1. Chapter 10 Chemical Quantities

2. All chemical reactions… Have two parts: Reactants - the substances you start with Products- the substances you end up with The reactants turn into the products. Reactants  Products A reaction can be described several ways: 1.In a sentence (every item is a word) Copper reacts with chlorine to form copper (II) chloride. 2. In a word equation (some symbols used) Copper + chlorine  copper (II) chloride

3. Symbols in equations The arrow separates the reactants  from the products Read as “reacts to form” or yields The plus + sign = “and” (s) after the formula = solid: AgCl (s) (g) after the formula = gas: CO 2(g) (l) after the formula = liquid: H 2 O (l) (aq) after the formula = dissolved in water, an aqueous solution: NaCl (aq) is a salt water solution

4. Symbols used in equations  used after a product indicates a gas has been produced: H 2 ↑  used after a product indicates a solid has been produced: PbI 2 ↓ indicates a reversible reaction shows that heat is supplied to the reaction is used to indicate a catalyst is supplied, in this case, platinum. Catalyst - A substance that speeds up a reaction, without being changed or used up by the reaction.

5. The Skeleton Equation 3.Uses formulas and symbols to describe a reaction but doesn’t indicate how many; this means they are NOT balanced All chemical equations are a description that describe reactions. 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.

6. Now, read these: Fe(s) + O 2 (g)  Fe 2 O 3 (s) Cu(s) + AgNO 3 (aq)  Ag(s) + Cu(NO 3 ) 2 (aq) NO 2 (g) N 2 (g) + O 2 (g)

7. Law of Conservation of M atter A natural law describing the fact that matter is neither created nor destroyed in any process The amount of matter that you start with has to equal to the amount of matter that you end with Atoms can’t be created or destroyed in an ordinary reaction: All the atoms we start with we must end up with A balanced equation has the same number of each element on both sides of the equation.

8. For Chemical Reactions This Means The amount of reactants has to equal the amount of products. Matter cannot be created or destroyed through a chemical reaction. Chemical equations have to be balanced.

9. Rules for balancing: 1.Assemble the correct formulas for all the reactants and products, use + and → 2.Count the number of atoms of each type appearing on both sides 3.Balance the elements one at a time by adding coefficients where needed (the numbers in front) - save balancing the H and O until LAST! 4.Check to make sure it is balanced.

10. Never change a subscript to balance an equation. –If you change the formula you are describing a different reaction. H 2 O is a different compound than H 2 O 2 Never put a coefficient in the middle of a formula 2NaCl is okay, but Na2Cl is not.

11. Balancing Chemical Equations Example: HCl + NaOH NaCl + H 2 OH=2 Cl=1 Na=1 O=1 The equation is balanced because the number of atoms in the reactants are equal to the number of atoms in the products.

12. Balancing Chemical Equations Example: H 2 + O 2 H 2 O H=2 O=2 H=2 O=1 H 2 + O 2 2 H 2 O H=2 O=2 H=4 O=2 2H 2 + O 2 2 H 2 O H=4 O=2 H=4 O=2

13. Balancing Chemical Equations Example: Cu + AgNO 3 Cu(NO 3 ) 2 + Ag Cu=1 Ag=1 N=1 O=3 Cu=1 Ag=1 N=2 O=6 Cu + 2AgNO 3 Cu(NO 3 ) 2 + Ag Cu=1 Ag=2 N=2 O=6 Cu=1 Ag=1 N=2 O=6 Cu + 2AgNO 3 Cu(NO 3 ) 2 + 2Ag Cu=1 Ag=2 N=2 O=6 Cu=1 Ag=2 N=2 O=6

14. NaHCO 3 + H 3 C 6 H 5 O 7 CO 2 + H 2 O + Na 3 C 6 H 5 O 7 Na=1 H=9 C=7 O=10 Na=3 H=7 C=7 O=10 3NaHCO 3 + H 3 C 6 H 5 O 7 CO 2 + H 2 O + Na 3 C 6 H 5 O 7 Na=3 H=11 C=9 O=16 Na=3 H=7 C=7 O=10 3NaHCO 3 + H 3 C 6 H 5 O 7 3CO 2 + H 2 O + Na 3 C 6 H 5 O 7 Na=3 H=11 C=9 O=16 Na=3 H=7 C=9 O=14 3NaHCO 3 + H 3 C 6 H 5 O 7 3CO 2 + 3H 2 O + Na 3 C 6 H 5 O 7 Na=3 H=11 C=9 O=16

15. Practice Balancing Examples _ AgNO 3 + _Cu  _Cu(NO 3 ) 2 + _Ag _Mg + _N 2  _Mg 3 N 2 _P + _O 2  _P 4 O 10 _Na + _H 2 O  _H 2 + _NaOH _CH 4 + _O 2  _CO 2 + _H 2 O

16. End of Section 11.1

17. Types of Reactions There are 5 major types of chemical reactions 1.Combination reaction 2.Decomposition reaction 3.Single Replacement reaction 4.Double Replacement reaction 5.Decomposition reaction Not all reactions fit into only one category Patterns of chemical reactions will help you predict the products of the reaction

18. Combination Reactions Combine = put together 2 substances combine to make one compound. Ca +O 2  CaO (2 elements form 1 compound) SO 3 + H 2 O  H 2 SO 4 (2 compounds form another) When 2 non metals react (or a transition metal and a non metal) in a combination reaction, often more than one product is possible. S (s) + O 2 (g)  SO 2 (g) 2S (s) + 3O 2 (g)  2SO 3 (g)

19. Complete and balance Ca + Cl 2  Fe + O 2  Al + O 2  Remember that the first step is to write the correct formulas – you can still change the subscripts at this point, but not later! Then balance by using the coefficients only

20. #2 - Decomposition Reactions decompose = fall apart one reactant breaks apart into two or more elements or compounds. NaCl Na + Cl 2 CaCO 3 CaO + CO 2 Note that energy (heat, sunlight, electricity, etc.) is usually required

21. #2 - Decomposition Reactions Can predict the products if it is a binary compound-Made up of only two elements –breaks apart into its elements: H 2 O HgO

22. #2 - Decomposition Reactions If the compound has more than two elements you must be given one of the products –The other product will be from the missing pieces NiCO 3 CO 2 + ___ H 2 CO 3 (aq)  CO 2 + ___ heat

23. #3 - Single Replacement One element replaces another Reactants must be an element and a compound. Products will be a different element and a different compound. Na + KCl  K + NaCl F 2 + LiCl  LiF + Cl 2

24. #3 Single Replacement Metals replace other metals (and they can also replace hydrogen) K + AlN  Zn + HCl  Think of water as: HOH –Metals replace one of the H, and then combine with the hydroxide. Na + HOH 

25. #3 Single Replacement We can even tell whether or not a single replacement reaction will happen: –Some chemicals are more “active” than others –More active replaces less active There is a list on page 333 - called the Activity Series of Metals Higher on the list replaces lower.

26. The Activity Series of the Metals Lithium Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Lead Hydrogen Bismuth Copper Mercury Silver Platinum Gold 1)Metals can replace other metals provided that they are above the metal that they are trying to replace. 2)Metals above hydrogen can replace hydrogen in acids. 3)Metals from sodium upward can replace hydrogen in water. Higher activity Lower activity

27. #3 Single Replacement Practice: Fe + CuSO 4  Pb + KCl  Al + HCl 

28. #4 - Double Replacement Two things replace each other. –Reactants must be two ionic compounds. –Usually in aqueous solution NaOH + FeCl 3  –The positive ions change place. NaOH + FeCl 3  Fe +3 OH - + Na +1 Cl -1 NaOH + FeCl 3  Fe(OH) 3 + NaCl

29. The Activity Series of the Halogens Fluorine Chlorine Bromine Iodine Halogens can replace other halogens in compounds, provided that they are above the halogen that they are trying to replace. 2NaCl (s) + F 2(g)  2NaF (s) + Cl 2(g) MgCl 2(s) + Br 2(g)  ??? No Reaction ??? Higher Activity Lower Activity

30. #4 - Double Replacement Has certain “driving forces” –Will only happen if one of the products: a) doesn’t dissolve in water and forms a solid (a “precipitate”), or b) is a gas that bubbles out, or c) is a molecular compound (usually water).

31. Complete and balance: assume all of the following reactions actually take place: CaCl 2 + NaOH  CuCl 2 + K 2 S  KOH + Fe(NO 3 ) 3  (NH 4 ) 2 SO 4 + BaF 2 

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

33. Practice Examples: H 2 + O 2  H 2 O  Zn + H 2 SO 4  HgO  KBr +Cl 2  AgNO 3 + NaCl  Mg(OH) 2 + H 2 SO 3 

34. #5 - Combustion Means “add oxygen” Normally, a compound composed of only C, H, (and maybe O) is reacted with oxygen – usually called “burning” If the combustion is complete, the products will be CO 2 and H 2 O. If the combustion is incomplete, the products will be CO (or possibly just C) and H 2 O.

35. Combustion Examples: C 4 H 10 + O 2  (assume complete) C 4 H 10 + O 2  (incomplete) C 6 H 12 O 6 + O 2  (complete) C 8 H 8 +O 2  (incomplete)

36. SUMMARY: an equation... Describes a reaction Must be balanced in order to follow the Law of Conservation of Mass Can only be balanced by changing the coefficients. Has special symbols to indicate physical state, if a catalyst or energy is required, etc.

37. Reactions Come in 5 major types. We can tell what type they are by looking at the reactants. Single Replacement happens based on the Activity Series Double Replacement happens if the product is a precipitate (insoluble solid), water, or a gas.

38. Section 11.3 Reactions in Aqueous Solution OBJECTIVES: –Describe the information found in a net ionic equation.

39. Section 11.3 Reactions in Aqueous Solution OBJECTIVES: –Predict the formation of a precipitate in a double replacement reaction.

40. Net Ionic Equations Many reactions occur in water- that is, in aqueous solution Many ionic compounds “dissociate”, or separate, into cations and anions when dissolved in water Now we are ready to write an ionic equation

41. Net Ionic Equations Example: –AgNO 3 + NaCl  AgCl + NaNO 3 1. this is the full equation 2. now write it as an ionic equation 3. can be simplified by eliminating ions not directly involved (spectator ions) = net ionic equation

42. Predicting the Precipitate Insoluble salt = a precipitate note Figure 11.11, p.342 General solubility rules are found: a)Table 11.3, p. 344 b)Reference section - page R54 (back of textbook) c)Lab manual Table A.3, page 332

43. End of Section 11.2

44. End of Chapter 7 End of Chapter 11