1 Chapter 7 Chemical Reactions

2 Section 7.1 Describing Chemical Change l OBJECTIVES: –Write equations describing chemical reactions, using appropriate symbols

3 Section 7.1 Describing Chemical Change l OBJECTIVES: –Write balanced chemical equations, when given the names or formulas of the reactants and products in a chemical reaction.

4 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

5 In a chemical reaction l The way atoms are joined is changed l Atoms aren’t created of destroyed. l Can be described several ways: 1. In a sentence Copper reacts with silver nitrate to form silver and copper (II) nitrate. 2. In a word equation Copper + silver nitrate  silver + copper (II) nitrate

6 Cu ( ) + AgNO 3 ( )  Ag ( ) + Cu(NO 3 ) 2 ( ) reactants products Or a skeleton equation Or a balanced equation

7 Symbols in equations-p.144 l the arrow separates the reactants from the products l Read “reacts to form” l The plus sign = “and” l (s) after the formula = solid l (g) after the formula = gas l (l) after the formula = liquid

8 Symbols used in equations l (aq) after the formula - dissolved in water, an aqueous solution.  used after a product indicates a gas (same as (g))  used after a product indicates a solid (same as (s))

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

10 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.

11 Skeleton Equation l Uses formulas and symbols to describe a reaction l doesn’t indicate how many. l All chemical equations are sentences that describe reactions.

12 Convert these to equations l Solid iron (III) sulfide reacts with gaseous hydrogen chloride to form iron (III) chloride and hydrogen sulfide gas. l Nitric acid dissolved in water reacts with solid sodium carbonate to form liquid water and carbon dioxide gas and sodium nitrate dissolved in water.

13 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)

14 l NO 2 (g) N 2 (g) + O 2 (g)

15 Balancing Chemical Equations

16 Balanced Equation l Atoms can’t be created or destroyed l 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.

17 C + O 2  CO 2 l This equation is already balanced l What if it isn’t? C + O O  C O O

18 C + O 2  CO l We need one more oxygen in the products. l Can’t change the formula, because it describes what it is (carbon monoxide in this example) C + O  C O O

19 l Must be used to make another CO l But where did the other C come from? C + O  C O O O C

20 l Must have started with two C 2 C + O 2  2 CO C + O  C O O O C C

21 Rules for balancing:  Assemble, write the correct formulas for all the reactants and products  Count the number of atoms of each type appearing on both sides  Balance the elements one at a time by adding coefficients (the numbers in front) - save H and O until LAST! E lement C arbon H ydrogen O xygen  Check to make sure it is balanced.

22 l Never change a subscript to balance an equation. –If you change the formula (subscripts) 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 –2 NaCl is okay, Na2Cl is not.

23 Example H 2 +H2OH2OO2O2  Make a table to keep track of where you are at

24 Example H 2 +H2OH2OO2O2  Need twice as much O in the product RP H O

25 Example H 2 +H2OH2OO2O2  Changes the O RP H O

26 Example H 2 +H2OH2OO2O2  Also changes the H RP H O

27 Example H 2 +H2OH2OO2O2  Need twice as much H in the reactant RP H O

28 Example H 2 +H2OH2OO2O2  Recount RP H O

29 Example H 2 +H2OH2OO2O2  The equation is balanced, has the same number of each kind of atom on both sides RP H O

30 Example H 2 +H2OH2OO2O2  This is the answer RP H O Not this

31 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

32 Section 7.2 Types of Chemical Reactions l OBJECTIVES: –Identify a reaction as combination, decomposition, single-replacement, double-replacement, or combustion

33 Section 7.2 Types of Chemical Reactions l OBJECTIVES: –Predict the products of combination, decomposition, single-replacement, double-replacement, and combustion reactions.

34 Types of Reactions l There are millions of reactions. l Can’t remember them all l 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 they will happen at all. l Will recognize them by the reactants

35 #1 - Combination Reactions l Combine - put together (synthesis) 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 they are two elements. Mg + N 2 

36 Write and balance Ca + Cl 2  Fe + O 2  iron (II) oxide Al + O 2  l Remember that the first step is to write the correct formulas l Then balance by using coefficients only

37 #2 - Decomposition Reactions l decompose = fall apart l one reactant falls apart into two or more elements or compounds. l NaCl Na + Cl 2 l CaCO 3 CaO + CO 2 l Note that energy is usually required to decompose

38 #2 - Decomposition Reactions l Can predict the products if it is a binary compound l Made up of only two elements l Falls apart into its elements lH2OlH2O l HgO

39 #2 - Decomposition Reactions l If the compound has more than two elements you must be given one of the products l The other product will be from the missing pieces l NiCO 3 CO 2 + ? H 2 CO 3 (aq)  CO 2 + ?

40 The Activity Series of the Metals Lithium Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Lead Hydrogen Bismuth Copper Mercury Silver Platinum Gold Metals can replace other metals provided that they are above the metal that they are trying to replace. Ex. Metals above hydrogen can replace hydrogen in acids. Cu + AgNO 3 Zn + NaCl Copy list to your periodic table CuNO 3 + Ag No reaction ???

41 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 ???

42 #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

43 #3 Single Replacement l Metals replace other metals (and hydrogen) K + AlN  Zn + HCl  l Think of water as HOH l Metals replace one of the H, combine with hydroxide. Na + HOH 

44 #3 Single Replacement l We can tell whether a reaction will happen l Some chemicals are more “active” than others l More active replaces less active l There is a list on page called the Activity Series of Metals l Higher on the list replaces lower.

45 #3 Single Replacement l Note the * concerning Hydrogen l H can be replaced in acids by everything higher l Li, K, Ba, Ca, & Na replace H from acids and water Fe + CuSO 4  Pb + KCl  Al + HCl 

46 #3 - Single Replacement l What does it mean that Hg and Ag are on the bottom of the list? l Nonmetals can replace other nonmetals l Limited to F 2, Cl 2, Br 2, I 2 (halogens) l Higher replaces lower. F 2 + HCl  Br 2 + KCl 

47 #4 - Double Replacement l Two things replace each other. l Reactants must be two ionic compounds or acids. l Usually in aqueous solution NaOH + FeCl 3  l The positive ions change place. NaOH + FeCl 3  Fe +3 OH - + Na +1 Cl -1 NaOH + FeCl 3  Fe(OH) 3 + NaCl

48 #4 - Double Replacement l Has certain “driving forces” –Will only happen if one of the products: –doesn’t dissolve in water and forms a solid (a “precipitate”), or –is a gas that bubbles out, or –is a covalent compound (usually water).

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

50 How to recognize which type l Look at the reactants: Combination Element + ElementCompound (s,l or g) (s,l or g) (s,l or g)

51 Decomposition Compound Element + Element (s,l or g) (s,l or g) (s,l or g)

52 Single Displacement Element + Compound (s,l or g) (aq) (s,l or g) (aq or s)

53 Compound + Compound (aq) (aq) (aq or s) (aq or s) Double Displacement

54 Combustion C x H y + O 2 (g) CO 2(g) + H 2 O (g)

55 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 

56 #5 - Combustion l Means “add oxygen” l A compound composed of only C, H, and maybe O is reacted with oxygen 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 (possibly just C) and H 2 O.

57 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)

58 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, and if a catalyst or energy is required.

59 Reactions l Come in 5 major types. l Can tell what type they are by the reactants. l Single Replacement happens based on the activity series l Double Replacement happens if the product is a solid, water, or a gas.

60 Section 7.3 Reactions in Aqueous Solution l OBJECTIVES: –Write and balance net ionic equations.

61 Section 7.3 Reactions in Aqueous Solution l OBJECTIVES: –Use solubility rules to predict the precipitate formed in double- replacement reactions.

62 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 can write a complete ionic equation

63 Net Ionic Equations l Example: –AgNO 3(aq) + NaCl(aq)  AgCl(s)+ NaNO 3(aq) 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

64 Predicting the Precipitate l Insoluble salt = a precipitate - note Figure 7.13, p.156 l General rules: Table 7.3, p. 161, Reference p.708 (back of textbook) l Sample problem 7-9, p.156