Chapter 11 “Chemical Reactions”

1. Chapter 11“Chemical Reactions”

2. Section 11.1p. 321Describing Chemical Reactions

3. All chemical reactions… • have two parts: • Reactants = substances you start with • Products = end with • reactants turn into products • Reactants ® Products

4. - Page 321 Products Reactants

5. In a chem rxn • Atoms not created or destroyed (Law of Conservation of Mass) • rxn described in a: #1.sentenceevery item is a word Copper reacts with chlorine to form copper (II) chloride. #2.word equation some symbols used Copper + chlorine ® copper (II) chloride

6. Symbols in equations? – Text page 323 • arrow (→) separates reactants from products (points to products) • Read as: “reacts to form” or yields • plus sign = “and” • (s) after formula = solid: Fe(s) • (g) = gas: CO2(g) • (l) = liquid: H2O(l)

7. Symbols used in equations • (aq) after formula = dissolved in water, aqueous solution: NaCl(aq) is salt water solution • ­ used after product - indicates gas produced: H2↑ • ¯ used after product - indicates solid produced: PbI2↓

8. Symbols used in equations • double arrow indicates a reversible reaction (more later) • shows that heat supplied to rxn • indicates catalyst supplied (here, platinum is catalyst)

9. What is a catalyst? • substance that speeds up rxn, w/o being changed or used up in rxn • Enzymes - biological or protein catalysts in your body

10. #3. The Skeleton Equation • Uses formulas and symbols to describe rxn • but doesn’t indicate how many; means they’re NOT balanced • All chem equations are descriptionof rxn

11. Write a skeleton equation for: • Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas. • Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

12. Now, read these equations: Fe(s) + O2(g)® Fe2O3(s) Cu(s) + AgNO3(aq)® Ag(s) + Cu(NO3)2(aq) NO2(g) N2(g) + O2(g)

13. #4. Balanced Chemical Equations • Atoms can’t be created or destroyed in an ordinary reaction: • All atoms we start with we must end up with (balanced!) • balanced equation has same # of each element on both sides of equation

14. Rules for balancing: • Assemble correct formulas for all reactants and products, using “+” and “→” • Count # of atoms of each type on both sides • Balance elements one at a time by adding coefficients (the numbers in front) where you need more - save balancing the H and O until LAST! (hint: I prefer to save O until the very last) • Double-Check make sure balanced

15. Never • Never change subscript (only change coefficients) • changing subscript (formula) describes different chemical • H2O different than H2O2 • Never put coefficient in middle of formula; only in front 2NaCl is okay, but Na2Cl is not.

16. Practice Balancing Examples 2 2 • _AgNO3 + _Cu ® _Cu(NO3)2 + _Ag • _Mg + _N2® _Mg3N2 • _P + _O2® _P4O10 • _Na + _H2O ® _H2 + _NaOH • _CH4 + _O2® _CO2 + _H2O 3 5 4 2 2 2 2 2

17. Balancing Equations Balancing Chemical Reactions Mark Rosengarten – 8:21

18. Section 11.2p. 330 Types of Chemical Reactions

19. Types of Reactions • 5 major types. • predict the products • predict whether or not they will happen at all • How? We recognize them by their reactants

20. #1 - Combination Reactions • Combine = put together • 2 substances combine to make one cmpd (also called “synthesis”) • Ca + O2® CaO • SO3 + H2O ® H2SO4 • predict products, especially if reactants are 2 elements • Mg + N2® _______ Mg3N2 (symbols, charges, cross)

21. Complete and balance: • Ca + Cl2® • Fe + O2®(assume iron (II) oxide is the product) • Al + O2® • Remember first step…write correct formulas – you can still change subscripts at this point, but not while balancing! • Then balance by changing just coefficients only

22. #1 – Combination Reactions • Additional Notes: a) Some nonmetal oxidesreact with H2O - produces acid: SO2 + H2O  H2SO3 b) Some metallic oxidesreact with H2O- produces base: CaO + H2O  Ca(OH)2 (how “acid rain” forms)

23. #2 - Decomposition Reactions • decompose = fall apart • one reactant breaks apart into 2 or more elements or cmpds • NaCl Na + Cl2 • CaCO3 CaO + CO2 • Note: energy (heat, sunlight, electricity, etc.) usually required

24. #2 - Decomposition Reactions • predict products if binary cmpd (made of 2 elements) • It breaks apart into the elements: • H2O • HgO

25. #2 - Decomposition Reactions • If cmpd has > 2 elements you must be given one of products • other product from the missing pieces • NiCO3 CO2 + ___ • H2CO3(aq) ® CO2 + ___ heat

26. #3 - Single Replacement Reactions • One element replaces another (new dance partner) • Reactants must be an element & cmpd • Products will be a different element and different cmpd • Na + KCl ® K + NaCl • F2 + LiCl ® LiF + Cl2 (Cations switched) (Anions switched)

27. #3 Single Replacement Reactions • Metals replace other metals (they can also replace H) • K + AlN ® • Zn + HCl ® • Think of water as: HOH • Metals replace first H, then combines w/ hydroxide (OH). • Na + HOH ®

28. #3 Single Replacement Reactions • can even tell whether or not single replacement rxn will happen: • b/c some chemicals more “active” than others • More active replaces less active • list – p. 333 - Activity Series of Metals • Higher on list replaces lower

29. The “Activity Series” of Metals • Lithium • Potassium • Calcium • Sodium • Magnesium • Aluminum • Zinc • Chromium • Iron • Nickel • Lead • Hydrogen • Bismuth • Copper • Mercury • Silver • Platinum • Gold • Metals can replace other metals, if they are above metal trying to replace (i.e. Zn will replace Pb) Higher activity • Metals above H can replace H in acids. • Metals from Na upward can replace hydrogen in H2O Lower activity

30. The “Activity Series” of Halogens Higher Activity Halogens can replace other halogens in compounds, if they are above halogen they are replacing Fluorine Chlorine Bromine Iodine Lower Activity 2NaF(s) + Cl2(g) 2NaCl(s) + F2(g) ??? MgCl2(s) + Br2(g) No Reaction! ???

31. #3 Single Replacement Reactions Practice: • Fe + CuSO4® • Pb + KCl ® • Al + HCl ®

32. #4 - Double Replacement Reactions • Two things replace each other. • Reactants must be two ionic compounds, in aqueous solution • NaOH + FeCl3® • positive ions change place (dance partners) • NaOH + FeCl3® Fe+3OH- + Na+1Cl-1 = NaOH + FeCl3® Fe(OH)3 + NaCl

33. #4 - Double Replacement Reactions • Have certain “driving forces”, or reasons • onlyhappens if one product: a) doesn’t dissolve in water & forms solid (a “precipitate”), or b) is gas that bubbles out, or c) is molecular compound (usually water)

34. Complete and balance: • assume all of the following reactions actually take place: CaCl2 + NaOH ® CuCl2 + K2S ® KOH + Fe(NO3)3® (NH4)2SO4 + BaF2®

35. How to recognize which type? • Look at the reactants: E + E = Combination C = Decomposition E + C = Single replacement C + C = Double replacement

36. Practice Examples: • H2 + O2® • H2O ® • Zn + H2SO4® • HgO ® • KBr + Cl2® • AgNO3 + NaCl ® • Mg(OH)2 + H2SO3 ®

37. #5 – Combustion Reactions • Combustion means “add oxygen” • Normally, a cmpd composed of only C, H, (and maybe O) is reacted with oxygen – called “burning” • Complete combustion, products are CO2 and H2O • If incomplete, products are CO (or possibly just C) and H2O

38. Combustion Reaction Examples: • C4H10 + O2® (assume complete) • C4H10 + O2® (incomplete) • C6H12O6 + O2® (complete) • C8H8 + O2® (incomplete)

39. SUMMARY: An equation... • Describes a rxn • Must be balanced (follows the Law of Conservation of Mass) • only balance by changing coefficients • special symbols to indicate physical state, catalyst or energy required, etc.

40. Reactions • 5 major types • We can tell what type they are by looking at reactants • Single Replacement happens based on the Activity Series • Double Replacement happens if one product is: 1) a precipitate (an insoluble solid), 2) water (a molecular compound), or 3) a gas

41. Section 11.3p. 342Reactions in Aqueous Solution KMnO4 NiCl2 Co(NO3)2 CuSO4 K2CrO4 K2Cr2O7

42. Net Ionic Equations • Many reactions occur in water- that is, in aqueous solution • When dissolved in water, many ionic cmpds “dissociate”, or separate, into cations & anions • Now write ionic equation

43. Net Ionic Equations • Example (needs to be a double replacement reaction) AgNO3 + NaCl  AgCl + NaNO3 1. this is the full balanced equation 2. next, write it as ionic equationby splitting the cmpds into their ions: Ag1+ + NO31- + Na1+ + Cl1- AgCl + Na1+ + NO31- Note that the AgCl did not ionize, because it is a “precipitate” (Table 11.3 p. 344)

44. Net Ionic Equations 3. simplify by crossing out ions not directly involved (called spectator ions) Ag1+ + Cl1- AgCl This is called the net ionic equation Let’s talk about precipitates before we do some other examples

45. Predicting the Precipitate • Insoluble salt is a precipitate [note Figure 11.11, p.342 (AgCl)] • General solubility rules are found: • Table 11.3, p. 344 in textbook • Reference section - page R54 (Table B.9)

46. Let’s do some examples together of net ionic equations, starting with these reactants:

47. BaCl2 + AgNO3 →

48. NaCl + Ba(NO3)2 →

49. Pb(NO3)2(aq) + H2SO4(aq)