1. Chapter 7
Chemical Reactions

2. Objectives Explain what a chemical reaction is
Describe indications of chemical reactions
Use state symbols in reactions
Write balanced chemical equations

3. What is a Chemical Reaction?
One or more substances are converted into new substances
New substances
must be formed!!!

4. All Chemical Reactions
have two parts
Reactants - the substances you start with
Products- the substances you end up with
The reactants turn into the products.
Reactants ® Products

5. Indications of Reactions
Change in Heat
- Exothermic or Endothermic Reactions
Light
Production of a Gas – Does not need to smell
Formation of a Precipitate
-Precipitate is a solid that is produced as a result of a chemical rxn in solution

6. Precipitate!

7. In a Chemical Reaction The way atoms are joined is changed
Atoms aren’t created of destroyed.
Can be described several ways
In a sentence
Copper reacts with chlorine to form copper (II) chloride.
In a word equation
Copper + chlorine ® copper (II) chloride

8. Chemical Equation
Represents with symbols and formulas, the identities and relative amounts of the reactants and products in a chemical rxn.

9. Symbols In Equations
An arrow separates the reactants from the products
Read “reacts to form” or “yields”
The plus sign = “and”
(s) = solid
(g) = gas
(l) = liquid
(aq) = aqueous solution
Dissolved in water

10. Symbols In equations = reversible reactions Equilibrium (More later)
= Reaction is heated
= Catalyst is used (Copper)
Catalysts speed up reactions but are not consumed.
Enzymes are biological catalysts
= Specific Pressure
2.00 atmospheres (1 atm is normal) Δ Cu
2.0 atm

11. How Do Reactions Happen?
Simple View
Particles must collide

12. Reactions Continued Particles are moving (Kinetic Energy)
-Higher temperature means a higher speed
Particles collide
-Energy is absorbed by particles
-Bonds are broken
-New bonds are formed
-Energy is released

13. Diatomic Elements 7 elements ALWAYS exist in diatomic state
Diatomic = 2 atoms
H2 , N2 , O2 , F2 , Cl2 , Br2 , I2
Elements in –ogen and –ine
1 + 7 pattern on the periodic table

14. Diatomic Elements

15. Converting To Formula Equ.
You will often have to convert word equations to formula equations.
Determine the reactants and products
Covert the words to equations
Include any state symbols that are given
If no state are given don’t worry about them.

16. Converting To Formula Equ.
Sodium metal and chlorine gas react to form solid sodium chloride

17. Converting To Formula Equ.
A solution of hydrochloric acid and solid sodium carbonate react to form solid sodium chloride and gaseous carbon dioxide and water vapor.

18. Convert To a Sentence Fe(s) + O2(g) ® Fe2O3(s)
Solid Iron and gaseous oxygen yields solid iron (III) oxide

19. Convert To a Sentence Cu(s) + AgNO3(aq) ® Ag(s) + Cu(NO3)2(aq)
Solid Copper and a solution of silver nitrate yields solid silver and a solution of copper (II) nitrate

20. Balanced Equation
Law of Conservation of Mass states “Mass cannot be created or destroyed.”
Thus, atoms can’t be created or destroyed
So, a balanced equation has the same number of each element on both sides of the equation.
Balance equations with coefficients
Number in front of a formula (Multiplier)

21. Is This Equation Balanced?
H2 + I2  HI

22. NO WAY!
Hydrogen and Iodine – 2 reactant atoms only 1 product atom

23. Balance With Coefficients
If there are 2 HI molecules the equation is balanced H2 + I2  2HI

24. Don’t Change the Formula
You make a different compound!!!!!

25. Writing Balanced Equations
Write the correct formulas for all the reactants and products
Count the number of atoms of each type appearing on both sides
Balance the elements one at a time by adding coefficients (the numbers in front)
Check to make sure it is balanced.

26. Never Never change a subscript to balance an equation.
If you change the formula you are describing a different reaction.
H2O is a different compound than H2O2
Never put a coefficient in the middle of a formula
2 NaCl is okay, Na2Cl is not.

27. Examples
H2 + O2  H2O

28. Examples
Ca(NO)3 + NaI  CaI2 + NaNO3

29. Examples
C2H6 + O2  CO2 + H2O

30. Homework
p. 264 #23,25,26,28,30-33

31. Objectives Predict a reaction type Predict the products of a reaction
Use the activity series
Predict solubility of compounds

32. Predicting the Products
Types of Reactions
Predicting the Products

33. Types of Reactions There are millions of reactions.
Can’t remember them all
Fall into several categories.
We will learn 5 types.
Will be able to predict the products.
For some we will be able to predict whether they will happen at all.
Will recognize them by the reactants

34. #1 Synthesis Reactions Combine - put together
2 elements, or compounds combine to make one compound.
Ca +O2 ® CaO
SO3 + H2O ® H2SO4
We can predict the products if they are two elements.
Mg + N2 ® Mg3N2

35. #2 Decomposition Reactions
decompose = fall apart
one reactant falls apart into two or more elements or compounds.
NaCl ® Na + Cl2
CaCO3 ® CaO + CO2

36. #3 Single Replacement One element replaces another
Reactants must be an element and a compound.
Products will be a different element and a different compound.
Na + KCl ® K + NaCl
F2 + LiCl ® LiF + Cl2

37. #3 Single Replacement We can tell whether a reaction will happen
Some are more active than other
More active replaces less active
Higher on the list replaces lower.
If the element by itself is higher, it happens, in lower it doesn’t

38. Activity Series Halogens - F2 Cl2 Br2 I2 Lithium Potassium Calcium
Sodium
Magnesium
Aluminum
Zinc
Chromium
Iron
Nickel
Lead
Hydrogen
Bismuth
Copper
Mercury
Silver
Platinum
Gold
Activity Series
Halogens - F2
Cl2
Br2 I2

39. #3 Single Replacement
What does it mean that Au And Ag are on the bottom of the list?
Nonmetals can replace other nonmetals
Limited to F2 , Cl2 , Br2 , I2
The order of activity is that on the table.
Higher replaces lower.

40. Solubility Some compounds dissolve in water.
We say they are Soluble
Examples
Sodium Chloride, Potassium Nitrate
Some compounds do not dissolve
We say they are Insoluble
Form Precipitates

41. Solubility Rules
Solubility Rules are a general list that tells us what kind of compounds are soluble or insoluble.
Follow from beginning to end
Rule one has precedence over rule two

42. Solubility Rules All Acids are soluble
Most nitrate and acetate salts are soluble.
Most salts containing the alkali metal ions (Li+, Na+, K+, Cs+, Rb+) and the ammonium (NH4+) ion are soluble.
Most chloride, bromide and iodide salts are soluble. Exceptions are salts containing the ions Ag+, Pb2+, and Hg2+.
Most sulfate salts are soluble. Notable exceptions are BaSO4, PbSO4, HgSO4 and CaSO4.

43. Solubility Rules Most hydroxide salts are insoluble.
Most sulfide (S2-), carbonate (CO32-), chromate (CrO42-) and phosphate (PO43-) salts are insoluble

44. Solubility Rules M I Soluble? Potassium Bromide Iron (III) Sulfate
Yes
Iron (III) Sulfate
Calcium Phoshate No
Zinc Acetate

45. #4 Double Replacement Two things replace each other.
Reactants must be two ionic compounds or acids.
Usually in aqueous solution
NaOH + FeCl3 ® Fe(OH)3 + NaCl

46. #4 Double Replacement Will only happen if one of the products
doesn’t dissolve in water and forms a solid
or is a gas that bubbles out.
or is a covalent compound usually water.

47. Examples H2 + O2 ® H2O ® Zn + H2SO4 ® HgO ® KBr +Cl2 ® AgNO3 + NaCl ®
Mg(OH)2 + H2SO3 ®

48. Last Type
Combustion
A compound composed of only C H and maybe O is reacted with oxygen
If the combustion is complete, the products will be CO2 and H2O.
CH4 + O2  CO2 + 2H2O

49. How to recognize which type
Synthesis – Only one product
Decomposition – Only one reactant
Single replacement – Element and Compound as reactants
Double replacement – Two compounds
Combustion – Something reacting with oxygen

50. Homework
p. 264 #'s , 46 without net ionic

51. Objectives Predict the type of double replacement reaction
Write net ionic equations

52. More on Double Replacement
Three Main Types
Precipitation (A solid forms)
Gas Forming (A gas forms)
Neutralization
Reaction of Acid and Base

53. Precipitation One of the products will be a solid
Determined by solubility rules
That’s it.

54. Gas Forming You see bubbles! Usually from decomposition of
Carbonic Acid - H2CO3  CO2 + H2O
Sulfurous Acid - H2SO3  SO2 + H2O
When you get one of these as a product replace it with the gas and water.
Ex. HNO3 Na2CO3  H2CO3 + NaNO3
Products become CO2 + H2O + NaNO3

55. Neutralization Reaction of Acid and a Base Acid usually starts with H
Base usually has hydroxide OH-
Product is salt and water
Ex HF + KOH HOH +KF

56. Dissociation Ionic Compounds separating into ions
AgNO3(aq)+Na2S(aq)Ag+ +NO3- +2Na+ +S-2

57. Net Ionic Equations
Used to remove “unimportant” reactants and products
What’s “unimportant”
Whatever is not used to make precipitates, liquids, or gases
They are spectator ions

58. Net Ionic Rules Dissociate all soluble compounds
According to balanced equation
Leave all solids, liquids, and gasses
Cross out common terms (Same)
Rewrite as net ionic

59. Example
Solutions of lead (II) nitrate & hydrochloric acid are in beakers. Draw what in is in the beakers.

60. Example
Solutions of lead (II) nitrate & hydrochloric acid are mixed. Write the balanced equation

61. Example
Draw what is in the beaker when the two solutions are mixed.

62. Example
Solutions of lead (II) nitrate & hydrochloric acid are mixed. Write the net ionic equation

63. Example
Solutions of copper (II) nitrate and sodium sulfide are mixed. Write the net ionic equation.

64. Homework
p. 266 #'s 46,62ab,68,75,82,98