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Chapter 7 Chemical Reactions

Chapter 7 Chemical Reactions. Section 7.1 Describing Chemical Change. OBJECTIVES: Write equations describing chemical reactions, using appropriate symbols. Section 7.1 Describing Chemical Change. OBJECTIVES:

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Chapter 7 Chemical Reactions

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  1. Chapter 7 Chemical Reactions

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

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

  4. 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

  5. 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 silver nitrate to form silver and copper (II) nitrate . 2. In a word equation Copper + silver nitrate ® silver + copper (II) nitrate

  6. Or a skeleton equation Cu( ) + AgNO3( )® Ag( ) + Cu(NO3)2( ) Or a balanced equation products reactants

  7. Symbols in equations-p.144 • 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

  8. 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))

  9. 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.

  10. 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.

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

  12. 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.

  13. Now, read these: • Fe(s) + O2(g) ® Fe2O3(s) • Cu(s) + AgNO3(aq) ® Ag(s) + Cu(NO3)2(aq)

  14. NO2 (g) N2(g) + O2(g)

  15. Balancing Chemical Equations

  16. 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.

  17. ® O + • C + O2® CO2 • This equation is already balanced • What if it isn’t? C C O O O

  18. ® O + • C + O2® CO • We need one more oxygen in the products. • Can’t change the formula, because it describes what it is (carbon monoxide in this example) C C O O

  19. C O ® O + • Must be used to make another CO • But where did the other C come from? C O C O

  20. C C O ® O + • Must have started with two C • 2 C + O2® 2 CO O C O 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! Element Carbon Hydrogen Oxygen • Check to make sure it is balanced.

  22. Never • Never change a subscript to balance an equation. • If you change the formula (subscripts) 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.

  23. Example H2 + O2 ® H2O Make a table to keep track of where you are at

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

  25. Example H2 + O2 ® 2 H2O R P 2 H 2 2 O 1 Changes the O

  26. Example H2 + O2 ® 2 H2O R P 2 H 2 2 O 1 2 Also changes the H

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

  28. Example 2 H2 + O2 ® 2 H2O R P 2 H 2 4 2 O 1 2 Recount

  29. 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

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

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

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

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

  34. 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

  35. #1 - Combination Reactions • Combine - put together (synthesis) • 2 substances combine to make one compound. • Ca +O2® CaO • SO3 + H2O ® H2SO4 • We can predict the products if they are two elements. • Mg + N2® __?__

  36. Write and balance • Ca + Cl2® • Fe + O2® iron (II) oxide • Al + O2® • Remember that the first step is to write the correct formulas • Then balance by using coefficients only

  37. #2 - Decomposition Reactions • decompose = fall apart • one reactant falls apart into two or more elements or compounds. • NaCl Na + Cl2 • CaCO3 CaO + CO2 • Note that energy is usually required to decompose

  38. #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

  39. #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 • NiCO3 CO2 + ? • H2CO3(aq) ® CO2 + ?

  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. CuNO3 + Ag Cu + AgNO3 ??? No reaction ??? Zn + NaCl Copy list to your periodic table

  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) + F2(g) 2NaF(s) + Cl2(g) ??? No Reaction MgCl2(s) + Br2(g) ???

  42. #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 • F2 + LiCl ® LiF + Cl2

  43. #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 ®

  44. #3 Single Replacement • 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 155 - called the Activity Series of Metals • Higher on the list replaces lower.

  45. #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 ®

  46. #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 ®

  47. #4 - 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+3OH- + Na+1Cl-1 • NaOH + FeCl3® Fe(OH)3 + NaCl

  48. #4 - 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).

  49. Complete and balance • assume all of the following reactions take place: CaCl2 + NaOH ® CuCl2 + K2S ® KOH + Fe(NO3)3® (NH4)2SO4 + BaF2®

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

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