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Unit 5 Chemical Equations & Reactions

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1 Unit 5 Chemical Equations & Reactions
Honors Chemistry

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

3 In a chemical reaction The way atoms are joined is changed
Atoms aren’t created of destroyed. 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

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

5 Symbols used in equations
(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))

6 Symbols used in equations
indicates a reversible reaction shows that heat is supplied to the reaction is used to indicate a catalyst is supplied, in this case, platinum.

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

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

9 Convert these to equations
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.

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

11 Balanced Equation Atoms can’t be created or destroyed
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.

12 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.

13 Never Never change a subscript to balance an equation.
If you change the formula you are describing a different reaction. H2O is a different compound than H2O2 Never put a coefficient in the middle of a formula 2 NaCl is okay, Na2Cl is not.

14 Example H2 + O2 H2O R P 2 H 2 2 O 1 Need twice as much O in the product

15 Example H2 + O2 2 H2O R P 2 H 2 2 O 1 Changes the O

16 Example H2 + O2 2 H2O R P 2 H 2 2 O 1 2 Also changes the H

17 Example H2 + O2 2 H2O R P 2 H 2 4 2 O 1 2 Need twice as much H in the reactant

18 Example 2 H2 + O2 2 H2O R P 2 H 2 4 2 O 1 2 Recount

19 Example 2 H2 + O2 2 H2O R P 4 2 H 2 4 2 O 1 2 The equation is balanced, has the same number of each kind of atom on both sides

20 Example R P 4 2 H 2 4 2 O 1 2 2 H2 + O2 ® 2 H2O This is the answer
Not this

21 Balancing Examples _Mg + _N2 ® _Mg3N2 _P + _O2 ® _P4O10
_AgNO3 + _Cu ® _Cu(NO3)2 + _Ag _Mg + _N2 ® _Mg3N2 _P + _O2 ® _P4O10 _Na + _H2O ® _H2 + _NaOH _CH4 + _O2 ® _CO2 + _H2O

22 Balancing Examples 2AgNO3 + Cu ® Cu(NO3)2 + 2Ag 3Mg +N2 ® Mg3N2
4P + 5O2 ® P4O10 2Na + 2H2O ® H2 + 2NaOH CH4 + 2O2 ® CO2 + 2H2O

23 Types of Reactions There are millions of reactions.
Can’t remember them all Fall into several categories. We will learn 5 major types. Will be able to predict the products. For some, we will be able to predict whether they will happen at all. Will recognize them by the reactants

24 CaO(s) + SO2(g)  CaSO3(s)
COMBINATION REACTION A reaction in which two or more substances combine to form a single product. A +B + C  ABC CaO(s) + SO2(g)  CaSO3(s)

25 Write and balance Ca + Cl2 ® Fe + O2 ® Al + O2 ®
Remember that the first step is to write the correct formulas Then balance by using coefficients only

A reaction in which a single compound reacts to give two or more substances, usually requiring a raise in temperature. ABC  A + B + C 2KClO3(s)  2KCl(s) + 3O2(g)

27 #2 - Decomposition Reactions
decompose = fall apart one reactant falls apart into two or more elements or compounds. NaCl Na + Cl2 CaCO CaO + CO2 Note that energy is usually required to decompose

28 #2 - Decomposition Reactions
Can predict the products if it is a binary compound Made up of only two elements Falls apart into its elements H2O HgO

29 Can predict the products if it is a binary compound
2H2O(l) → Can predict the products if it is a binary compound Made up of only two elements Falls apart into its elements Fig. 8-13, p. 215

30 #2 - Decomposition Reactions
If the compound has more than two elements you must be familiar with page 10 of your packet NiCO3 ® KClO3(aq) ®

31 Single Displacement (or Replacement) Reactions
pp. 218, 220

32 Fig. 8-15, p. 218

33 Single Displacement (or Replacement) Reactions
PREDICT THE PRODUCT 1. Ca + HCl  2. ZnBr2 + I2  3. Cu + AgNO3  Answers are on the next slide.

34 #3 Single Replacement Metals replace other metals (and hydrogen)
K + AlN ® Zn + HCl ® Think of water as HOH Metals replace one of the H, combine with hydroxide. Na + HOH ®

35 We can tell whether a reaction will happen
Some chemicals are more “active” than others More active replaces less active There is a list on page 11 - called the Activity Series of Metals Higher on the list replaces lower.

36 #3 Single Replacement Note the * concerning Hydrogen
H can be replaced in acids by everything higher Li, K, Ba, Ca, & Na replace H from acids and water Fe + CuSO4 ® Pb + KCl ® Al + HCl ®

37 #3 - Single Replacement What does it mean that Hg and Ag are on the bottom of the list? Nonmetals can replace other nonmetals Limited to F2 , Cl2 , Br2 , I2 (halogens) Higher replaces lower. F2 + HCl ® Br2 + KCl ®

38 Double Displacement (or Replacement) Reactions
pp. 220, 223

39 Double Displacement (or Replacement) Reactions
PREDICT THE PRODUCT & BALANCE 1. MgSO4 + LiOH  ___________ 2. Pb(NO3)2 + Na2CO3  ____________ 3. HNO3 + Ba(OH)2  ___________ Answers are on the next slide.

40 Double Displacement (or Replacement) Reactions
ANSWERS 1. MgSO4 + 2 LiOH  Mg(OH)2 + Li2SO4 2. Pb(NO3)2 + Na2CO3  PbCO3 + 2 NaNO3 3. 2 HNO3 + Ba(OH)2  Ba(NO3)2 + 2 H2O Exchange cations

41 Double Replacement Two things replace each other.
Reactants must be two ionic compounds or acids. Usually in aqueous solution NaOH + FeCl3 ® The positive ions change place. NaOH + FeCl3 ® Fe+3 OH- + Na+1 Cl-1 NaOH + FeCl3 ® Fe(OH)3 + NaCl

42 Double Replacement 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).

43 Oxidation and Reduction Reactions
Early chemists saw oxidation as the combination of a material with oxygen Today, many of these reactions may not even involve oxygen Redox currently says that electrons are transferred between reactants Always occur simultaneously. Called redox reactions.

44 Oxidation and Reduction Reactions
Oxidation – refers to a loss of electrons, or gain of oxygen Reduction – refers to a gain of electrons, or loss of oxygen Oxidizing agent – the substance being reduced Reducing agent – the substance being oxidized

45 Oxidation and Reduction Reactions
Loss of Electrons is Oxidation Gain of Electrons is Reduction LEO SAYS GER

46 Rules For Assigning Oxidation States
Oxidation numbers are used to keep track of how many electrons are lost or gained by each atom

47 Balancing Redox Reactions
If the oxidation number changes, that element has undergone oxidation or reduction Reaction as a whole must be redox Half-reaction – equation showing just the oxidation or just the reduction

48 Balancing Redox Reactions (in an acid)
SO2 + MnO4- →SO42- + Mn2+ Divide the equation into half-reactions. SO2 → SO MnO4- → Mn2+ Oxidation Reduction Balance all atoms except O and H

49 Balancing Redox Equations
Balance O by adding H2O. 2H2O + SO2 →SO MnO4- →Mn2+ + 4H2O Balance H by adding H+. 2H2O + SO2 → SO H+ 8H+ + MnO4- → Mn2+ + 4H2O

50 Balance net charge by adding e-.
2H2O + SO2 →SO H+ + 2e- 5e- + 8H+ + MnO4- →Mn2+ + 4H2O Step 6. Make e- gain equal e- loss; then add half- reactions. 5(2H2O + SO2 →SO H+ + 2e-) 2(5e- + 8H+ + MnO4- →Mn2+ + 4H2O) ______________________________________ 10H2O + 5SO2 + 10e- + 16H+ + 2MnO4- →2Mn2+ + 8H2O + 5SO H+ + 10e-

51 Cancel anything that's the same on both sides.
2H2O + 5SO2 +2MnO4- →2Mn2+ + 5SO H+

52 Balancing Redox Reactions (in a base)
Follow steps 1 – 7 then: 8. Identify the number of proton (H+) in the acidic answer. Add the same number of OH+ ions to BOTH sides of the equation. 9. If H+ and OH+ appear on the same side of the equation, they will react in a 1:1 ratio to form H2O. 10. Cancel out water molecules that appear on both sides of the equation. 11. Check that all is balanced

A reaction where an insoluble solid is formed during a reaction between two aqueous solutions. (aq) (aq)  (aq) (s) 2KI(aq) + Pb(NO3)2(aq)  2KNO3(aq) + PbI2(s) Exchange cations

54 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 AgNO3 + NaCl  AgCl + NaNO3 1. this is the full equation 2. Write it as an ionic equation 3. Eliminating ions not directly involved (spectator ions) = net ionic equation

55 Molecular, Ionic, Net Ionic Equations

56 How Do We Predict Solubility?

A reaction between an acid and a base which results in the production of a salt and water. HA + BOH  (metal/nonmetal) + H2O HNO3(aq) + KOH(aq)  KNO3(aq) + H2O(l)

58 2C4H10(g) + 13O2(g)  8CO2(g) + 10H2O(g)
COMBUSTION REACTION Means “add oxygen” A compound composed of only C, H, and maybe O is reacted with oxygen If the combustion is complete, the products will be CO2 and H2O. If the combustion is incomplete, the products will be CO (possibly just C) and H2O. A reaction of a substance with oxygen, usually the rapid release of heat produces a flame. CH + O2  CO2 + H2O 2C4H10(g) + 13O2(g)  8CO2(g) + 10H2O(g)

59 2C4H10(g) + 13O2(g)  8CO2(g) + 10H2O(g) + heat (in Joules)
Combustion Many times in a combustion reaction, heat energy is given off. In chemical terms this is called an exothermic reaction. Thermochemistry is field of chemistry which studies the transfer of heat in a reaction. The thermodynamic equation representing this exothermic reaction is: 2C4H10(g) + 13O2(g)  8CO2(g) + 10H2O(g) + heat (in Joules)

60 Examples C4H10 + O2 ® (assume complete) C4H10 + O2 ® (incomplete)
C6H12O6 + O2 ® (complete) C8H8 +O2 ® (incomplete)

A reaction that produces a gas from reactants not in the gaseous state. 2 HCl (aq) + ZnS (s)  ZnCl2 (aq) + H2S (g) Zn (s) + 2 HCl (aq)  ZnCl2 (aq) + H2 (g) Many gas formation reactions involve two steps, first the double displacement reaction then the decomposition reaction of an unstable substance. Na2CO3 + 2HCl  2 NaCl + H2CO3 H2CO3  CO2 + H2O Besides carbonic acid (H2CO3), sulfurous acid (H2SO3) also decomposes into SO2 and water.

NH4OH → NH3 (g) + H2O (l) H2CO3 → CO2 (g) + H2O (l) H2SO3 → SO2 (g) + H2O (l)

63 Remember 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, and if a catalyst or energy is required.

64 Table 8-3, p. 223

65 Examples H2 + O2 ® H2O ® Zn + H2SO4 ® HgO ® KBr +Cl2 ® AgNO3 + NaCl ®
Mg(OH)2 + H2SO3 ®

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