
2. Write the Formula Equation:
CH4 + O2  CO2 + H2O

22. Balancing Chemical Equations
3. Change the Coefficients as needed.
4. NEVER change a subscript !
5. Balance different types of atoms one
at a time.

23. Balancing Chemical Equations
6. Balance elements that only appear once
on each side of the equation first.
7. Balance polyatomic ions that appear on both sides of the equation as one unit.
8. Balance “H” and “O” last.

24. 9. Elements Use the symbol from the periodic table for an element.
Na Au S C Sn P
The following elements always appear as diatomic molecules in chemical equations.
H2 N2 O2 F2 Cl2 Br2 I2

25. Checking Your Work
Count atoms to be sure that the equation is balanced.
Counting = Coefficient x subscript

26. Example One
Sodium metal (solid) combines with chlorine gas to produce solid sodium chloride.

27. Example One
Sodium metal (solid) combines with chlorine gas to produce solid sodium chloride.
Na (s) + Cl2 (g)  NaCl (s)
Remember that chlorine is diatomic!

28. Example One
Sodium metal (Solid) combines with chlorine gas to produce solid sodium chloride.
Na (s) + Cl2 (g)  NaCl (s)
2 Na (s) + Cl2 (g)  2 NaCl (s)

29. Example Two
When copper metal (solid) reacts with aqueous silver nitrate, the products formed are aqueous copper (II) nitrate and silver metal (solid).

30. Example Two
When copper metal (solid) reacts with aqueous silver nitrate, the products formed are aqueous copper (II) nitrate and silver metal (solid).
Cu (s) + AgNO3 (aq)  Cu(NO3)2 (aq) + Ag(s)

31. Example Two
When copper metal (solid) reacts with aqueous silver nitrate, the products formed are aqueous copper (II) nitrate and silver metal (solid).
Cu (s) + AgNO3 (aq)  Cu(NO3)2 (aq) + Ag(s)
Cu (s) + 2 AgNO3 (aq)  Cu(NO3)2 (aq) Ag(s)

32. Example Three
Solid Iron (III) oxide and carbon monoxide gas react to form solid iron and carbon dioxide gas.

33. Example Three
Solid Iron (III) oxide and carbon monoxide gas react to form solid iron and carbon dioxide gas.
Fe2O3 (s) + CO(g)  Fe(s) + CO2(g)

34. Example Three
Solid Iron (III) oxide and carbon monoxide gas react to form solid iron and carbon dioxide gas.
Fe2O3 (s) + CO(g)  Fe(s) + CO2(g)
Fe2O3 (s) + 3 CO(g)  2 Fe(s) + 3 CO2(g)

35. THE END