1. Chemical Equations & Reactions
Chemistry 6.0

3. I. Chemical Reactions Color change Formation of a precipitate, ppt
Definition: a process by which 1 or more substances, called reactants, are changed into 1 or more substances, called products, with different physical & chemical properties.
Evidence of a Chemical Reaction
Color change
Formation of a precipitate, ppt
Release of a gas
Energy change – heat, light, sound
Odor change
Reactions are started by the addition of energy

5. II. Chemical Equation Form Reactant + Reactant  Product + Product
Symbols: (s), (l), (g), (aq) NR

6. Interpretation of a Balanced Equation
2Mg(s) + O2(g)  2MgO(s)
2 atoms of solid magnesium react with
1 molecule of oxygen gas to form
2 formula units of solid magnesium oxide OR
2 moles of solid magnesium react with
1 moles of oxygen gas to form
2 moles of solid magnesium oxide

7. B. Energy & Chemical Equations
Exothermic reactions – release energy; energy a product
H2O(g)  H2O(l) kJ
Endothermic reactions – absorbs energy; energy a reactant
H2O(s) kJ  H2O(l)

8. C. Characteristics of A Balanced Chemical Equations
The equation must represent known facts. All substances have been identified.
The equation must contain the correct symbols and/or formulas for the reactants and products
Can be either a word equation or a formula equation
The law of conservation of mass must be satisfied. This provides the basis for balancing chemical equations. 1st formulated by Antoine Lavoisier
TOTAL MASS REACTANTS = TOTAL MASS PRODUCTS
Number of atoms of EACH element is the SAME on both sides of the equation.

10. D. Balancing Chemical Equations
Balance using coefficients after correct formulas are written.
Coefficients are usually the smallest whole number – required when interpreted at the molecular level
Balance atoms one at a time
Balance the atoms that are combined and appear only once on each side.
Balance polyatomics that appear on both sides
Balance H and O atoms last
NEVER CHANGE SUBSCRIPTS!!!
**Count atoms to be sure that the
equation is balanced**

11. BALANCING Examples sodium + chlorine  sodium chloride
CH4 (g) + O2 (g)  CO2 (g) + H2O(l)
K(s) + H2O(l)  KOH(aq) + H2(g)
AgNO3(aq) + Cu(s)  Cu(NO3)2(aq) + Ag(s)

13. Synthesis Reactions
Two or more substances combine to form a more complex product.
A + B → AB (only ONE PRODUCT)
A.K.A. Direct Combination Reactions, or composition reactions
Ex. Fe + S → FeS
Ex. CaO + H2O → Ca(OH)2

14. Synthesis Reaction

16. Sodium Metal plus Chlorine Gas Video
2 Na + Cl2  2 NaCl
Synthesis Reaction

17. Decomposition Reactions
Single Reactant breaks down to into a simpler substance.
AB → A + B (only ONE REACTANT)
The opposite of a synthesis reaction.
Ex. 2HgO → 2Hg + O2
Ex. CaCO3 → CaO + CO2

19. Single Replacement Reaction
Atoms of one element replace atoms of another element in a compound.
A + BX → AX + B
A more active element will replace a less active element. (See activity series)
Ex. Fe + CuSO4 → FeSO4 + Cu
Ex. Mg + CuSO4 → MgSO4 + Cu

21. Double-Replacement Reactions
Atoms or ions from 2 different compounds replace each other.
AX + BY → AY + BX
Ex. CaCO HCl → CaCl2 + H2CO3

24. Combustion Reactions
One substance reacts with oxygen to produce oxide compounds.
Occurs when burning.
Combustion reactions are often classified as synthesis reactions.
These reactions are usually exothermic, releasing a large amount of energy as light, heat, or sound.

25. Combustion Reactions (cont.)
When a hydrocarbon is involved in a combustion reaction, H2O and CO2 are the products.
Ex. CH O2 → CO2 + 2H2O kJ
Ex. S + O2 → SO2

26. Combustion Reaction

27. 5 Types of Chemical Reactions Video

28. III. Classifying Chemical Reactions
Pattern for prediction based
on the kind of reactants
Combustion or Burning – complete combustion always produces carbon dioxide and water!
Hydrocarbons
CxHy + O2  CO2 + H2O
Alcohols
CxHyOH + O2  CO2 + H2O
Sugars
C6H12O6 + O2  CO2 + H2O
C12H22O O2  CO2 + H2O

29. B. Synthesis or Composition
2/more reactants  1 product
Element + Element  Compound
A + B  AB
2 Na + Cl2  2 NaCl
4 Al O2  2 Al2O3

30. 2. Compound + Compound  Compound
EXAMPLE 1: metal oxide + carbon dioxide  a carbonate CaO + CO2  CaCO3 EXAMPLE 2: metal oxide + water  a base (hydroxide) Na2O + H2O  2 NaOH EXAMPLE 3: metal oxide + sulfur trioxide  metal sulfate Na2O + SO3  Na2SO4 EXAMPLE 4: metal oxide + sulfur dioxide  metal sulfite Na2O + SO2  Na2SO3 EXAMPLE 5: nonmetal oxide + water  an acid SO3 + H2O  H2SO4

31. Decomposition - Binary Compounds
1. Binary Compound  2 elements AB  A + B 2 H2O  2 H2 + O2 2 HgO  2 Hg + O2

32. Decomposition - Ternary Compounds
2. Ternary Compound  Compound + Element/Compound
a. metal carbonate  metal oxide + carbon dioxide
CaCO3  CaO + CO2
b. metal hydroxide  metal oxide + water
(Except Group IA metals)
Mg(OH)2  MgO H2O
c. metal chlorate  metal chloride + oxygen
2KClO3  2KCl O2
d. metal nitrate  metal nitrite + oxygen
2NaNO3  2NaNO O2
e. acids  nonmetal oxide + water (HINT: MUST USE CHARGE OF H and O)
H2CO3  CO2 + H2O
f. Other 2H2O2  2H2O + O2

33. Single Replacement or Single Displacement
Element + Compound  New Compound + New Element
Metals
A BC  AC + B
Active metals displace less active metals or hydrogen from their compounds in aqueous solution. Refer to the Activity Series.
a. Zn + CuSO4  ZnSO4 + Cu
b. metal + H2O  metal hydroxide + H2
Metals include the alkali metals and calcium.
2Na + 2H2O  2NaOH + H2

34. Reactivity or Activity of Metals
The reactivity of a metal is based on its ability to replace another in a compound.
If a single replacement reaction occurs, the metal that “cuts in” is MORE reactive than the one that was removed or replaced.
An activity series of metals is a listing that ranks metals according to their reactivity.
The most active metal is at the TOP of the list
The least active metal is at the BOTTOM of the list

35. Which is more active nickel or iron?
The ACTIVITY SERIES is listed below:
lithium
potassium
barium
strontium
calcium
sodium
magnesium
aluminum
manganese
zinc
iron
cadmium
cobalt
nickel
tin
lead
hydrogen
copper
silver
mercury
gold
The most active metal is LITHIUM
The least active metal is GOLD
Which is more active nickel or iron?
IRON

36. 3CuCl2 + 2Al  2AlCl3 + 3Cu

37. Single Replacement or Single Displacement
Element + Compound  New Compound + New Element
Metals
A BC  AC + B
Active metals displace less active metals or hydrogen from their compounds in aqueous solution. Refer to the Activity Series.
a. Zn + CuSO4  ZnSO4 + Cu
b. metal + H2O  metal hydroxide + H2
Metals include the alkali metals and calcium.
2Na + 2H2O  2NaOH + H2

38. Single Replacement or Single Displacement
2. Nonmetals D + EF  ED + F Cl2 + 2NaBr  2NaCl + Br2 Many nonmetals displace less active nonmetals from combination with a metal or other cation. Order of decreasing activity is F2  Cl2  Br2  I2

39. Double Replacement or Double Displacement or Metathesis:
Compound + Compound  New Compound + New Compound
AB CD  AD CB
AgNO NaCl  AgCl NaNO3
Special Notes:
If NH4OH is one of the products, it breaks down into NH3 and H2O.
If H2CO3 is one of the products, it breaks down into CO2 and H2O.

40. Double Replacement Reactions and Precipitates
If the reactants are both aqueous in a double replacement reaction, a precipitate may form.
Use solubility rules to determine the identity of the precipitate.
AgNO3 (aq) + NaCl (aq)  AgCl (s) NaNO3 (aq)

41. Exothermic Reactions Release heat into the surroundings
Heat is a product of the reaction
Combustion reactions are exothermic
C3H8 + 5O2 → 3CO H2O kJ

42. Endothermic Reactions
Heat is absorbed by the reactants and stored in the chemical bonds of the product.
Heat acts as a reactant.
C + H2O + 113kJ → CO + H2
Motorcycle Helmet

44. Molar Heat Capacity
The heat absorbed or released during a reaction depends on a difference in a quantity called enthalpy. (Total energy content of a sample.)
The symbol for enthalpy is H.
When reactions take place at standard temperature and pressure, q = H.
Stand. temp. = 25°C Stand. Pres. = 1 atm
Purest form of a substance = most stable form
Enthalpy change at STP denoted H°

45. Enthalpy Change ∆H = H products – H reactants
Type of reaction ∆H
H products
H reactants
Exothermic
Endothermic -
>
<
> +
<

46. Exothermic Reaction
Endothermic Reaction

47. What is the energy of the reactants? kJ
What is the energy of the reactants?
What is the energy of the products?
Is the forward reaction exothermic or endothermic?
What is the ΔH for the forward reaction?
What is the ΔH for the reverse reaction?
150 kJ
50 kJ
exothermic
-100 kJ
100 kJ

48. Hess’s Law
Hess’s Law states that if a series of reactions are added together, the enthalpy change of the net reaction will be the sum of the enthalpy changes of the individual steps.

49. Steps for using Hess’s Law
Identify the compounds
Locate the compounds on the periodic table
Write a reaction from the table.
Write the appropriate “sub equation.”
If needed, multiply equation and enthalpy change.
If you reverse the reaction, change sign of enthalpy change.
Add equations.
Add enthalpy changes.

50. Bond Energy
Bond energy is the strength of a chemical bond between atoms, expressed as the amount of energy required to break it apart. (Unit kJ/mol)
It is as if the bonded atoms were glued together: the stronger the glue is, the more energy would be needed to break them apart. A higher bond energy, therefore, means a stronger bond.
Ionic bonds are stronger than covalent bonds. Among covalent bonds, triple bonds are stronger than double bonds and double are stronger than single bonds.

51. Bond Energy for Carbon Bonds
Bond lengths (Å):
Single > double > triple
Bond Energy (kJ/mol)
Single < double < triple

52. Using Bond Energy in Calculations
On the REACTANT side, determine the number and type of each bond.
Determine the total amount of energy needed to break them. (endothermic – positive sign)
On the PRODUCT side, determine the number and type of each bond.
Determine the total amount of energy needed for bonds to form. (exothermic – negative sign)
Calculate the enthalpy change for the entire reaction.
∆H = reactants + products