Balancing Chemical Reactions

1. Balancing Chemical Reactions

2. Objectives • Write equations describing chemical reactions using appropriate symbols • Write balanced chemical equations when given the names or formulas of the reactants and products in a chemical reaction • Define chemical equation, catalyst, aqueous solution, skeleton equation, coefficients, and balanced equation

3. Chemical Equations • Chemical equations – using chemical formulas to write equations • Reactants (left side of arrow) • Products (right side of arrow) • Arrow means yields, gives, or reacts to produce • Reactants  Products • Catalyst (a substance that speeds up the rate of the reaction but that is not used up in the reaction) should be written above the arrow

4. Can indicate the physical state of a substance in the equation by putting a symbol after each formula Solid – (s) Liquid – (l) Gas – (g) Aqueous solution: a substance dissolved in water – (aq)

5. Iron reacts with oxygen to produce rust. Iron + oxygen  iron(III) oxide • Hydrogen peroxide reacts to form water and oxygen. (Bubbles=oxygen gas) Hydrogen peroxide  water + oxygen • Residential heating Methane + oxygen  carbon dioxide + water • Green color on the Statue of Liberty (copper exposed to moist air) Copper + carbon dioxide + water  basic copper(II) carbonate

6. Glucose is fermented by yeast to form ethanol and carbon dioxide C6H12O6 C2H5OH(aq) + CO2 (g) Yeast should be written above the arrow because it is neither a reactant nor a product. • Plants carry out photosynthesis - the creation of glucose and oxygen from carbon dioxide, water, and sunlight Carbon dioxide + water  glucose + oxygen Sunlight should be written above the arrow because it is neither a reactant nor a product.

7. Skeleton Equation • A chemical equation that does not indicate the relative amounts of the reactants and products involved in the reaction • Examples: a. Fe(s) + O2(g)  Fe2O3(s) b. H2O2(aq)  H2O(l) + O2(g) Manganese(IV) oxide is a catalyst, so MnO2 should be written above the arrow. See Figure 8.3 on page 205 in text.

8. Link to Human Physiology • Hydrogen peroxide decomposes to oxygen and water when it comes into contact with blood. The enzyme catalase is a catalyst that contains an iron(II) ion. When hydrogen peroxide is poured on a cut, it reacts with the iron(II) ions of catalase, thus releasing energetic oxygen atoms that produce the antiseptic effect of hydrogen peroxide.

9. Write the skeleton equation for hydrogen peroxide linked to human physiology • Answer: H2O2(aq)  H2O(l) + O2(g) Iron(II) ion is a catalyst, so Fe2+ should be written above the arrow. • This is an example of a decomposition reaction : a single compound is broken down into two or more products

10. Link to Art • Hydrogen peroxide can also be used to restore the clarity of old paintings. • Lead-based paints darken with time (PbS). • Hydrogen peroxide converts PbS to PbSO4.

11. Write a Skeleton Equation • Solid sodium hydrogen carbonate reacts with hydrochloric acid to produce aqueous sodium chloride, water, and carbon dioxide gas. Include appropriate symbols. 1. Write the correct formula for each substance in the reaction. 2. Separate the reactants from the products. 3. Indicate the physical state of each substance.

12. Answer • NaHCO3(s) + HCl(aq)  NaCl(aq) + H2O(l) + CO2(g)

13. A Balanced Equation • An equation that gives the correct quantity of each reactant and product • Coefficients (numbers placed in front of the symbols) are used • Must obey law of conservation of mass: Each side of the equation has the same number of atoms of each element • Example: A standard bicycle is composed of one frame, two wheels, one handlebar, and two pedals F + 2W + H + 2P  FW2HP2

14. Rules for Balancing Equations • 1. Determine the correct formulas for all of the reactants and products. In some cases, also list in parenthesis the physical state of matter. • 2. List reactants on the left side of the arrow (Use plus sign (+) when there is more than one reactant) • 3. List the products on the right side of the arrow (Use plus sign (+) when there is more than one product) • 4. Steps 1-3 provide a skeleton equation. (Note: Sometimes this is also the balanced equation. For example: C + O2 CO2)

15. 5. Count the number of atoms of each element in the reactants and products. For simplicity, a polyatomic ion appearing unchanged on both sides of the arrow is counted as a single unit. • 6. Balance the elements one at a time by using coefficients. DO NOT CHANGE THE SUBSCRIPTS. • 7. Check each atom or polyatomic ion to be sure that the equation is balanced. • 8. Make sure that all the coefficients are in the lowest possible ratio that balances.

16. Problem: Hydrogen and oxygen react to form water. Write a balanced equation. • Reactants: H2(g) + O2(g) • Products: H2O(l) • H2(g) + O2(g)  H2O(l) • Count the atoms Left side Right side H – 2 H – 2 O – 2 O – 1 • Use coefficient to get 2 oxygen on the right side: H2(g) + O2(g)  2 H2O(l) Left side Right side H – 2 H – 4 O – 2 O – 2

17. Need 4 hydrogen atoms, so place a coefficient of 2 in front of H2 2H2(g) + O2(g)  2 H2O(l) Left side Right side H – 4 H – 4 O – 2 O – 2 • Check number of atoms • Check that the coefficients are in the lowest possible ratio • The equation is balanced

18. Problems • 1. Balance the following equations: a. SO2 + O2 SO3 b. Al + O2 Al2O3 • 2. Rewrite the word equation as a balanced chemical equation: Aluminum sulfate and calcium hydroxide react to form aluminum hydroxide and calcium sulfate.

19. Answers 1a) 2SO2 + O2 2SO3 1b) 4Al + 3O2 2Al2O3 2) Word equation to balanced chemical equation: Al2(SO4)3 + 3Ca(OH)2 2Al(OH)3 + 3CaSO4

20. Practice Problems • 1. __NaCl + __BeF2 __NaF + __BeCl2 • 2. __FeCl3 + __Be3(PO4) 2 __BeCl2 + __FePO4 • 3. __AgNO3 + __LiOH __AgOH + __LiNO3 • 4. __CH4 + __O2 __CO2 + __H2O • 5. __Mg + __Mn2O3 __MgO + __Mn

21. Types of Chemical Reactions Objectives: • 1. Identify a reaction as combination, decomposition, single-replacement, double-replacement, or combustion • 2. Predict the products of combination, decomposition, single-replacement, double-replacement, and combustion reactions

22. Classifying Reactions • For combination (synthesis: combination of parts into a whole) and decomposition, compare the number of reactants and products • For combustion, check for oxygen • For single- and double-replacement, look for a cation swap or the formation of a precipitate • Not all chemical reactions fit uniquely into only one of these classes

23. Combination Reactions • Two or more substances combine to form a single substance • Reactants are usually either two elements or two compounds • The product is always a compound (Can be an ionic compound or a molecular compound) • Some nonmetal oxides react with water to produce an acid (hydrogen ions in aqueous solution) • Some metal oxides react with water to form a base (hydroxide ions)

24. Worksheet Handout • Refer to the worksheet handout. • Identify the combination reactions.

25. Decomposition Reactions • A single compound is broken down into two or more products • The products can be any combination of elements and compounds • Most decomposition reactions require energy in the form of heat, light, or electricity • Extremely rapid decomposition reactions that produce gaseous products and heat are often the cause of explosions

26. Worksheet Handout • Refer to the worksheet handout. • Identify the decomposition reactions.

27. Single-replacement Reactions • Also called single-displacement reactions • One element replaces a second element in a compound • Can be compared to partners cutting in on each other at a dance: A person who is alone approaches a dancing couple and cuts in…that person replaces one member of the couple, who is now left alone

28. Whether one metal will displace another metal from a compound can be determined by the relative reactivities of the two metals. (Memorize the symbols and activity series of metals on page 217.) • A reactive metal will replace any metal listed below it in the activity series • Examples: Iron will displace copper from a copper compound in solution. Magnesium does not replace lithium from aqueous solutions of their compounds.

29. Refer to Table 8.2 on page 217 • Will magnesium displace zinc from a zinc compound in solution? • Will magnesium displace silver from a silver compound in solution? • Important Note: 1. Metals from lithium to lead will replace hydrogen from acids. 2. Metals from lithium to sodium will also replace hydrogen from water.

30. Single-Replacement (cont’d) • A nonmetal can also replace another nonmetal from a compound • This replacement is usually limited to the halogens (Group 7A): F2 (most activity) Cl2 . Br2 . I2 (least activity) • The activity of the halogens decreases as you go down group 7A on the periodic table

31. Worksheet Handout • Refer to the worksheet handout. • Identify the single-replacement reactions.

32. Double-replacement Reactions • Involves an exchange of positive ions between two reacting compounds • Often characterized by the production of a precipitate (ppt.-insoluble substance that “falls out” of a solution) • Product may be a gas that “bubbles” out of the mixture • Product may be a molecular compound, such as water

33. Worksheet Handout • Refer to the worksheet handout. • Identify the double-replacement reactions.

34. Combustion Reactions • An element or a compound reacts with oxygen, often producing energy as heat and light • Commonly involve hydrocarbons (compounds of hydrogen and carbon) • The complete combustion of a hydrocarbon produces carbon dioxide and water • If the supply of oxygen during a reaction is insufficient, combustion will be incomplete

35. During incomplete combustion, elemental carbon and toxic carbon monoxide may be additional products • Reaction between some elements and oxygen Example: Both magnesium and sulfur will burn by reaction with oxygen • Refer to worksheet handout. • Identify the combustion reactions.

36. Make a Chemistry Foldable • 1. Fold a sheet of notebook paper to the red margin line. • 2. Using scissors, cut the folded section into five equal parts. • 3. Label each section with the name of one of the five types of reactions. • 4. Open each flap and put in three characteristics and one example (include balanced equation). • 5. Write the title : Types of Chemical Reactions on the top of the sheet. Add your name and class. • 6. Use the chemistry foldable as a study guide.

37. Predicting Products of a Chemical Reaction • Recognize the possible type of reaction that the reactants can undergo • Some reactions do not fit any one of the five general types (Example: redox reactions) • Oxidation-reduction (redox) reactions will be discussed during the second semester OIL RIG – oxidation is the loss of electrons and reduction is the gain of electrons LEO the lion says GER – loss of electrons is oxidation and gain of electrons is reduction

38. Reactions in Aqueous Solution • Objectives: 1. Write and balance net ionic equations 2. Use solubility rules to predict the precipitateformed in double replacement reactions

39. Net Ionic Equations • Most ionic compounds dissociate, or separate, into cations and anions when they dissolve in water. • Refer to question #21 on the worksheet handout. Use this equation to answer #22 on the worksheet handout. • Write a complete ionic equation that shows dissolved ionic compounds as their free ions. • Eliminate ions that do not participate in the reaction by canceling ions that appear on both sides of the equation. These are called spectator ions.

40. Ions that are not directly involved in a reaction are called spectator ions. • Rewrite the equation, leaving out the canceled spectator ions. • Balance the atoms and the charges of the ions. (In this case, the number of atoms and the net ionic charge on each side of the equation is zero and it is therefore balanced.) • A net ionic equation indicates only those particles that actually take part in the reaction. • Record your answer to #23 on the worksheet handout.

41. Practice Problem • Write a balanced net ionic equation for the following reaction: Pb(s) + AgNO3 (aq)  Ag (s) + Pb(NO3)2 (aq) Answer: 1. The nitrate ion is the spectator ion. 2. The number of atoms balance, but the charges on the ions do not balance. 3. Place a coefficient 2 in front of Ag+ (aq) to balance the charges. 4. A coefficient of 2 in front of Ag (s) rebalances the atoms. 5. Pb(s) + 2Ag+ (aq)  2Ag (s) + Pb2+ (aq) is the balanced net ionic equation

42. Predicting the Formation of a Precipitate • Use the general rules for solubility of ionic compounds. • Examples: 1. Sodium nitrite will not form a precipitate because alkali metal salts and nitrate salts are soluble (Rules 1 and 2) 2. Rule 3 (Exceptions) indicates that barium sulfate is insoluble and therefore will precipitate.

43. Solubility Rules for Ionic Compounds

44. Practice Problem • Identify the precipitate formed and write the net ionic equation for the reaction of aqueous potassium carbonate with aqueous strontium chloride. 1. Write the reactants showing each as dissociated free ions. Balance the charges. 2. Using solubility rules, look at possible new pairings of cation and anion that give an insoluble substance. 3. Eliminate the spectator ions and write the net ionic equation.

45. Answer • 1. Reactants as dissociated free ions 2K+ (aq) + CO32- (aq) + Sr2+ (aq) + 2Cl- (aq) Charges must be balanced to equal 0. • 2. Of the two possible combinations, KCl is soluble (Rules 1 and 4) and SrCO3 is insoluble (Rule 5) • 3. The net ionic equation must be balanced for the number of atoms of each element and the charges on the ions. Sr2+ (aq) + CO32- (aq)  SrCO3 (s)