1. Diatomic Gases and Halogens
H, O, N, F, Cl occur as diatomic gases
Covalent bond, attain shared octet
H2 (g)
O2 (g)
N2 (g)
F2(g), Cl2 (g)
Remaining halogens also diatomic
Br2(l), I2(s)

2. Chapter 9 – Chemical Reactions
9.1 Reactions and Equations
9.2 Classifying Chemical Reactions
9.3 Reactions in Aqueous Solutions

3. Section 9.1 Reactions and Equations
Chemical reactions are represented by balanced chemical equations.
Recognize evidence of chemical change.
Represent chemical reactions with equations.
Balance chemical equations.
Know which elements occur as diatomic molecules and the states of these elements at room temperature; correctly represent these elements when writing chemical equations which involve them. [not in chapter]

4. Section 9.1 Reactions and Equations
Key Concepts
Some physical changes are evidence that indicate a chemical reaction has occurred.
Word equations and skeleton equations provide important information about a chemical reaction.
A chemical equation gives the identities and relative amounts of the reactants and products that are involved in a chemical reaction.
Balancing an equation involves adjusting the coefficients until the number of atoms of each element is equal on both sides of the equation.

5. Chemical Reactions
Process by which atoms of one or more substances are rearranged to form different substances

6. reactant 1 + reactant 2  product 1 + product 2
Chemical Equations
Reactants are starting substances
Products are substances formed
reactant 1 + reactant 2  product 1 + product 2
+ separates 2 or more reactants or products
 separates reactants from products
Read as “reacts to produce” or “yields”

7. Parts of Balanced Equation
Subscript - Shows how many atoms of an element are present in a formula unit or molecule of this substance
Coefficient – Shows how many formula units or molecules are needed to balance the equation

8. 2 Li(s) + 2 H2O(l)  2 LiOH(aq) + H2(g)
Chemical Equations
2 Li(s) + 2 H2O(l)  2 LiOH(aq) + H2(g)
(s) identifies solid state
(l) identifies liquid state
(g) identifies gaseous state
(aq) identifies solution in water = aqueous solution
LiOH(l) not the same as LiOH(aq)

9. Evidence of Chemical Reactions
Temperature change
Color change
Flame, smoke
Odor
Gas evolution (bubbles)
Appearance of new phase (precipitate)

10. Word and Skeleton Equations
Word equation
Iron(s) + chlorine(g)  iron(III) chloride(s)
“Solid iron and chlorine gas react to produce solid iron(III) chloride”
Skeleton equation
Chemical formulas in place of words
Fe(s) + Cl2(g)  FeCl3(s)

11. Word and Skeleton Equations
Solid carbon reacts with solid sulfur to form liquid carbon disulfide
C(s) + S(s)  CS2(l)
Note: CS2 is not an ionic compound and ionic compound naming rules do not apply

12. Practice Write skeleton equations Problems 1-3, page 284

13. Equations and Atoms Fe(s) + Cl2(g)  FeCl3(s)  +
As written, 1 chlorine atom has been “created” – matter not conserved  +
One iron atom Two chlorine atoms
One iron atom Three chlorine atoms

14. Balanced Chemical Equations
2Fe(s) + 3Cl2(g)  2FeCl3(s) + 
Both sides have: Two iron atoms Six chlorine atoms

15. Balanced Chemical Equations
2Fe(s) + 3Cl2(g)  2FeCl3(s)
Skeleton equation has been balanced by inserting correct coefficients in front of reactants and/or products
Integers, not written if = 1
Lowest whole number ratio of amounts of reactants and products

16. Balancing Chemical Equations
Steps (see table 9.2, p 286)
1 – Write skeleton equation
H2(g) + Cl2(g)  HCl(g)  +
One H atom One Cl atom
Two H atoms
Two Cl atoms

17. Balancing Chemical Equations
2 – Count the atoms of elements in the reactants
H2(g) + Cl2(g)  HCl(g)
2 atoms H
2 atoms Cl

18. Balancing Chemical Equations
3 – Count the atoms of elements in the products
H2(g) + Cl2(g)  HCl(g)
1 atoms H
1 atoms Cl

19. Balancing Chemical Equations
4 – Change coefficients to make # of atoms of each elements equal on both sides of the equation
H2(g) + Cl2(g)  2HCl(g)
2 atoms H (both sides)
2 atoms Cl (both sides)

20. Balancing Chemical Equations
5 – Write the coefficients in their lowest possible whole number ratio
H2(g) + Cl2(g)  2HCl(g)
Coefficients as written already in lowest possible ratio

21. Balancing Chemical Equations
6 – Check your work.
Chemical formulas correctly written?
Number of atoms same on both sides of equation?
All states specified?

22. Balancing Chemical Equations Figure 9.6, p 288

23. Balancing Chemical Equations
When balanced, matter conserved
H2(g) + Cl2(g)  2HCl(g) + 
Two H atoms
Two Cl atoms
Two H atoms Two Cl atoms

24. Balancing Equation Strategy
Balance elements that occur in only one compound on each side first
Balance free elements last
Balance unchanged polyatomic ions as groups (NO-3, etc)
Fractional coefficients are acceptable as intermediate result; clear at end by multiplication by common divisor

25. Balancing Chemical Equations
Combustion of propane
C3H8(g) + O2(g)  CO2(g) + H2O(g)
Count atoms on each side
C H O C H O
Not balanced; try to balance C

26. Balancing Chemical Equations
C3H8(g) + O2(g)  3CO2(g) + H2O(g)
C H O C H O
Not balanced; try to balance H
C3H8(g) + O2(g)  3CO2(g) + 4H2O(g)

27. Balancing Chemical Equations
C3H8(g) + O2(g)  3CO2(g) + 4H2O(g)
C H O C H O
Not balanced; balance oxygen
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)
Done

28. Balancing Practice NH3(g) + O2(g)  N2(g) + H2O(l) ?
Balance N: 2 1  1 1
Balance H: 2 1  1 3
Balance O: 2 3/2  1 3
Integers:  2 6
4NH3(g) + 3O2(g)  2N2(g) + 6H2O(l)

29. Balancing Practice C4H10(g) + O2(g)  CO2(g) + H2O(g)
Balance C: 
Balance H: 
Balance O: 1 13/2 
Integers: 
2C4H10(g) + 13O2(g)  8CO2(g) + 10H2O(g)

30. Practice
Note: instruction to “write chemical equation” is understood to mean “write a balanced chemical equation”
Problems 4-6 page 287
Problems 3-8, page 980
Problems 13 page 283
Problems 64, 66, page 312

31. Chapter 9 – Chemical Reactions
9.1 Reactions and Equations
9.2 Classifying Chemical Reactions
9.3 Reactions in Aqueous Solutions

32. Section 9.2 Classifying Chemical Reactions
There are 5 classes of chemical reactions: synthesis, combustion, decomposition, single replacement, and double replacement. The replacement reactions have subclassifications.
Classify chemical reactions into one or more of 5 possible classes.
Identify the characteristics of different classes of chemical reactions, including any subclassifications that may apply (redox, formation of hydrogen in acid, etc.).

33. Section 9.2 Classifying Chemical Reactions
(cont.)
Use the activity series for metals and for halogens to correctly predict if a given pair of reactants will undergo a single replacement reaction.
Know that the products of the complete combustion of any hydrocarbon or carbohydrate are carbon dioxide and water.

34. Section 9.2 Classifying Chemical Reactions
Key Concepts
Classifying chemical reactions makes them easier to understand, remember, and recognize.
Activity series of metals and halogens can be used to predict if single-replacement reactions will occur.

35. Five Classes of Reactions
Synthesis
Combustion
Decomposition
Single Replacement (4 types)
Double Replacement (3 types)

36. CaO(s) + H2O(l)  Ca(OH)2(s)
Synthesis Reactions
Two or more substances react to form a single product A + B  C
2Na(s) + Cl2(g)  2 NaCl
CaO(s) + H2O(l)  Ca(OH)2(s)
2SO2(g) +O2(g)  2SO3(g)

37. Combustion Reaction Oxygen combines with substance
A + O2(g)  C or C + D
Energy released rapidly in form of heat & light
2H2(g) + O2(g)  2H2O(g)
(Also a synthesis reaction)
Other reactions involve combination with O2(g) but slow process (rusting)

38. CH4(g) + 2 O2(g)  CO2(g) + 2H2O(g)
Combustion
C(s) + O2(g)  CO2(g)
CH4(g) + 2 O2(g)  CO2(g) + 2H2O(g)

39. Combustion – Special Cases
Complete combustion of any hydrocarbon yields CO2 & H2O as products
CH4(g) + 2 O2(g)  CO2(g) + 2H2O(g)
2C4H10(g) + 13O2(g)  8CO2(g) + 10H2O(g)
Same holds true for complete combustion of a carbohydrate (contains CHO only)
C2H7OH = CH3CH2OH ethyl alcohol (ethanol)
4C2H7OH(l) + 13O2(g)  8CO2(g) + 14H2O(g)

40. Practice Synthesis & Combustion Problems 14-17, page 291 *
* Problem 17 has alternate accounting system for aqueous system

41. Decomposition Reactions
Single compound breaks down into two or more elements or new compounds
AB  A + B
Often require source of energy to start
Heat, light, electricity

42. Decomposition Reactions
NH4NO3(s)  N2O(g) + 2H2O(g)
2NaN3(s)  2Na(s) + 3N2(g)

43. Practice Problems 18-20 page 292 Problems 11-12 page 980
Problem 87, page 313

44. Replacement Reactions
Single-Replacement (SR)
Atoms of one element replace atoms of another element
A + BC  AC + B
A has replaced B in compound BC
SR reactions are also redox (oxidation reduction) reactions – will examine in more detail in section 19.1

45. Single-Replacement Reactions
Categories (will discuss each in detail)
Metal replaces H atom in H2O
Metal replaces H atom in HX(aq)
One metal (solid) replaces 2d metal in an aqueous ionic compound
Replacement of nonmetal in compound by another nonmetal (halogens)

46. Single-Replacement Reactions
Metal replaces H atom in H2O
Think of H2O as HOH
2Li(s) + 2H2O(l)  2LiOH(aq) + H2(g)
Li has replaced H in HOH
Reaction occurs readily for active metals – Li, Na, K, Rb, Ca, etc.
Elemental metal becomes a cation

47. Single-Replacement Reactions
Metal replaces H atom in HX(aq)
Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
Zn has replaced H in HCl
Reaction occurs readily for more metals than for reaction with water alone
Depends upon T and acid strength
Elemental metal becomes a cation

48. Single-Replacement Reactions
One metal (solid) replaces 2d metal in an aqueous ionic compound
M1(s) + M2X(aq)  M2(s) + M1X(aq)
M1 = metal # 1 – solid elemental metal
M2 = metal # 2 – in ionic compound
Replacement only happens when M1 is more active (more reactive) than M2

49. Single-Replacement Reactions
M1(s) + M2X(aq)  M2(s) + M1X(aq)
M1 = metal # 1 – solid elemental metal
M2 = metal # 2 – in ionic compound
M1 transformed from elemental metal to being a cation in an ionic compound
M2 transformed from cation in an ionic compound to being an elemental metal

50. Single-Replacement Reactions
One metal (as solid) replaces another metal in a compound dissolved in water
Cu(s) + 2AgNO3(aq)  2Ag(s) + Cu(NO3)2(aq)
Cu has replaced Ag in AgNO3
Cu more active than Ag (see following)
Cu(s)  Cu Ag+  Ag(s)
Note: Product could be CuNO3(aq)

51. Activity Series for Metals
Lithium
Rubidnium
Potassium
Calcium
Sodium
Magnesium
Aluminum
Manganese
Zinc
Most Active
Iron
Nickel
Tin
Lead
Copper
Silver
Platinum
Gold
Least Active

52. Single-Replacement Reactions
One metal (as solid) replaces another metal in a compound dissolved in water
Ni(s) + NaNO3(aq)  NR
NR = no reaction
Nickel is less active than sodium (see following slide)

53. Activity Series for Metals
Lithium
Rubidnium
Potassium
Calcium
Sodium
Magnesium
Aluminum
Manganese
Zinc
Most Active
Iron
Nickel
Tin
Lead
Copper
Silver
Platinum
Gold
Least Active

54. Practice – from Problem 9.2
Predict products and balance equation
Fe(s) + CuSO4(aq)  ???
Is Fe more active than Cu?
Yes
Fe(s) + CuSO4(aq)  FeSO4(aq) + Cu(s)
Balanced?
Fe(s)  Fe Cu2+  Cu(s)

55. Practice – from Problem 9.2
Predict products and balance equation
Mg(s) + AlCl3(aq)  ???
Is Mg more active than Al?
Yes
Mg(s) + AlCl3(aq)  Al(s) +MgCl2(aq)
Balanced? No
3Mg(s) + 2AlCl3(aq)  2Al(s) + 3MgCl2(aq)
Mg(s)  Mg Al3+  Al(s)

56. Single-Replacement Reactions
Replacement of nonmetal in compound by another nonmetal (halogens)
Most common for halogens
X21 + 2AX2  X22 + 2AX1
X1 = halogen number 1 X2 = halogen number 2
Replacement only happens when X1 is more active (more reactive) than X2
X1 elemental halogen becomes anion
X2 anion halogen becomes an element

57. Single-Replacement Reactions
Replacement of one halogen by another halogen
F2(g) + 2NaBr(aq)  2NaF(aq) + Br2(aq)
F has replaced Br in NaBr
Fluorine is more active than bromine (see following slide)
Elemental fluorine has become fluoride
Bromide has become elemental bromine

58. Activity Series for Halogens
Most Active
Fluorine
Chlorine
Bromine
Iodine
Least Active

59. Single-Replacement Reactions
Replacement of one halogen by another halogen
Br2(l) + MgCl2(aq)  ???
Is bromine more active than chlorine? No
Br2(l) + MgCl2(aq)  NR

60. Activity Series for Halogens
Most Active
Fluorine
Chlorine
Bromine
Iodine
Least Active

61. Practice Problems 21-24, page 295 Problem 88, page 313
Problems 13-15, pages

62. Single-Replacement Reactions Summary
Active metal replaces H atom in H2O
Metal replaces H atom in HX(aq)
One metal (solid) replaces 2d metal in an aqueous ionic compound
Replacement of nonmetal in compound by an elemental nonmetal (halogens)

63. Double-Replacement Reactions
Exchange of ions between two compounds in aqueous solution
In example shown above, hydroxide ion and chloride ion have exchanged

64. Double-Replacement Reactions
One of the products is always:
A gas
Water
Precipitate – solid that comes out of solution
Unlike single replacement reactions, no elemental forms produced – elements start as part of compound and also end that way

65. Double-Replacement Reactions
KCN(aq) + HBr(aq)  KBr(aq) + HCN(g)
Gaseous Product
Ca(OH)2(aq) + 2HCl(aq)  CaCl2(aq) + 2H2O(l)
Water is product
2NaOH(aq) + CuCl2(aq)  2NaCl(aq) + Cu(OH)2(s)
Precipitate is product

66. Double-Replacement Reactions
Steps to determine balanced equation
See table 9.3, p 297
Step 1 – Write reactants in skeleton
Al(NO3)3(aq) Na2CO3(aq)
Step 2 – Identify anions and cations
Al+3 NO Na+ CO32-
Step 3 – Swap cations
Na+ NO Al+3 CO32-

67. Double-Replacement Reactions
Step 4 – Write formulas for products
Al2(CO3)3(s) NaNO3(aq)
Step 5 – Write complete equation
Al(NO3)3(aq) + Na2 CO3(aq)  Al2(CO3)3(s) + NaNO3(aq)
Step 6 – Balance equation
2Al(NO3)3(aq) + 3Na2 CO3(aq)  Al2(CO3)3(s) + 6NaNO3(aq)

68. Practice
Problems 25-28, page 297

69. Predicting Products of Chemical Reactions Table 9.4, p 298

70. Chapter 9 – Chemical Reactions
9.1 Reactions and Equations
9.2 Classifying Chemical Reactions
9.3 Reactions in Aqueous Solutions

71. Section 9.3 Reactions in Aqueous Solutions
Double-replacement reactions occur between substances in aqueous solutions and produce precipitates, water, or gases.
Describe aqueous solutions.
Write complete ionic and net ionic equations for chemical reactions in aqueous solutions.
Predict whether reactions in aqueous solutions will produce a precipitate, water, or a gas.

72. Section 9.3 Reactions in Aqueous Solutions
(cont.)
Know the solubility rules for common cations and anions and use them to predict the occurrence of and formula for a precipitate.

73. Section 9.3 Reactions in Aqueous Solutions
Key Concepts
In aqueous solutions, the solvent is always water. There are many possible solutes. Many molecular compounds form ions when they dissolve in water. When some ionic compounds dissolve in water, their ions separate.
When two aqueous solutions that contain ions as solutes are combined, the ions might react with one another. The solvent molecules do not usually react.
Reactions that occur in aqueous solutions are double-replacement reactions.

74. Aqueous Solutions Water is solvent
Water-soluble substance is the solute
Solutes can be:
Molecules (covalent) that remain intact
Sugar, ethanol
Compounds that form ions
Ionic compounds – NaCl
Covalent compounds – HCl, NH3

75. Aqueous Solutions - Dissociation
NaCl(s)  Na+(aq) + Cl-(aq)
NaOH(s)  Na+(aq) + OH-(aq)
Ionic compound dissociates into solvated (hydrated) ions that can separate from each other
All soluble ionic compounds do this
Insoluble ionic compounds do not form ions – otherwise, they would dissolve!

76. Hydration Process for Ionic NaCl
Hydrated Ions
Na Ions
Cl Ions
H2O Molecules
Crystal Lattice

77. HCl(g)  H+(aq) + Cl-(aq)
Aqueous Solutions
HCl(g)  H+(aq) + Cl-(aq)
Despite being primarily a covalent compound, HCl in water dissociates into individual ions that can separate

78. Net Ionic Equations & Aqueous Reactions
Three types of double replacement reaction have a net ionic equation
Reactions that produce:
A solid (precipitate)
Water (neutralization)
A gas

79. Reactions Forming a Precipitate Solubility Rules
For certain common reactions, must learn to predict when a precipitate will form in aqueous solution
Possible to do if know solubility rules – these are not in this chapter but are listed on page 974 (Table R-8)
Are responsible for those rules on the following slide [need to know all for AP]

80. Ionic Compounds: Solubility in H2O
Compounds of alkali metal ions and ammonium ions are soluble
NaCl(aq), LiOH(aq), (NH4)2CO3(aq)
Nitrates and bicarbonates compounds are soluble
Mg(NO3)2(aq), NaHCO3(aq)
Carbonates, phosphates, hydroxides, and oxides are insoluble, except for compounds with alkali metals or ammonium ion
CaCO3(s), AlPO4(s) but Na2CO3(aq)

81. Ionic Compounds: Solubility in H2O
Soluble: Alkali metal, ammonium, nitrate, bicarbonate
Insoluble: Carbonate, phosphate, hydroxide, oxide (except with cations above)
Na2S ?
NH4OH ?
Ca3(PO4)2 ?
K2CO3 ?
MgCO3 ?
(NH4)3PO4 ?
Ba(OH)2 ?
Soluble (alkali)
Soluble (ammonium)
Insoluble (phosphate)
Insoluble (carbonate)
Insoluble (hydroxide)

82. Precipitation Reactions
2NaOH(aq) + CuCl2(aq) 
2NaCl(aq) + Cu(OH)2(s)
Solubility rules tell you NaCl must be soluble (alkali metal) and Cu(OH)2 must be insoluble (hydroxide, not alkali or NH4+)
However, if you forgot hydroxide rule but were told a precipitate forms, know the hydroxide must be the insoluble compound because NaCl can’t be (rule for alkali metal)

83. Precipitation Reactions
2NaOH(aq) + CuCl2(aq) 
2NaCl(aq) + Cu(OH)2(s)
Equation does not reveal ionic states
Use complete ionic equation to do this
2Na+(aq) + 2OH-(aq) + Cu2+(aq) + 2Cl-(aq) 
2Na+(aq) + 2Cl-(aq) + Cu(OH)2(s)
This does not dissociate (it is not soluble)

84. Net Ionic Equations 2Na+(aq) + 2OH-(aq) + Cu2+(aq) + 2Cl-(aq) 
2Na+(aq) + 2Cl-(aq) + Cu(OH)2(s)
Sodium and chloride ions are spectator ions (don’t participate in reaction)
Dropping them leads to net ionic equation
2OH-(aq) + Cu2+(aq)  Cu(OH)2(s)

85. Practice Problems 35-39, Page 302 Problems 99, 100 page 313

86. Reactions That Form Water
H+ and OH- ions can combine to form H2O (covalent)
H+(aq) + OH-(aq)  H2O(l)
These reactions also known as acid-base or neutralization reactions
No new phase formed
Already in aqueous phase

87. Reactions That Form Water
HnX type compounds are acids that dissociate into nH+(aq) ions and Xn-(aq) ions
X=SO4 sulfuric acid n=2
X=Cl hydrochloric acid n=1
X=PO4 phosphoric acid n=3
X=NO3 nitric acid n=1
X=CO3 carbonic acid n=2

88. Reactions That Form Water
Me(OH)m type compounds are bases that dissociate into mOH-(aq) ions (hydroxide) and Me+m(aq) metal ions
Me=Na sodium hydroxide m=1
Me=K potassium hydroxide m=1
Me=Ca calcium hydroxide m=2
Me=Mg magnesium hydroxide m=2

89. Reactions That Form Water
Equation
HBr(aq) + NaOH(aq)  H2O(l) + NaBr(aq)
Complete ionic equation
H+(aq) + Br-(aq) + Na+(aq) + OH-(aq) 
H2O(l) + Na+(aq) + Br-(aq)
Spectator ions are Na+(aq), Br-(aq)
Net ionic equation
H+(aq) + OH-(aq)  H2O(l)

90. Practice Problems 40-44, page 304 Problems 100-102, page 313

91. Reactions That Form Gases
Carbon dioxide CO2(g)
Hydrogen cyanide HCN(g)
Hydrogen sulfide H2S(g)
Carbon dioxide comes from the decomposition reaction of carbonic acid
H2CO3(aq)  CO2(g) + H2O(l)

92. Reactions That Form Gases
Equation
2HI(aq) + Li2S(aq)  H2S(g) + 2LiI(aq)
Complete Ionic Equation
2H+(aq) + 2I-(aq) + 2Li+(aq) + S2-(aq) 
H2S(g) + 2Li+(aq) + 2I-(aq)
Spectator ions are Li+(aq), I-(aq)
Net ionic equation
2H+(aq) + S2-(aq)  H2S(g)

93. Reactions That Form CO2
Occur when acid added to either a bicarbonate or a carbonate
Equation 1 – double replacement
HCl(aq) + NaHCO3(aq) 
H2CO3(aq) + NaCl(aq)
Equation 2 – decomposition
H2CO3(aq)  H2O(l) + CO2(g)
Can construct overall equation by combining individual equations

94. Double-Replacement Followed by Decomposition for CO2

95. Reactions That Form CO2 Overall equation
HCl(aq) + NaHCO3(aq) + H2CO3(aq) 
H2CO3(aq) + NaCl(aq) + H2O(l) + CO2(g)
Cancel substances appearing on both sides of the equation
Overall equation
HCl(aq) + NaHCO3(aq) 
NaCl(aq) + H2O(l) + CO2(g)

96. Reactions That Form CO2 Overall equation Complete ionic equation
HCl(aq) + NaHCO3(aq) 
NaCl(aq) + H2O(l) + CO2(g)
Complete ionic equation
H+(aq) + Cl-(aq) + Na+(aq) + HCO3-(aq) 
Na+(aq) + Cl-(aq) + H2O(l) + CO2(g)
Cancel to get net ionic equation
H+(aq) + HCO3-(aq)  H2O(l) + CO2(g)

97. Equation 1 – double replacement
CO2 from Carbonate
Equation 1 – double replacement
2HCl(aq) + Na2CO3(aq) 
H2CO3(aq) + 2NaCl(aq)
Equation 2 – decomposition
H2CO3(aq)  H2O(l) + CO2(g)
Can construct overall equation by combining individual equations

98. CO2 from Carbonate Overall equation
2HCl(aq) + Na2CO3(aq) + H2CO3(aq) 
H2CO3(aq) + 2NaCl(aq) + H2O(l) + CO2(g)
Cancel substances appearing on both sides of the equation
Overall equation
2HCl(aq) + Na2CO3(aq) 
2NaCl(aq) + H2O(l) + CO2(g)

99. CO2 from Carbonate Cancel to get net ionic equation Overall equation
2HCl(aq) + Na2CO3(aq) 
2NaCl(aq) + H2O(l) + CO2(g)
Complete ionic equation
2H+(aq) + 2Cl-(aq) + 2Na+(aq) + CO3-(aq) 
2Na+(aq) + 2Cl-(aq) + H2O(l) + CO2(g)
Cancel to get net ionic equation
2H+(aq) + CO32-(aq)  H2O(l) + CO2(g)

100. Practice
Problems , page 306