3. Effects of chemical reactions: Chemical reactions rearrange atoms in the reactants to form new products. The identities and properties of the products are completely different from that of the reactants. Production of gases and color changes are signs of chemical reactions.

4. Energy and Reactions zEnergy must be ADDED to BREAK bonds. zEnergy is RELEASED when bonds are FORMED. zChemical energy is CONSERVED in chemical reactions.

5. Exo- vs. Endo- zEXOTHERMIC REACTIONS: release energy (More energy is released as the products form bonds than is absorbed to break the bonds in the reactants.) zENDOTHERMIC REACTIONS: absorb energy

6. Chemical Equations Chemical equations are used to represent or describe chemical reactions. For example when hydrogen H 2 burns, it reacts with oxygen, O 2, in the air to form water. We write the chemical equation for this reaction as follows: 2H 2 + O 2 —> 2H 2 O

7. Chemical Equations An equation shows… Formulas of reactants Formulas of products Molar ratios of all compounds in the reaction.

8. Chemical Equations We read the (+) sign as “reacts with” and the arrow (—>) as “produces” or “yields”. 2H 2 + O 2 —> 2H 2 O ReactantsProducts

9. To show physical states of each substance: (s) or  - solid (l) - liquid (g) or  - gas (aq) - aqueous aqueous means dissolved in water

10. To show physical states of each substance: Consider the reaction of iron with oxygen to form iron (III) oxide, or rust. Fe (s) + O 2 (g)  Fe 2 O 3 (s)(unbalanced)

11. Coefficients & Subscripts COEFFICIENTS: numbers in front of compound that represents the number of molecules/moles of that compound SUBSCRIPTS: small numbers that help define the compound. 2H 2 SO 4 Coefficient Subscript

12. H 2 OOne molecule of water 2H 2 OTwo molecules of water H 2 O 2 One molecule of Hydrogen Peroxide

13. During a chem. rxn.; atoms are rearranged (NOT created or destroyed!) Chemical equations must be balanced to show the relative amounts of all substances. Balanced means: each side of the equations has the same # of atoms of each element. CH 4 + O 2 —> H 2 O + CO 2 Unbalanced CH 4 + 2O 2 —> 2H 2 O + CO 2 Balanced

14. In order to balance… Write correct formulas for all reactants and products Reactants  Products Count the number of atoms of each element in reactants & products. Balance one at a time using coefficients. Check for balance Are the coefficients in the lowest possible ratio?

15. Balancing Equations NOTE: When balancing equations, you may change coefficients as much as you need to, but you may never change subscripts because you can’t change what substances are involved.

16. Fe (s) + O 2 (g)  Fe 2 O 3 (s)(unbalanced) 4Fe (s) + 3O 2 (g)  2Fe 2 O 3 (s)(balanced)

17. Sample Problem 1 Water is decomposed (broken down) to form the gaseous products hydrogen, H 2, and oxygen, O 2. Write the balanced equation for this reaction. H 2 O  H 2 + O 2 2H + 1O  2H + 2O O is not balanced 2H 2 O  2H 2 + O 2 4H + 2O  4H + 2O The equation is balanced!

18. Sample Problem 2 Chlorine gas, Cl 2, reacts with potassium bromide, KBr, to form potassium chloride and bromine, Br 2. Write the balanced equation for this reaction, Cl 2 + KBr  KCl + Br 2 2Cl + 1K + 1Br  1Cl + 1K +2Br Cl and Br are not balanced Cl 2 + 2KBr  2KCl + Br 2 2Cl + 2K + 2Br  2Cl + 2K +2Br The equation is balanced!

19. Balancing equations involves a great deal of “trial and error” at first, but there are some tricks…

20. For example….. Sodium metal reacts with water to produce sodium hydroxide and hydrogen gas. Na + H 2 O —> NaOH + H 2 Note that on the product side (right side) there are an odd number of hydrogens. On the reactant side (left side) there is an even number. This implies there must be an even coefficient in front of the NaOH. Lets start with 2

21. _Na + _H 2 O —> 2NaOH + _H 2 Now lets balance sodium; we need a 2 in front of the Na… 2Na + _H 2 O —> 2NaOH + _H 2 Now consider hydrogen… 2Na + 2H 2 O —> 2NaOH + H 2

22. Check to see if it balances… 2 Na on the left2 Na on the right 4 hydrogen 2 + 2 = 4 hydrogen2 oxygen  the equation is balanced.

23. Examples CuCl 2 (aq) + Al(s)  Cu(s) +AlCl 3 (aq) 3CuCl 2 (aq) + 2Al(s)  3Cu(s) +2AlCl 3 (aq) (3:2:3:2)

24. Examples Propane, C 3 H 8, burns in oxygen, O 2, to form carbon dioxide and water. C 3 H 8 + O 2  CO 2 + H 2 O Balance C – then H – then O C 3 H 8 + 5O 2  3CO 2 + 4H 2 O (1:5:3:4)

25. Examples Pentane, C 5 H 12, burns in oxygen, O 2, to form carbon dioxide and water. C 5 H 12 + O 2  CO 2 + H 2 O Balance C – then H – then O C 5 H 12 + 8O 2  5CO 2 + 6H 2 O (1:8:5:6)

26. Examples Silver nitrate reacts with copper to produce silver and copper (II) nitrate. AgNO 3 + Cu  Ag + Cu(NO 3 ) 2 2AgNO 3 + Cu  2Ag + Cu(NO 3 ) 2 (2:1:2:1)

27. Examples Phosphorus reacts with oxygen gas to produce diphosphorus pentoxide. P + O 2  P 2 O 5 4P + 5O 2  2P 2 O 5 (4:5:2)

28. Examples C 7 H 14 + O 2  CO 2 + H 2 O Balance C – then H – then O C 7 H 14 + 10½O 2  7CO 2 + 7H 2 O 2C 7 H 14 + 21O 2  14CO 2 + 14H 2 O (2:21:14:14)

30. Types of Chemical Reactions Synthesis / Combination Decomposition Single Replacement Double Replacement Combustion

31. Synthesis / Combination Reactions Definition: Reaction where two or more substances react to form a single substance. A + B  AB Examples: 2K (s) + Cl 2 (g)  2KCl (s) SO 2 (g) + H 2 O (l)  H 2 SO 3 (aq)

32. Decomposition Reactions Definition: Reaction where a single compound is broken down into two or more products. AB  A + B Examples: 2H 2 O (l)  2H 2 (g) + O 2 (g) CaCO 3  CaO + CO 2

33. Single-Replacement Reactions Definition: Reaction where atoms of one element replace atoms of a second element in a compound. XA + B  BA + X Note: A reactive metal will replace any metal listed below it in the activity series. Generally, nonmetal replacement is limited to the halogens. The activity of the halogens decreases as you go down Group 7A of the periodic table. See handout. Examples: 2AgNO 3 + Mg  Mg(NO 3 ) 2 +2Ag Mg+LiNO 3  no reaction

34. Li K Ca Na Mg Al Zn Fe Pb (H)* Cu Hg Ag Increasing Activity Any element will replace any element below it. Activity Series *Metals from Li to Na will replace H from acids and water; from Mg to Pb they will replace H from acids only

35. For Example… Ca + MgO  CaO + Mg The Ca will replace the Mg because Ca is more active than Mg. That is to say…Ca is above Mg on the activity list.

36. Double-Replacement Reactions Definition: Reaction that involves an exchange of positive ions between two compounds. XA + BY  BA + XY Note: These reactions generally take place between two ionic compounds in aqueous solution, and are often characterized by one of the products coming out of solution in some way. Examples: 2NaCN (aq) +H 2 SO 4 (aq)  2HCN (g) +Na 2 SO 4 (aq) Na 2 S (aq) +Cd(NO 3 ) 2 (aq)  CdS (s) +2NaNO 3 (aq)

37. Combustion Reactions Definition: Reaction where an element or a compound reacts with oxygen, often producing energy in the form of heat and light. Examples: CH 4 +2O 2  CO 2 +2H 2 O + heat + light 2Mg (s) +O 2 (g)  2MgO (s)

38. Combustion of Hydrocarbons If the reactant is a hydrocarbon, the products are always carbon dioxide and water. CH 4 + 2O 2  CO 2 + 2H 2 O

39. Ionic Equations When a soluble substance is dissolved in water, the substance often breaks into ions. This solution is said to be an aqueous solution. Pb(NO 3 ) 2 (aq)  Pb 2 + + 2NO 3 - NaI (aq)  Na + + I -

40. Ionic Equations Consider the reaction… Pb(NO 3 ) 2 (aq) + NaI (aq)  PbI 2 (s) + NaNO 3 (aq) What is really going on is… Pb 2+ + NO 3 - + Na + + I -  PbI 2 (s) + Na + + NO 3 - Note that the Na + ion and the NO 3 - ion are not reacting. They are said to be spectator ions.

41. Net Ionic Equations It is often useful to write an equation showing only the species that are actually reacting. This is called a net ionic equation. It does not show the spectator ions. Pb 2+ + NO 3 - + Na + + 2I -  PbI 2 (s) + Na + + NO 3 - becomes…. Pb 2+ + 2I -  PbI 2 (s)

42. From the types of rxns wkst…