1. Chemical Reactions and Equations
Chapter 9

2. Chemical Reaction
A process in which one or more substances are converted into new substances with different chemical and physical properties.
Not reversible.

3. Chemical Reactions Examples:
Breathing
Burning gasoline
Baking
All chemical reactions involve two types of substances:
Reactant – a substance that enters into a chemical reaction
Product – a substance that is produced by a chemical reaction

4. Reasons for Reactions
The arrangement of electrons in an atom determines whether it will bond with other atoms and with which atoms it will bond.
An atom with a full set of valence electrons will not form bonds.
An atom with an incomplete set of valence electrons will bond.

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6. Chemical Equations
9-2

7. Chemical Equations
Chemical reactions are represented by sentences known as chemical equations.
A chemical equation identifies the reactants and products in a chemical reaction.

8. Chemical Equations
A chemical reaction is the process by which one or more substances are changed into one or more different substances
It can be represented by an equation.
It shows what changes have taken place
It also shows the relative amounts of the various elements and compounds that take part in these reactions

9. Chemical Equations
The starting substances in a chemical reaction are the reactants
The substances that are formed by the chemical reaction are the products
Reactant(g) + Reactants(g)  Products(g) + Products(g)

10. Chemical Equations
The letters in parentheses indicate the physical state of each substance involved.
(g) means gas
(l) means liquid
(s) means solid
(aq) means that it is dissolved in water

11. Word Equations Gives names of the reactants and names of products.
Example:
calcium + oxygen  calcium oxide
+ means “reacts with”
 means “yields”, indicates direction of reaction.

12. Formula Equations Writing a chemical equation:
First determine the symbols and formulas that describe the reactants and products.
Substitute into the word equation.
Example: calcium + oxygen  calcium oxide
Ca + O2  CaO

13. Practice Problems
Silver (I) nitrate reacts with copper to form copper (II) nitrate and silver.
Hydrogen peroxide (H2O2) decomposes to form water and oxygen gas.

14. Balancing Chemical Equations
Law of conservation of mass:
Matter is neither created nor destroyed.
Total mass of the reactants must be equal to the total mass of the products.
For mass to remain constant before and after a chemical reaction, the number of atoms of each element must be the same before and after a chemical reaction.

15. Balancing Chemical Equations
The same number of atoms of each element must appear on both sides of the arrow in a chemical equation.
Example:
If reactants have 4 hydrogen and 1 carbon, products must also have 4 hydrogen and 1 carbon.

16. Balancing Chemical Equations
Rules for writing balanced equations
1. Determine the reactants and the products
2. Assemble the parts of the chemical equation
H2(g) O2(g)  H2O(l)
Reactants Yields Products

17. Balancing Chemical Equations
3. Write a balanced equation
Balancing means showing an equal number of atoms for each element on both sides of the equation
Remember nothing is lost or gained!
Only Transferred
When balancing equations only change the coefficients
Never change the subscripts!!
Finally make sure that all the coefficients are in the lowest possible ratio.

18. Balancing Equations Ca + O2  CaO Not balanced.
2 oxygen on the reactants, 1 oxygen in the products.
*Equation cannot be balanced by changing subscripts*
Changing a subscript changes the identity of the substance.

19. Balancing Equations
To balance an equation correctly we need to use coefficients.
Whole numbers written before the formula for reactants and products.
Ca + O2  CaO – not balanced
2 Ca + O2  2 CaO – balanced

20. Balancing Equations Example: Replace words with symbols:
Methane + oxygen  carbon dioxide + water
Replace words with symbols:
CH O2  CO2 + H2O

21. Balancing Equations CH4 + O2  CO2 + H2O To balance:
Start with those elements that occur in only one substance on each side of the equation.
**Number of atoms found by multiplying the subscript by the coefficient.
Balance by trial and error.
Balanced equation should have coefficients in the lowest whole number ratio possible.
CH O2  CO H2O

22. Practice Problems Balance the following equation:
H2(g) + O2(g)  H2O(l)

23. Practice Problems 2H2(g) + O2(g)  2H2O(l) H = 4 H = 4 O = 2 O = 2
The equation is balanced

24. Practice Problems Balance the following equation:
C2H6O + O2  CO2 + H2O

25. Practice Problems C2H6O + 3O2  2CO2 + 3H2O C = 2 C = 2 H = 6 H = 6
The equation is balanced

26. Practice Problems Balance the following equations:
1. Zn + HCl ----> ZnCl2 + H2
2. Al + O > Al2O3
3. Al + CuSO > Al2(SO4)3 + Cu
4. Li + H2O > LiOH + H2

27. Practice Problems
Aluminum reacts with oxygen to produce aluminum oxide.
Sodium nitrate reacting with calcium chloride to produce sodium chloride and calcium nitrate
Dinitrogen pentoxide reacts with water to produce nitric acid (HNO3)

28. Writing Complete Chemical Equations
A complete reaction must be balanced and include the physical state of each reactant and product.
Gas (g), solid (s), liquid (l)
Example:
CH4 (g) + 2O2 (g)  CO2 (g) + 2H2O (g)

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30. Classifying Chemical Reactions
9-3

31. Type of Reactions Four types of chemical reactions.
Direct combination (Synthesis)
Decomposition
Single-replacement
Double-replacement

32. Synthesis
Two or more reactants come together to form a single product.
Synthesis reactions.
A + B  AB
Product is always more complex than either of the reactants.

33. Synthesis
Synthesis – two or more substances combine to form one new substance
It follows this general form:
Element or compound + Element or compound  Compound

34. Synthesis Examples of synthesis: NH3 + HCl  NH4Cl CaO + SiO2  CaSiO3
2H2 + O2  2H2O
2Na + Cl2  2 NaCl
CO2 + H2O  H2CO3

35. Decomposition Reactions
Something breaking down into smaller parts.
A reaction in which a single compound breaks into two or more smaller compounds or elements.
Identified by only one reactant:
AB  A + B
AB represents a compound
A and B represent elements or simpler compounds.

36. Decomposition Reactions
Decomposition – when a compound is breaking apart or decompose into simpler substances when energy is supplied
Energy may be supplied in the form of heat, light, mechanical shock, or electricity
It follows this general form:
Compound  Two or more elements or compounds

37. Decomposition Reactions
Examples of Decomposition:
2H2O2 → 2H2O + O2
CaCO3 → CaO + CO2
H2CO3 → H2O + CO2
2KClO3 → 2KCl + 3O2
2 H2O  2 H2 + O2
CaCO3  CaO + CO2

38. Single-Replacement Reaction
Single Replacement – one element displaces another in a compound.
It follows this general form:
A BC  AC + B
Element + Compound  Element + Compound
Example of single displacement:
2AgNO3(aq) + Zn(s) → 2Ag(s) + Zn(NO3)2(aq)

39. Single-Replacement Reaction

40. Single-Replacement Reaction
An uncombined element displaces an element that is part of a compound.
Reactants: one element and one compound
A + BX  AX + B
Generally ionic compounds.

41. Single-Replacement Reaction
Example:
Mg + CuSO4  MgSO4 +Cu
A more active element will replace a less active element.
Figure 9-19 – activity series:
Used to predict whether or not a single-replacement reaction will occur.
An element can replace any element that is below it.
Metals replace metals or hydrogen
Nonmetals replace nonmetals
Cl2 + 2 KI  ?
Cl2 + 2 KI  KCl + I2

42. Double-Replacement Reaction
Double Replacement – is similar to single displacement but it uses two compounds instead of one element
It follows this general form:
Compound + Compound  Compound + Compound
AB + CD → AD + CB

43. Double-Replacement Reaction
Atoms or ions from two different compounds replace each other.
Identifying feature: two compounds as reactants and two compounds as products.
AX + BY  AY + BX
Example:
CaCO3 + 2HCl  CaCl2 + H2CO3

44. Double-Replacement Reaction
Examples of Double Displacement:
AgNO3 + HCl  AgCl + HNO3
Fe2O3 + HCl  FeCl3 + H2O

45. Double-Replacement Reaction
Conditions:
Most reactions will not occur unless the reactants are dissolved in water so that the ions can separate into ions.
Reactions are more likely to take place if one of the products is a molecular compound, a precipitate, or a gas.

46. Combustion
Combustion – occurs when a compound burns in air, it is actually reacting with the oxygen in the air.
The products of the oxidation of the hydrocarbon under normal conditions are carbon dioxide and water vapor.
Combustion reactions are refered to as oxidation reactions.
They follow this general form:
Hydrocarbon + Oxygen  Carbon Dioxide + Water (MOST USED)
Hydrocarbon + Flourine  Carbon Flouride + Hydrogen Flouride

47. Combustion Examples of Combustion/Oxidation: CH4 + 2O2 → CO2 + 2H2O
CH2S + 6F2 → CF4 + 2HF + SF6
CH4 + 2O2 + 7N2 → CO2 + 2H2O + 7N2 + heat

48. Combustion Not all reactions take one of these five general forms.
There are other classes of reactions, but those we will talk about in a later chapter.

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