1 Chapter 7 Chemical Reactions Killarney High School

2 Section 8.1 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 chlorine to form copper (II) chloride. 2. In a word equation Copper + chlorine  copper (II) chloride

6 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

7 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))

8 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.

9 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.

10 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.

11 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.

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

13 Balancing Chemical Equations

14 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.

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

16 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

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

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

19 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!  Check to make sure it is balanced.

20 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 –2 NaCl is okay, Na2Cl is not.

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

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

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

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

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

26 Example H 2 +H2OH2OO2O2  Recount RP H O

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

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

29 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

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

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

32 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

33 #1 - Combination 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 they are two elements. Mg + N 2 

34 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

35 #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

36 #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

37 #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 + ?

38 #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

39 #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 

40 #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.

41 #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 

42 #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 

43 #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

44 #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).

45 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 

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

47 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 

48 #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.

49 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)

50 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.

51 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.

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

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

54 ION Group 1A NH 4 + NO 3 - CH 3 COO - Alkali’sH+H+ Soluble (aq) All Insoluble (s) None ION Cl - Br – I - SO 4 2- S 2- OH - PO 4 3- SO 3 2- CO 3 2- Soluble (aq) Most Group 1A & IIA NH 4 + Group 1A NH 4 + Sr 2+ Ba 2+ Group 1A NH 4 + Insoluble (s) Ag + Pb 2+ Hg + Cu + Ag + Pb 2+ Ca 2+ Ba 2+ Sr 2+ Most SOLUBILITY RULES

55 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

56 Net Ionic Equations l Example: –AgNO 3(aq) + NaCl (aq)  AgCl (s) + NaNO 3(aq) 1. this is the full equation

57 2. now write it as an total ionic equation l Ag + NO 3 - (aq) + Na + Cl - (aq)  AgCl (s) + Na + NO 3 - (aq)

58 3. can be simplified by eliminating ions not directly involved (spectator ions) = net ionic equation l Ag + (aq) + Cl - (aq)  AgCl (s) This is the only change occurring in the reaction. What were the spectator ions? l Ag + NO 3 - (aq) + Na + Cl - (aq)  AgCl (s) + Na + NO 3 - (aq)

59 Predicting the Precipitate l Insoluble salt = a precipitate -note Figure 7.14, p.156 l General rules: Table A.7, p. 673, - (back of textbook) or Your Table of Ions