Ch. 9 Chemical Reactions & Equations

1. Ch. 9 Chemical Reactions & Equations SAVE PAPER AND INK!!! When you print out the notes on PowerPoint, print "Handouts" instead of "Slides" in the print setup. Also, turn off the backgrounds (Tools>Options>Print>UNcheck "Background Printing")! Zn I2 Product Zn + I2 Reactants

2. Introduction • Chemical reactions occur when bonds between atoms are formed or broken • Chemical reactions involve changes in matter, the making of new materials with new properties, and energy changes. • Symbols represent elements, • Formulas represent compounds, • Chemical equations represent chemical reactions

3. Combustion of MethaneCH4 + 2O2  CO2 + 2H2O Atoms are rearranged!

4. Chemical Equations Yields or produces! Their Job: Depict the kind of reactants and products and their relative amounts in a reaction. 4Al (s) + 3O2 (g) > 2Al2O3 (s) The numbers in the front of formulas are called COEFFICIENTS The letters (s), (g), and (l) are the physical states of compounds. What’s created Starting Materials

5. The charcoal used in a grill is basically carbon. The carbon reacts with oxygen to yield carbon dioxide.

6. Lavoisier, 1788 Word Equations: show the names of the reactants and the products This is the “Word Equation” for that reaction carbon + oxygen carbon dioxide A skeleton equation does NOT indicate the relative amounts of reactants and products (no coefficients!) The skeleton equation for that reaction is: C + O2 CO2 Chemical Equations must be balanced in order to conform to the Law of Conservation of Mass - same # & type of atoms on each side of the yield arrow.

7. Word Equation: Methane + Oxygen gas carbon dioxide + water Skeleton equation: CH4 + O2 CO2 + H2O

8. Symbols Used in Equations • Solid (s) • Liquid (l) • Gas (g) • Aqueous solution (aq) (dissolved in water) • Catalyst H2SO4 • Escaping gas () • Change of temperature () • Precipitate ( )

9. Balancing Equations • When balancing a chemical reaction you may add coefficients in front of the compounds to balance the reaction, butyou may not change the subscripts. • Changing the subscripts changes the compound. Subscripts are determined by the valence electrons (charges for ionic or sharing for covalent) • Never put a coefficient in the middle of a formula • 2NaClis okNa2Clis not.

10. Subscripts vs. Coefficients • The subscripts tell you how many atoms of a particular element are in a compound. The coefficient tells you about the quantity, or number, of molecules of the compound.

11. There are a number of ways to interpret balanced equations 2 H2(s) + O2(g) ---> 2 H2O(s) This equation means: 2 molecules H2 + 1 molecules O2 ---->2 molecules H2O(g)

12. Steps to Balancing Equations There are four basic steps to balancing a chemical equation. • Write the correct formula for the reactants and the products. DO NOT TRY TO BALANCE IT YET! You must write the correct formulas first. And most importantly, once you write them correctly DO NOT CHANGE THE FORMULAS! • Find the number of atoms for each element on the left side. Compare those against the number of the atoms of the same element on the right side. • Determine where to place coefficients in front of formulas so that the left side has the same number of atoms as the right side for EACH element in order to balance the equation. • Check your answer to see if: • The numbers of atoms on both sides of the equation are equal. • The coefficients are in the lowest possible whole number ratios. (reduced)

13. Some Suggestions to Help You Helpful Hints for balancing equations: • Take one element at a time. • Save pure elements for last • IF everything balances except for that last pure element, and there is no way to balance it with a whole number, double all the coefficients and try again. • (Shortcut) Polyatomic ions that appear on both sides of the equation should be balanced as independent units That is, don’t separate them into individual atoms!

14. Example Make a table to keep track of atoms H2 + O2 H2O Need another O on the product side

15. Example Place a coefficient of 2 in front of H2O H2 + O2 2H2O Changes the O

16. Example H2 + O2 2H2O 2 This also changes the H

17. Example H2 + O2 2H2O 4 2 Now we need twice as much H in the reactant

18. Example 2H2 + O2 2H2O 4 2 Add a coefficient of 2 in front of H2

19. Example 2H2 + O2 2H2O Your answer 4 4 2 Recount to check

20. Balancing Equations Sodium phosphate + iron (III) oxide  sodium oxide + iron (III) phosphate Na3PO4 + Fe2O3 ---> Na2O + FePO4 3 2 2 R L 2 1 1 1 6 2 2 3 3 1 2 3 6 2 Na PO4 Fe O

21. Balancing Equations __C3H8 + __O2 ----> __CO2 + __ H2O 3 5 4

22. EXAMPLE:Leave oxygen for last! Balance B first. Now balance H How many O’s are now on the right? Can you get that many on the left? 11 NO 5 2 __B4H10 + ___O2 ---->___ B2O3 + ___H2O 2 11 4 10 When this happens, try doubling everything! How many O’s on the right, now? Can we get that on the left? 22 YES!!!

23. Types of Reactions • There are millions of chemical reactions. The only way to be sure what your products will be is to carry them out in the lab! • Not very practical – or cost effective, however… • There are five types of chemical reactions we can make some predictions for: • Combination reactions (Synthesis) • Decomposition reactions • Single Replacement reactions • Double Replacement reactions • Combustion reactions • You need to be able to identify the type of reaction and predict the product(s) given the reactants

24. 1. Combination Reactions • Combination reactions occur when two or more substances combine to form a single compound. • (Sometimes these are called synthesis reactions.) reactant + reactant  1 product • Basically: A + B  AB • Example: 2H2 + O2  2H2O • Example: C+ O2  CO2

25. Combination Reactions (Synthesis) 2 Mg + O2 2 MgO Mg O Mg2+ O2- O O2- Mg2+ Mg General form: A + B  AB element or element or compound compound compound

26. Synthesis Reactions • example

27. Balancing Equations ___ Al(s) + ___ Br2(l) ---> 2AlBr3(s) 2 3

28. Practice • Predict the products. Write and balance the following synthesis reaction equations. • Sodium metal reacts with chlorine gas Na(s) + Cl2(g)  • Solid Magnesium reacts with fluorine gas Mg(s) + F2(g)  • Aluminum metal reacts with fluorine gas Al(s) + F2(g)  2 2 NaCl MgF2 2 3 2 AlF3

29. 2. Decomposition Reactions • Decomposition reactions occur when a single compound breaks up into two or more simpler substances • 1 Reactant  Product + Product • In general: AB  A + B • Example: 2 H2O  2H2 + O2 • Example: 2 HgO  2Hg + O2

30. Decomposition Reactions • Another view of a decomposition reaction:

31. Example: Decomposition of hydrogen peroxide. • Word equation: hydrogen peroxide  water + oxygen

32. Example: Decomposition of hydrogen peroxide. • Skeleton equation: H2O2  H2O + O2

33. Example: Decomposition of hydrogen peroxide. • Balanced Equation: 2H2O2  2H2O + O2

34. Example: Decomposition of hydrogen peroxide in the presence of a catalyst. Balanced equation showing • the catalyst (MnO2) • the state of the reactants and products: MnO2 2H2O2(l)  2H2O (l) + O2 (g)

35. Decomposition Types • These are hardest to predict products for! • There are many special cases, but we will not explore those in Chemistry. For Ex.: • Carbonates and chlorates are special case decomposition reactions that do not go to the elements. • Carbonates (CO32-) decompose to carbon dioxide and a metal oxide • Example: CaCO3  CO2 + CaO • Chlorates (ClO3-) decompose to oxygen gas and a metal chloride • Example: 2 Al(ClO3)3  2 AlCl3 + 9 O2 • I

36. Practice • Predict the products. Then, write and balance the following decomposition reaction equations: • Solid Lead (IV) oxide decomposes PbO2  • Aluminum nitride decomposes AlN  Pb + O2 2 2 Al + N2

37. 3. Single Replacement Reactions • Single Replacement Reactions occur when one element replaces another in a compound. • A metal can replace a metal (+) OR a nonmetal can replace a nonmetal (-). • element + compound element + compound A + BC  AC + B (if A is a metal)OR A + BC  BA + C (if A is a nonmetal) *(remember the cation always goes first!) *Use Activity Series of Metals to see if reaction works.

38. Single Replacement Reactions • Another view:

39. Single Replacement Reactions • Single Replacement Reaction Mg + CuSO4 MgSO4 + Cu General form: A + BC  AC + B

40. 2 Al + Fe2O3 2 Fe + Al2O3 Here is a mixture of aluminum powder and iron (III) oxide powder. Once ignited the aluminum rips the oxygen off of the iron oxide resulting in molten iron, which falls down to the pot of sand below. One time someone used wet sand, which was a mistake because the molten iron turned the water to steam instantly. This sprayed molten iron all over the front two rows of students. Write the balanced equation for this reaction.

41. No, Ni is below Na Yes, Li is above Zn Yes, Al is above Cu Yes, Fe is above Cu We have looked at several reactions: Such experiments reveal trends. The Activity Series of Metals ranks the relative reactivity of metals. Used predict if single displacement reactions will occur:metals near the top are most reactive and will displacing metals near the bottom. Q: Which of these will react? Fe + CuSO4  Ni + NaCl  Li + ZnCO3  Al + CuCl2  Cu + Fe2(SO4)3 NR (no reaction) Zn + Li2CO3 Cu + AlCl3 p. 295

42. Activity Series Metal Reactivity Ca Li Rb K Ba Ca Na Mg Al Mn Zn Cr Fe Ni Sn Pb H2 Cu Hg Ag Pt Au H is the only nonmetal listed. H2 may be displaced from acids or can be given off when a metal reacts with H2O. Halogen Reactivity F2 Cl2 Br2 I2 Foiled again – Aluminum loses to Calcium

43. Single Replacement Reactions • Write and balance the following single replacement reaction equation: • Zinc metal reacts with aqueous hydrochloric acid Zn(s) + HCl(aq) ZnCl2 + H2(g) Note: Zinc replaces the hydrogen ion in the reaction 2

44. Single Replacement Reactions • Sodium chloride solid reacts with fluorine gas NaCl(s) + F2(g)  NaF(s) + Cl2(g) Note that fluorine replaces chlorine in the compound • Aluminum metal reacts with aqueous copper (II) nitrate Al(s)+ Cu(NO3)2(aq) 2 2

45. Complete these reactions: Cu + Sn2S  Ca + Cu2SO4  Na + PbO  NR 2 Cu + CaSO4 2 Pb + Na2O

46. 4. Double Replacement Reactions • Double Replacement Reactions occur when a metal in one compound replaces a metal in another compound. compound + compound  compound + compound • AB + CD  AD + CB

47. Double replacement reactions usually occur between 2 ionic compounds in aqueous solution. They are usually driven by the formation of either a gas, a precipitate (insoluble solid), or a molecular compound (like water).

48. Double Replacement Reactions CaCO3 + 2 HCl  CaCl2 + H2CO3 Double-replacement reaction General form: AB + CD  AD + CB

49. Double Replacement Reactions • Think about it like “foil”ing in algebra, first and last ions go together + inside ions go together • Example: AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq) • Another example: K2SO4(aq) + Ba(NO3)2(aq)  KNO3(aq) + BaSO4(s) 2