1. REACTIONS

2. Reactions  Chemical equation: equation that shows the rearrangement of atoms that occurs in a chemical reaction  Reactants: original substances  Products: resulting substances A + B Yields AB Reactants Products

3. Symbols You’ll See SymbolMeaning (s)Solid state (l)Liquid state (g)Gas state (aq)Aqueous solution  Yields + Two reactants or two products ⇄ Reversible reaction

4. Law of Conservation of Mass  Recall that mass cannot be created or destroyed.  In other words reactant mass = product mass 28.014 amu + 6.0474 amu 34.061 amu = 14.007 amu1.0079 amu

5. Four Indications of Chemical Change 1. Energy absorbed or released in the form of heat or light. 2. Change in color or odor. 3. Production of gas (bubbling). 4. Formation of precipitate (solid).

6. Balancing Chemical Equations  Balanced chemical equations have equal numbers of atoms of each element in reactant and products.  Never change subscripts!  Coefficient: number of particles of a substance (can change).  Subscript: number of atoms in a substance (does not change). Subscript Indicate number of atoms of an element in a substance Ex) Three atoms of H Changing changes chemical identity 2 NH 3 Coefficient Indicate the number of particles of a substance Ex) Two molecules of NH 3 Change to balance formulas 2 atoms N & 6 atoms H

7. Steps for Balancing Chemical Equations 1. Write skeleton equation. 2. Write all elements present. 3. Count and write number of atoms of each element on reactant side. 4. Repeat step 3 for product side. 5. Add coefficients to balance. 6. Recount atoms to check work. 7. Make sure that coefficients are in lowest whole number ratio. A 2 B + C 2  CB 2 + A 2 ABCABC 2 1 2 2 2 1 2 4 2 2 4 2 2 4 4 4 8 8 4

8. How to Reduce Error  Go in this order: 1. Metals 2. Polyatomic Ions 3. Nonmetals 4. Hydrogen and Oxygen last! H 2 S + AgNO 3  HNO 3 + Ag 2 S H S Ag NO 3 21112111 11211121 2 2 2 2 2 2 H 2 S + AgNO 3  HNO 3 + Ag 2 S

9. Day 2—Types of Reactions

10. Types of Chemical Reactions 1. Synthesis 2. Decomposition 3. Single Replacement 4. Double Replacement 5. Combustion

11. Synthesis Two or more reactants form one product. A + B  AB

12. Decomposition One reactant forms two or more products. AB  A + B

13. Single Replacement A more reactive element replaces a similar, but less reactive, element in a compound. A + BC  AC+ B To check if a reaction will occur, check your reactivity chart.

14. Double Replacement Two compounds exchange ions and form two new compounds. AB + CD  AD + CB

15. Combustion Substance reacts with oxygen and releases energy in the form of heat and light. C x H y + O 2  CO 2 + H 2 O

16. Day 3—Reactivity Series

17. Reactivity Series  Please turn to the page in your notebook with the activity series. This is the same page as the list of polyatomic ions.  The activity series shows a list of metals (and hydrogen). The top metals are highly reactive and will take the place of another metal in a bond.  The bottom metals are fairly unreactive and will usually be found by themselves. (Notice that gold is down by the bottom.)

18. Activity Series  We’ve learned about single replacement reactions. These can only occur if the metal that is being added to the compound is more active than the original.  So for example,  Al + AgNO 3  Ag + Al(NO 3 ) 3 will occur.  But  Cu + FeSO 4  CuSO 4 + Fe will not. Iron is more active than Copper, and will not give up its bond with sulfate.

19. Solubility  Another important table is the solubility chart. This shows whether a compound is soluble (or can be dissolved in water), or insoluble (and will form a precipitate).  This chart is important in double-replacement reactions.  When you show the double replacement products, you can determine what phase they will be in.

20. Day 4—Acid-Base Reactions

21. Acid-Base Reactions Salt: ionic compound composed of the cation of a base and anion of acid Acid Base Reactions: are also double replacement, and also called “neutralization” Acid + Base  Water + Salt AcidBaseSalt HCl(aq) NaOH(aq) NaCl Water (aq) + HOH(l ) +  Ex)

22. Types of Reactions  Synthesis A + B  AB  Decomposition AB  A + B  Single Replacement AB + C  A + CB  Double Replacement AB + CD  AD + CB  Combustion C x H y + O 2  CO 2 + H 2 O  Acid-Base Neutralization Hanion + cationOH  H 2 O + salt

23. Day 5—Redox Reactions

24. Oxidation-Reduction Reactions Oxidation: Loss of electrons Reduction: Gain of electrons Memory Aids: OIL RIG: Oxidation I Lose, Reduce I Gain LEO goes GER: Lose Electrons, Oxidize…Gain Electrons, Reduce

25. Oxidation Numbers  Oxidation Number: value used to represent the number of electrons transferred.  Identify “redox” reactions by observing changes in oxidation number.

26. Assigning Oxidation Numbers If a compound were composed of ions, oxidation numbers would be the charges. Ex) H 2 SO 4 +1–2 = 0 x = +6 (2 ∙ +1)+ x + (4 ∙ –2 )

27. General Rules  Neutral elements = 0  Monatomic ions = ionic charge  Oxygen (when not by self) = -2  Hydrogen with nonmetal = +1  Hydrogen with metal = -1  Sum of states = the chemical formula’s overall charge

28. General Rules  Ask, “What was oxidized? What was reduced?” 1. Assign oxidation numbers 2. Compare oxidation numbers on each side a) Oxidation number decreases = reduced, gained electrons b) Oxidation number increases = oxidized, lost electrons c) No change = oxidation didn’t occur Ex) Mg + H 2 SO 4  MgSO 4 + H 2 –2 +6 +10 0 –2 +2 +6 gained 1 e – lost 2 e – H was reduced Mg was oxidized S & O were not oxidized or reduced

29. Oxidation Reduction Reactions Oxidizing Agent: substance that causes oxidation (whatever was reduced). Reducing Agent: substance that causes reduction (whatever was oxidized). Ex) Net ionic equation: Mg + 2H +  Mg 2+ + H 2 H + was reduced Mg was oxidized H + was the oxidizing agent Mg was the reducing agent

30. What Are Redox Reactions? Always: combustion, single replacement, synthesis of ions from pure elements. Never: double replacement—the oxidation numbers will remain the same.