1. Chapter 10
Chemical Reactions

2. Objectives Recognize evidence of chemical change.
Represent chemical changes using equations.

3. Evidence of Chemical Change
What are some things you can look for to determine if a chemical change has taken place?

4. Signs of Chemical Reactions
There are five main signs that indicate a chemical reaction has taken place:
release
input
change in color
change in odor
production of new
gases or solid
input or release
of energy
difficult to reverse

5. Representing Chemical Reactions
We can use equations (like math) to show a reaction.
Reactants are on the left; products are on the right.
Instead of =, we write 

6. We use arrows because they show the direction of the reaction.
How would we read this?
reactant 1 + reactant 2  product 1 + product 2

7. Symbols are used to represent reactants and products.
They can be solid, liquids, gases, or aqueous (dissolved in water)
When writing equations, use s, l, g, or aq to show the state of the reactant/product.

8. Word Equations Iron is a solid Chlorine is a gas
The brown cloud is composed of solid particles of the product.
iron(s) + chlorine(g)  iron(III) chloride (s)

9. Skeleton Equations Written using only the chemical symbols
iron(s) + chlorine(g)  iron(III) chloride (s)
Fe(s) + Cl2(g)  FeCl3(s)
***Why is it Cl2?
It is very important that you understand why it is Cl2!!!

10. Practice Write the skeleton equation for:
hydrogen(g) + bromine(g)  hydrogen bromide(g)
carbon monoxide(g) + oxygen(g)  carbon dioxide(g)
potassium chlorate(s)  potassium chloride(s) + oxygen(g)
**Don’t forget about the diatomic molecules!!!

11. Chemical Equations What’s wrong with this skeleton equation?
Fe(s) + Cl2(g)  FeCl3(s)
Look at the number of each atom…
Hint: Conservation of mass….

12. Chemical equations show the number of each atom required in a reaction.
When one is balanced, there are equal numbers of atoms on the left and right sides.
Fe(s) + Cl2(g)  FeCl3(s)
Balanced: 2Fe(s) + 3Cl2(g)  2FeCl3(s)

13. Fe(s) + Cl2(g)  FeCl3(s)
2Fe(s) + 3Cl2(g)  2FeCl3(s)

14. Balancing Equations
You can only change the coefficients (number in front).
You can NEVER change the subscripts!!!
So, you ask, how could you, an attentive and hard working chemistry student, balance equations all by yourself?!?! ….

15. Steps to Balancing Equations
1. Write the skeleton equation (no coefficients).
H2(g) + Cl2(g)  HCl(g)
2. Count atoms in the reactants.
2-H 2-Cl
3. Count atoms in the products.
1-H 1-Cl
If they’re the same, you’re done.
These work for most, if not all. Also, teach ½ multiplier, then multiply everything by 2.

16. 4. Change coefficients to make the number of atoms on each side of the  equal.
H2(g) + Cl2(g)  2HCl(g)
5. Write coefficients in their lowest ratio. (Reduce)
6. Check your work.
2-H Cl  2-H, 2-Cl

17. Practice
Write the balanced equation in which sodium hydroxide and calcium bromide react to produce sodium bromide and solid calcium hydroxide. The reaction occurs in water.
1. Skeleton reaction
NaOH(aq) + CaBr2(aq) Ca(OH)2(s) + NaBr(aq)

18. 2. Count atoms on reactants side.
1-Na 1-O 1-H 1-Ca 2-Br
3. Count atoms on products side.
1-Na 2-O 2-H 1-Ca 1-Br
4. Adjust coefficients.
2NaOH(aq) + CaBr2(aq) Ca(OH)2(s) + 2NaBr(aq)

19. 5. Write coefficients in their lowest ratio.
-It’s already reduced.
6. Check your work.

20. In water, iron(III) chloride reacts with sodium hydroxide, producing solid iron(III) hydroxide and sodium chloride.
Liquid carbon disulfide reacts with oxygen gas, producing carbon dioxide gas and sulfur dioxide gas.
Solid zinc and aqueous hydrogen sulfate react to produce hydrogen gas and aqueous zinc sulfate.

21. Checkpoint:
List three types of evidence that a chemical reaction has occurred.
What’s different about a skeleton equation and a chemical equation?
Why must a chemical equation be balanced?
Why can’t you change the subscripts when balancing chemical equations?

22. Types of Reactions
Chemical reactions can be classified based on type. The types are:
Synthesis
Combustion
Decomposition
Single Replacement
Double Replacement

23. Synthesis Two (or more) substances react to produce a single product.
A + B  AB
Easy to spot, since there’s only one product.
Can be two compounds, two elements, or a compound and an element.

24. Synthesis of Magnesium Oxide
Magnesium and Oxygen yield Magnesium Oxide
2 Mg (s) + O2 (g)  2MgO (s)
Product is new and different from the original reactants.

25. Decomposition Reaction
One substance breaks down into two or more substances. (Opposite of Synthesis)
AB  A + B
In this reaction, H2O2 decomposes into water and oxygen gas, with KI as a catalyst.
2H2O2 (aq)  2H2O (l) + O2 (g)

26. Combustion
In all combustion reactions, oxygen combines with a substance and energy is released as heat and light.
Burning coal is a combustion reaction:
C(s) + O2 (g) CO2 (g)
This is how burning coal produces greenhouse gases.

27. When they burn in Oxygen, the products are always CO2 and water.
Hydrocarbons are prone to combustion reactions. They have Hydrogen and Carbon.
When they burn in Oxygen, the products are always CO2 and water.
CxHx + O2 (g)  CO2 + H2O
*Watch demonstration of combustion reaction

28. Single Replacement
One element replaces the atoms of the other substance.
A + BX  AX + B
Usually, a metal replaces a metal, like this:
CuCl2 + Mg  MgCl2 + Cu

29. Metals won’t always change places with each other, because some metals don’t react as readily.
This is called activity series.
Metals can replace anything below them, but nothing above.

30. Will these reactions occur?
Ag (s) + Cu(NO3)2  ??
Cu(s) + Ag(NO3)2  ??
F2 + NaBr  ??
Br2 + NaF  ??

31. Double Replacement Ions are exchanged between two compounds.
AX + BY  AY + BX
2NaOH(aq) + CuCl2(aq)  2NaCl(aq) + Cu(OH)2(s)
In this reaction, the copper produced is a solid. When a solid is produced, it’s called a precipitate.

32. Synthesis Combustion Decomposition Single Replacement
A + B  AB
Combustion
? + O2  Energy
Cx Hx + O2  CO2 + H2O
Decomposition
AB  A+ B
Single Replacement
AX + B  A + BX
Double Replacement
AX + BY  AY + BX

33. Practice WS: Balancing and Classifying Chemical Equations

34. 10.3: Reactions in Solutions
Let’s get the definitions out of the way:
Solution:
Homogeneous (same throughout) mixture
Solute:
The material that is dissolved
Solvent:
The most plentiful substance; usually water

35. Aqueous Solutions In aqueous solutions, water is ALWAYS the solvent.
There can be many different solutes, like sugar or ethanol, which are covalent molecules.
Other solutes might form ions in solution:

36. Since the ions can move around, they can react.
When ionic compounds dissolve in water, the ions can separate and move around.
The equation might look like this:
HCl(g)  H+(aq) + Cl-(aq)
Since the ions can move around, they can react.

37. These are always double-replacement reactions.
When two solutions are mixed, and the ions interact, new products are formed.
The water – or solvent – does not react.
These are always double-replacement reactions.
Three possible products:
Precipitate, water, or gas.

38. Reactions forming precipitates
Consider this reaction:
2NaOH(aq) + CuCl2(aq) 2 NaCl(aq) + Cu(OH)2(s)
Both reactants are ionic compounds, so when they are in solution, what happens to them?

39. 2NaOH(aq) + CuCl2(aq) 2 NaCl(aq) + Cu(OH)2(s)
…the ions are free to move around and interact.
The equation could be represented by the ions only, like this:
2Na+(aq) + 2OH-(aq) + Cu2+(aq) + 2Cl-(aq) 
2Na+(aq) + 2Cl-(aq) + Cu(OH)(s)
This is a complete ionic equation: it shows all of the ions.

40. 2Na+(aq) + 2OH-(aq) + Cu2+(aq) + 2Cl-(aq) 
If some of the ions don’t change forms, then they are called spectator ions.
We usually don’t write these in ionic equations.
The resulting reaction is the “net ionic equation”.
2Na+(aq) + 2OH-(aq) + Cu2+(aq) + 2Cl-(aq) 
2Na+(aq) + 2Cl-(aq) + Cu(OH)(s)
Which are the spectator ions? Cross them out.

41. The resulting, net ionic equation is:
2OH-(aq) + Cu2+(aq)  Cu(OH)(s)

42. Practice Problems
Write the balanced equation, complete ionic equation, and net ionic equation for the following:
Aqueous solutions of barium nitrate and sodium carbonate form the precipitate barium carbonate.
Ba(NO3)2(aq) + Na2CO3(aq)  BaCO3(s)+ NaNO3(aq)
Balance the equation.

43. Write the complete ionic equation.
Ba2+(aq) + NO3-(aq) + 2Na+(aq) +CO32-(aq) 
BaCO3(s) + 2Na+(aq) + 2NO3-(aq)
Write the net ionic equation. (Cross out spectator ions.)
Ba2+(aq) +CO32-(aq)  BaCO3(s)

44. Write: Balanced, Complete Ionic, Net Ionic Equations
KI(aq) + AgNO3(aq)  KNO3(aq) + AgI(s)
AlCl3(aq) + NaOH(aq)  Al(OH)3(s) + NaCl(aq)

45. Reactions that form water
Sometimes, the result of a double- replacement reaction in solution is simply water.
HBr(aq) + NaOH(aq)  H2O(l) + NaBr (aq)
**Write the complete ionic equation, then write the net ionic equation.

46. H+(aq) + O2-(aq)  H2O(l)
Write the balanced, complete ionic, and net ionic equations for:
H2SO4(aq) + KOH(aq)  H2O(l) + K2SO4(aq)

47. Reactions that form gases
Some double-replacement reactions that occur in water may produce gases. This is often evident because of bubbles.
A reaction of this type looks like this:
HCl(aq) + Na2S(aq)  H2S(g) + NaCl(aq)
For the above, write the balanced, complete ionic, and net ionic equations.

48. Checkpoint Describe an aqueous solution.
Distinguish between a complete ionic equation and a net ionic equation.
What are three common types of products formed from the results of aqueous reactions?
Elaborate: Why would reactions that occur in solutions probably not occur otherwise?