1. Unit 7: Chemical Reactions
Chapter 6

2. Section 1: The Nature of Chemical Reactions Objectives
Recognize some signs that a chemical reaction may be taking place.
Explain chemical changes in terms of the structure and motion of atoms and molecules.
Describe the differences between endothermic and exothermic reactions.
Identify situations involving chemical energy.

3. What are chemical reactions?
As we know, there are two types of changes that matter can undergo: physical and chemical.
Recall that a physical change may change the appearance of a substance, but does not change its composition.
However, when a substance undergoes a chemical change, a new substance with new properties is formed.
A chemical reaction is a chemical change.

4. Signs of chemical reactions
Often, there are visible signs that a chemical reaction has taken place.
Things to look for include:
Production of gas bubbles (such as CO2).
The formation of a solid precipitate (heavier substance that sinks to the bottom of the substance)
The production of light or heat
Change of color as a result of heating

5. The parts of a chemical reaction can be split into two categories: the products and the reactants
The reactants are the substances that a reaction begins with
The products are the substances that are formed in the chemical reaction
The products and reactants contain the same types and quantity of atoms, but they have different arrangements
Remember: you can’t create or destroy matter, so the atoms you end up with have to come from the atoms you start with

6. Energy and Reactions
As you may remember from the previous units, compounds are different from mixtures in several ways. One of these is that you can’t just pull the pieces of a compound apart like you can with a mixture.
In fact, the only way to separate the parts of a compound is to add energy to break the bonds between the atoms.
This energy can be heat, light, sound, or electricity.

7. If a reaction releases energy, it is said to be exothermic.
Likewise, when a bond is formed, energy is released. Again, this energy can be in the form of light, heat, sound, or electricity.
If a reaction releases energy, it is said to be exothermic.
If a reaction takes in energy, it is endothermic.
EXOTHERMIC
ENDOTHERMIC

8. Section 2: Reaction Types Objectives
Distinguish among five general types of chemical reactions.
Predict the products of some reactions based on the reaction type.
Describe reactions that transfer or share electrons between molecules, atoms, or ions.

9. Types of reactions There are 5 types of chemical reactions:
Synthesis reactions
Decomposition reactions
Combustion reactions
Single-displacement reactions
Double-displacement reactions

10. Synthesis reactions
In a synthesis reaction, many small molecules join to form fewer larger molecules (polymers).
A + B  AB is the general form
2Na + Cl2  2NaCl is a common example

11. Decomposition reactions
A reaction in which large substances are broken apart into smaller atoms or molecules
AB  A + B is the general form
2H2O  2H2 + O2 is a common example

12. Combustion reactions
Require oxygen as a reactant, so at least one product of the reaction contains oxygen
This usually involves organic compounds, and heat is often released
There is no general form for this type of reaction, except that O2 is always a reactant
2CH4 + 4O2  2CO2 + 4H2O is a common example, where methane combines with oxygen in the air to form carbon dioxide and water

13. Single-displacement reactions
When the atoms of one element appear to move into a compound, and atoms of the other element appear to move out, it is a single-displacement
AX + B  BX + A is the general form
3CuCl2 + 2Al  2AlCl3 + 3Cu is a common example

14. Double-displacement reactions
Here, ions appear to be exchanged between compounds
AX + BY  AY + BX is the general form
Pb(NO3)2 + K2CrO4  PbCrO4 + 2KNO3 is a common example

15. Name That Reaction!
A) S8 + 8O2  8SO2 + heat B) 6CO2 + 6H2O  C6H12O6 + 6O2 C) 2NaHCO3  Na2CO3 + H2O + CO2 D) Zn + 2HCl  ZnCl2 + H2

16. A) synthesis B) double-displacement C) decomposition D) single-displacement

17. Section 3: Balancing Chemical Equations Objectives
Demonstrate how to balance chemical equations.
Interpret chemical equations to determine the relative number of moles of reactant needed and moles of products formed.
Explain how the law of definite proportions allows for predictions about reaction amounts.

18. Chemical Equations
A chemical equation is used to describe what happens during a chemical reaction.
It makes use of the element symbols and compound formulas that we’ve already learned.
The reactants are always on the left-hand side and the products are always on the right-hand side. The arrow () means “yields”.
Reactants  Products

19. Conservation of Mass
Remember that matter cannot be created or destroyed, according to the law of conservation of mass.
That means that we can’t make more atoms than we start with, and we can’t lost atoms along the way.
So, the number of atoms in the reactants must equal the number of atoms in the products.
Balancing equations helps us to do this.

20. How to balance chemical equations
A few things to keep in mind when trying to balance equations:
You cannot change the formula of the compound or molecule
You cannot change the subscripts in the compound or molecule
You can use coefficients (numbers placed in front of a compound or molecule to tell how many there are) to balance
A coefficient multiplies everything in that compound or molecule by the given number
Anything in parentheses acts like a single atom (remember polyatomic ions?) and if you multiply one part of it, you must multiply all of it

21. Start with: CH4 + O2  CO2 + H2O
Notice that on the reactant side, there is one C, two O, and four H
On the product side, there is one C, three O (two of them in the first compound and one in the second), and two H.
Therefore, this equation is not balanced.

22. To balance this equation, we must use coefficients
To balance this equation, we must use coefficients. Let’s start with the H atoms.
CH4 + O2  CO2 + 2H2O
The reactant side now has one C, four H and two O.
The product side now has one C, four H, and three O.
We’re almost balanced!

23. Now, to balance the O atoms, we must do one more thing:
CH4 + 2O2  CO2 + 2H2O
This gives us one C on each side of the arrow, four H on each side, and four O on each side.
Notice that the 2 in front of H2O multiplies both the H and the O.
Our equation is balanced!

24. A balanced equation demonstrates the law of conservation of mass because the combined masses of all the atoms on the reactant side now equals the masses of all the atoms on the product side
A balanced equation also tells us how many molecules of a substance are required for a reaction to take place. In the example from before, methane will only combust if two molecules of oxygen gas are present, and will always produce (or yield) one molecule of carbon dioxide and two molecules of water. That reaction will always happen in the same proportions.

25. Let’s try another! Mg + O2  MgO
Right now, there is one atom of Mg and two atoms of O on the reactant side, and only one atom of each on the product side.
Therefore, we must balance!

26. The next step is to decide where we need to place coefficients
The next step is to decide where we need to place coefficients. Remember that a coefficient goes in front of a compound or molecule. You cannot separate a compound, nor can you change the formula or subscripts.
__Mg + __O2  __MgO

27. By placing a 2 in front of the MgO, we now get:
__Mg + __O2  2MgO
This means we now have one Mg and two O to start with, and two Mg and two O to end with.
Not quite finished!

28. To balance out the O in our final step, we must use a coefficient of __. This will give us:
2Mg + __O2  2MgO

29. Now, let’s try a little trickier one.
Pb(NO3)2 + KI  KNO3 + PbI2
What’s different about this equation?
Let’s count atoms. On the reactant side, we have one Pb, two N, six O, one K, and one I. On the product side, we have one K, one N, three O, one Pb, and two I.
Definitely needs some work!

30. To continue, we need to pick an element and start balancing
To continue, we need to pick an element and start balancing. Let’s start with our I atoms.
__Pb(NO3)2 + 2KI  __KNO3 + __PbI2
By placing a 2 in front of the KI, we have doubled our number of K and I. That means we now have one Pb, two N, six O, two K, and two I as reactants, and one Pb, one N, three O, one K, and two I as products.
Almost there.

31. Now let’s balance our K atoms.
__Pb(NO3)2 + 2KI  2KNO3 + __PbI2
By placing a 2 in front of our KNO3 atom, we now have one Pb, two N, six O, two K, and two I as reactants, and one Pb, two N, six O, two K, and two I as products.
Success!

32. The Law of Definite Proportions
The law of definite proportions says that:
A compound always contains the same elements in the same proportions, regardless of how the compound is made or how much of the compound is formed.
This means that water will always have twice as many H as O, whether it’s made in a synthesis reaction, decomposition, combustion, or displacement, and regardless of what it’s reacting with.

33. Section 4: Rates of Change Objectives
Describe the factors affecting reaction rates.
Explain the effect a catalyst has on a chemical reaction.
Explain chemical equilibrium in terms of equal forward and reverse reaction rates.
Apply LeChatelier’s principle to predict the effect of changes in concentration, temperature, and pressure in an equilibrium process.

34. The conditions under which most chemical reactions take place are very specific. For many of them, they can only occur if the temperature is just right, or if there is the right amount of other substances present, or if the wind is blowing in the right direction.
Ok, maybe not that last one, but the others are true.

35. Other things that can affect reaction rate are surface area, concentration of the substance, pressure, size of the molecules and the presence of catalysts.
In general, higher temperature, surface area, concentration, and pressure make a reaction go faster.
Heavier molecules usually react more slowly, and catalysts (enzymes or chemicals that can speed up a reaction by reducing the amount of energy needed) make reactions occur more quickly.

36. Equilibrium Not all reactions are terminal
That means that sometimes, a reaction is not finished when the products are made. Sometimes, the reaction will start over backwards.
This is called an equilibrium system.
The reaction can go in either direction, depending on conditions such as temperature, amount of pressure, and amount of the substances involved.

37. LeChatelier’s Principle
LeChatelier’s Principle says that:
If a change is made to a system in chemical equilibrium, the equilibrium shifts to oppose the change until a new equilibrium is reached.
This principle can be used to control reactions because those conditions can be manipulated (such as increasing temperature or decreasing pressure) to cause the reaction to go in the direction you desire.