1. What is the difference between a chemical reaction and physical change?
When you watch a reaction occur, what are some hints that it is a chemical reaction?

2. Ch. 11 Chemical Equations Reactions
Describing Chemical Reactions

3. Objectives
List three observations that suggest that a chemical reaction has taken place.
List three requirements for a correctly written chemical equation.
Write a word equation and a formula equation for a given reaction.
Balance a formula equation by inspection.

4. Chemical Reactions when a substance changes identity
reactants- original
products- resulting
law of conservation of mass
total mass of reactants = total mass of products

5. Chemical Reactions chemical equation
represents identities and relative amounts of reactants and products in the chemical reaction
uses symbols and formulas

6. Hints of Chemical Rxn heat or light gas bubbles precipitate
can also happen with physical changes
gas bubbles
means a gas is being created as product
precipitate
solid is being created
color change

7. Writing Chemical Equations
most pure elements
written as elemental symbol
diatomic molecules
molecule containing only 2 atoms
some elements normally exist this way
H2, O2, N2, F2, Cl2, Br2, I2
other exceptions
sulfur: S8
phosphorus: P4

8. Word Equations uses names instead of formulas
helps you to write formula equation

9. Example
Description:
Solid sodium oxide is added to water at room temperature and forms sodium hydroxide.
Word Equation:
sodium oxide + water  sodium hydroxide
Formula Equation:
Na2O + H2O  NaOH

10. Symbols Used in Equations
after a formula:
(s) solid
(l) liquid
(aq) aqueous: dissolved in water
(g) gas
yields
reversible
above arrow:
or heat heated
MnO2 or Pt catalyst
25°C specific T requirement
2 atm specific P requirement

11. Text Pg. 323 Chart of symbols used in chemical equations

12. List three observations that suggest that a chemical reaction has taken place.

13. HCl hydrochloric acid H2SO4 sulfuric acid HNO3 nitric acid
Acids you have to know!
HCl hydrochloric acid
H2SO4 sulfuric acid
HNO3 nitric acid
H3PO4 phosphoric acid
HC2H3O2 acetic acid

14. Write the chemical equation from the following description:
Zinc metal is added to hydrochloric acid to create zinc chloride and hydrogen gas.

15. Aluminum reacts with oxygen to produce aluminum oxide
Al + O  Al2O3
Al + O2  Al2O3
Al3 + O  Al2O3

16. Aluminum reacts with oxygen to produce aluminum oxide
Al + O  Al2O3
Al + O2  Al2O3
Al3 + O  Al2O3

17. Phosphoric acid is produced through the reaction between tetraphosphorus decoxide and water
H3PO4  P4 + H2O
H3PO4 + H2O  P4
P4O10 + H2O  H3PO4

18. Phosphoric acid is produced through the reaction between tetraphosphorus decoxide and water
H3PO4  P4 + H2O
H3PO4 + H2O  P4
P4O10 + H2O  H3PO4

19. Iron(III)oxide reacts with carbon monoxide to produce iron and carbon dioxide
FeO + CO  Fe + CO2
Fe2O3 + CO  Fe + CO2
Fe + CO  Fe2O3 + CO2

20. Iron(III)oxide reacts with carbon monoxide to produce iron and carbon dioxide
FeO + CO  Fe + CO2
Fe2O3 + CO  Fe + CO2
Fe + CO  Fe2O3 + CO2

21. Coefficients whole numbers in front of formula
distributes to numbers of atoms in formula
specifies the relative number of moles and molecules involved in the reaction
used to balance the equation

22. Balancing Equations ONLY add/change coefficients- NEVER subscripts!!!
balance one type of atom at a time
balance polyatomic ions first
balance atoms that appear only once second
balance H and O last
simplify if you can
Check at end!

23. Rules for writing and balancing equations – Pg. 327 in text.

24. Writing Equations
Write Word equations to help you organize reactants and products
Be sure to include symbols showing states of each reactant and product
Be sure to write the correct formula for each (crossing over for ionic compounds!)
Check your balancing of the equation when you are finished

25. Example 1
Description:
Aqueous iron III oxide reacts with hydrogen gas to produce iron metal and liquid water
Word Equation:
Iron III oxide + hydrogen gas 
iron + water

26. Example 1 Formula Equation: Fe2O3 (aq) + H2 (g)  Fe (s) + H2O (l)
Balanced Formula Equation
Fe2O3 (aq) + 3H2 (g)  2Fe (s) + 3H2O (l)

27. Example 2
Solid calcium metal reacts with water to form aqueous calcium hydroxide and hydrogen gas.
calcium + water 
calcium hydroxide + hydrogen
Ca(s) + H2O(l)  Ca(OH)2(aq) + H2(g)
Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g)

28. Example 3
solid zinc metal reacts with aqueous copper (II) sulfate to produce solid copper metal and aqueous zinc sulfate
zinc + copper (II) sulfate 
copper + zinc sulfate
Zn(s) + CuSO4 (aq)  Cu (s) + ZnSO4 (aq)
Zn(s) + CuSO4 (aq)  Cu(s) + ZnSO4 (aq)

29. Example 4
Hydrogen peroxide in an aqueous solution decomposes to produce oxygen and water
hydrogen peroxide  oxygen + water
H2O2 (aq)  O2 (g) + H2O (l)
2H2O2 (aq)  O2 (g) + 2H2O (l)

30. Example 5
Solid copper metal reacts with aqueous silver nitrate to produce solid silver metal and aqueous copper (II) nitrate
copper + silver nitrate 
silver + copper (II) nitrate
Cu (s) + AgNO3 (aq)  Ag (s) + Cu(NO3)2 (aq)
Cu (s) + 2AgNO3 (aq)  2Ag (s) + Cu(NO3)2 (aq)

31. Example 6
Carbon dioxide gas is bubbled through water containing solid barium carbonate, creating aqueous barium bicarbonate
carbon dioxide + water + barium carbonate  barium bicarbonate
CO2 (g) + H2O (l) + BaCO3 (s)  Ba(HCO3)2 (aq)

32. Example 7
Acetic acid solution is added to a solution of magnesium bicarbonate to create water, carbon dioxide gas, and aqueous magnesium acetate.
acetic acid + magnesium bicarbonate 
water + carbon dioxide + magnesium acetate
HCH3COO (aq) + Mg(HCO3)2 (aq) 
H2O(l) + CO2 (g) + Mg(CH3COO)2 (aq)
2HCH3COO (aq) + Mg(HCO3)2 (aq) 
2H2O(l) + 2CO2 (g) + Mg(CH3COO)2 (aq)

33. Write the balanced formula equation for:
Lithium metal is added to a solution of aluminum sulfate to make aqueous lithium sulfate and aluminum metal.

34. Types of Chemical Reactions

35. Types of Chemical Reactions
5 basic types discussed here
not all reactions fall in these categories
you should be able to:
categorize a reaction
predict the product(s)

36. 1. Synthesis also called combination reaction reactants:
more than one
can be elements or compounds
products: only one compound
A + X  AX
where A is the cation and X is anion

37. 1. Synthesis Rubidium and sulfur Rb (s) + S8 (s)  Rb2S (s)
Magnesium and oxygen
Mg (s) + O2 (g)  MgO (s)
Sodium and chlorine
Na (s) + Cl2 (g)  NaCl (s)
Magnesium and fluorine
Mg (s) + F2 (g)  MgF2 (s)

38. 1. Synthesis calcium oxide and water CaO(s) + H2O(l)  Ca(OH)2 (aq)
sulfur dioxide and water
SO2 (g) + H2O (l)  H2SO3 (aq)
calcium oxide and sulfur dioxide
CaO (s) + SO2 (g)  CaSO3 (s)

39. 2. Decomposition opposite of synthesis usually require energy
reactants: only one compound
products: more than one
usually elements but can be compounds
AX  A + X

40. 2. Decomposition water H2O (l)  H2 (g) + O2 (g) calcium carbonate
CaCO3 (s)  CaO (s) + CO2 (g)
calcium hydroxide
Ca(OH)2 (s)  CaO (s) + H2O (l)
carbonic acid
H2CO3 (aq)  CO2 (g) + H2O (l)

41. 3. Single Replacement
an element replaces a similar element in a compound
reactants: 1 element & 1 compound
products: 1 element & 1 compound
A + BX  B + AX
Y + AX  X + AY

42. 3. Single Replacement zinc and hydrochloric acid
Zn (s) + HCl (aq)  ZnCl2 (aq) + H2 (g)
iron and water
Fe (s) + H2O (l)  FeO (aq) + H2 (g)
magnesium and lead (II) nitrate
Mg (s) + Pb(NO3)2 (aq)  Mg(NO3)3 (aq) + Pb (s)
chlorine and potassium bromide
Cl2 (g) + KBr (s)  KCl (s) + Br2 (g)

43. 4. Double Replacement two similar elements switch places
reactants: 2 compounds
products: 2 compounds
AX + BY  BX + AY

44. 4. Double Replacement barium chloride and sodium sulfate
BaCl2 (aq) + Na2SO4 (aq)  NaCl (aq) + BaSO4 (s)
iron sulfide and hydrochloric acid
FeS (aq) + HCl (aq)  FeCl2 (aq) + H2S (g)
hydrochloric acid and sodium hydroxide
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
potassium iodide and lead (II) nitrate
KI (aq) + Pb(NO3)2 (aq)  KNO3 (aq) + PbI2 (s)

45. 5. Combustion Only responsible for one type
releases energy in form of heat/light
reactants: hydrocarbon + O2
H2O and CO2 as the only products
Ex: CH4 + O2  CO2 + H2O

46. C3H8(g) + O2(g)  CO2(g) + H2O(g)
Combustion
propane and oxygen
C3H8(g) + O2(g)  CO2(g) + H2O(g)

47. Practice
Classify each of the following reactions one of the five basic types:
Na2O + H2O  NaOH
synthesis
Zn (s) + 2HCl (aq)  ZnCl2 (aq) + H2 (g)
single replacement
Ca (s) + 2H2O (l)  Ca(OH)2 (aq) + H2 (g)

48. Practice C2H4 (g) + O2 (g)  CO2 (g) + H2O (g)
2H2O2 (aq)  O2 (g) + 2H2O (l)
decomposition
Cu (s) + 2AgNO3 (aq)  2Ag (s) +Cu(NO3)2 (aq)
single replacement
C2H4 (g) + O2 (g)  CO2 (g) + H2O (g)
combustion
ZnO (s) + C (s)  2Zn (s) + CO2 (g)

49. Practice Na2O (s) + 2CO2 (g) + H2O (l)  NaHCO3 (s)
synthesis
Ca(s) + H2O(l)  Ca(OH)2 (aq) + H2 (g)
single replacement
KClO3 (s)  KCl (s) + O2 (g)
decomposition
H2SO4 (aq) + BaCl2 (aq)  HCl (aq) + BaSO4 (s)
double replacement

50. Activity Series

51. Activity Series Pg. 333 Activity easier it reacts, higher the activity
ability of an element to react
easier it reacts, higher the activity
activity series
list of elements organized according to activities
from highest to lowest

52. Activity Series metals nonmetals
greater activity, easier to lose electrons
easier to become a cation
nonmetals
greater activity, easier to gain electrons
easier to become an anion

53. Activity Series
used to predict whether single replacement reactions will occur
most active is on top
an element can replace anything below it but not any above it

54. Practice zinc and hydrofluoric acid calcium and lead (II) nitrate
Zn (s) + HCl (aq)  ZnCl2 (aq) + H2 (g)
calcium and lead (II) nitrate
Ca (s) + Pb(NO3)2 (aq)  Ca(NO3)2 (aq) + Pb (s)
copper and lithium sulfate
Cu (s) + Li2SO4 (aq)  no reaction
bromine and iron (II) chloride
Br2 (l) + FeCl2 (aq)  no reaction