1. 1 Chapter 11 “Chemical Reactions”

2. 2 All chemical reactions… l have two parts: –Reactants - the substances you start with –Products- the substances you end up with l The reactants turn into the products. Reactants  Products

3. 3 - Page 321 Reactants Products

4. 4 In a chemical reaction l Atoms aren’t created or destroyed. l 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

5. 5 Symbols in equations - page 323 l the arrow separates the reactants from the products –Read as “reacts to form” or yields l The plus sign = “and” l (s) after the formula = solid: AgCl (s) l (g) after the formula = gas: CO 2(g) l (l) after the formula = liquid: H 2 O (l)

6. 6 Symbols used in equations l (aq) after the formula = dissolved in water, an aqueous solution: NaCl (aq) is a salt water solution  used after a product indicates a gas has been produced: H 2 ↑  used after a product indicates a solid has been produced: PbI 2 ↓

7. 7 Symbols used in equations ■ indicates a reversible reaction (more later) ■ shows that heat is supplied to the reaction ■ is used to indicate a catalyst is supplied, in this case, platinum.

8. 8 What is a catalyst? l A substance that speeds up a reaction, without being changed or used up by the reaction. l Enzymes are biological or protein catalysts.

9. 9 3. The Skeleton Equation l Uses formulas and symbols to describe a reaction –but doesn’t indicate how many; this means they are NOT balanced l All chemical equations are a description that describe reactions.

10. 10 Write a skeleton equation for: 1. Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas. 2. Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

11. 11 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)

12. 12 4. Balanced Chemical Equations l Atoms can’t be created or destroyed in an ordinary reaction: –All the atoms we start with we must end up with l A balanced equation has the same number of each element on both sides of the equation.

13. 13 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! (I prefer to save O until the very last) 4)Check to make sure it is balanced.

14. 14 l 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 l Never put a coefficient in the middle of a formula 2NaCl is okay, but Na2Cl is not.

15. 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. 16 Sec 2: Types of Reactions l There are millions of reactions. l We can’t remember them all, but luckily they will fall into several categories. l We will learn 5 major types. l Will be able to predict the products. l For some, we will be able to predict whether or not they will happen at all. l We recognize them by their reactants

17. 17 #1 - Synthesis Reactions l Combine = put together l 2 substances combine to make one compound. Ca +O 2  CaO SO 3 + H 2 O  H 2 SO 4 l We can predict the products if the reactants are two elements. Mg + N 2 

18. 18 Complete and balance: Ca + Cl 2  Fe + O 2  (assume iron (II) oxide is product) Al + O 2  l Remember that the first step is to write the correct formulas – you can still change the subscripts at this point, but not later! l Then balance by using the coefficients only

19. 19 #1 - Synthesis l Note: a) Some nonmetal oxides react with water to produce an acid: SO 2 + H 2 O H 2 SO 3 b) Some metallic oxides react with water to produce a base: CaO + H 2 O Ca(OH) 2

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

21. 21 #2 - Decomposition Reactions l Can predict the products if it is a binary compound-Made up of only two elements –breaks apart into its elements: lH2OlH2O l HgO

22. 22 #2 - Decomposition Reactions l 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 l NiCO 3 CO 2 + ___ H 2 CO 3 (aq)  CO 2 + ___ heat

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

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

25. 25 #3 Single Replacement l 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 l There is a list on page 333 - called the Activity Series of Metals l Higher on the list replaces lower.

26. 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. 27 #3 Single Replacement Practice: Fe + CuSO 4  Pb + KCl  Al + HCl 

28. 28 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

29. 29 #4 - Double Replacement l 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

30. 30 #4 - Double Replacement l 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. 31 Complete and balance: l 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. 32 How to recognize which type l Look at the reactants: E + E =Combination C =Decomposition E + C =Single replacement C + C =Double replacement

33. 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. 34 #5 - Combustion l Means “add oxygen” l Normally, a compound composed of only C, H, (and maybe O) is reacted with oxygen – usually called “burning” l If the combustion is complete, the products will be CO 2 and H 2 O. l If the combustion is incomplete, the products will be CO (or possibly just C) and H 2 O.

35. 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. 36 SUMMARY: an equation... l Describes a reaction l Must be balanced in order to follow the Law of Conservation of Mass l Can only be balanced by changing the coefficients. l Has special symbols to indicate physical state, if a catalyst or energy is required, etc.

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

38. 38 Sec. 3: Net Ionic Equations l Many reactions occur in water- that is, in aqueous solution l Many ionic compounds “dissociate”, or separate, into cations and anions when dissolved in water l Now we are ready to write an ionic equation

39. 39 Net Ionic Equations l 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

40. 40 Total Ionic Equations Molecular Equation: K 2 CrO 4 + Pb(NO 3 ) 2  PbCrO 4 + 2 KNO 3 Soluble SolubleInsoluble Soluble Total Ionic Equation: 2 K + + CrO 4 -2 + Pb +2 + 2 NO 3 -  PbCrO 4 (s) + 2 K + + 2 NO 3 -

41. 41 Predicting the Precipitate l Insoluble salt = a precipitate note Figure 11.11, p.342 l General solubility rules are found: a)Table 11.3, p. 344 b)Reference section - page R54 (back of textbook)

42. 42 Net Ionic Equations l These are the same as total ionic equations, but you should cancel out ions that appear on BOTH sides of the equation Total Ionic Equation: 2 K + + CrO 4 -2 + Pb +2 + 2 NO 3 -  PbCrO 4 (s) + 2 K + + 2 NO 3 - Net Ionic Equation: CrO 4 -2 + Pb +2  PbCrO 4 (s)

43. 43 Net Ionic Equations l Try this one! Write the molecular, total ionic, and net ionic equations for this reaction: Silver nitrate reacts with Lead (II) Chloride in hot water. Molecular: Total Ionic: Net Ionic:

44. 44 Practice Problems l Write the formula for pure iron(III) chloride, including the state l Write an equation showing what happens to iron(III) chloride when it is added to water

45. 45 Net Ionic Equations Revisited: 1)Write the (balanced!) molecular equation first -Reaction products: swap cations and anions -Predict solubility (using Solubility rules) 2)Write the complete ionic equation next -s) compounds don’t ionize -( (aq) compounds do ionize ion subscripts in the molecular equation become coefficients in the complete ionic equation! 3)Write the net ionic equation next -cancel spectator ions The net ionic equation is a “simplified” form of the complete ionic equation

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