Fundamentals of General, Organic, and Biological Chemistry, 8e (McMurry)
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Chapter 5 Classification and Balancing of Chemical Reactions
Overview In Chapter 5, students will learn how to write, balance, and classify equations. They will learn about solubility guidelines and be able to write and balance equations for acidbase reactions. Redox is introduced here, and students should become proficient at oxidationreduction reactions as well as balancing net ionic equations.
Introduction for Instructors
• Chapter 5 is the introduction to chemical reactions, perhaps the meat of any course in chemistry. Thorough preparation here will give students confidence and the ability to achieve serious work and understanwding later.
• If you expect a certain number of students to drop out, they’ll drop during this section. If you expect most of your students to succeed, you may have to tempt, cajole, and threaten a little here. If you help them to develop a good understanding of these techniques, it will pay off later in student success and also in high course and instructor evaluations.
• Remember to review earlier information as you go through this material. It’s worthwhile to remember that most students don’t end up chemistry majors, so the material must have meaning to non-chemists as well, and this may be your greatest challenge. In some cases, simply stating the problem or the process of solution in different terms will help. Redox is tough, but beautiful. If you approach it step by step and with humor and understanding, your students are likely to reward you with hard work and appreciation.
Chapter Goals
• To be able to understand the meaning of, write and balance chemical equations
• To be able to recognize precipitation, acid-base, and redox reactions
• To be able to determine oxidation numbers of atoms and determine which substances are being reduced and which oxidized in a reaction
• To be able to recognize spectator ions and write net ionic equations
Lecture Outline
5.1 Chemical Equations
• A chemical equation tells us what’s happening as well as how much is happening.
• By convention, reactants are on the left, products on the right.
• The law of conservation of mass states that the sum of the masses of products must equal the sum of the masses of the reactants.
5.2 Balancing Chemical Equations
• Four steps are necessary in balancing an equation:
- Write an unbalanced equation, using correct formulas for reactants and products.
- Add coefficients to balance the numbers of atoms of each element (remember that subscripts cannot be changed).
- Check to make certain that the same numbers and kinds of atoms appear on both sides.
- Be certain that the coefficients can’t be divided by some small whole number.
Hands-On Chemistry 5.1
• Look up a recipe for your favorite cookies.
5.3 Precipitation Reactions and Solubility Guidelines
• Compounds containing most Group 1 elements such as sodium and potassium, or ammonium ions are usually soluble.
• Most halogen compounds are soluble except those containing silver, mercury(II), or lead (II) ions.
• Nitrates, perchlorates, acetates, and sulfates are usually soluble except those containing barium, mercury (II), or lead (II) ions.
Chemistry in Action
Kidney Stones: A Problem in Solubility
• Nucleic acids are degraded in the body to uric acid. The uric acid is excreted as sodium urate, which is only moderately soluble.
• Sodium urea is sometimes deposited in the joints of the big toe, causing a painful condition known as gout.
• Sodium urate may also be deposited in kidneys or in ducts as painful stones.
5.4 Acids, Bases, and Neutralization Reactions
• Acids and bases react to form salts and water.
• The neutralization process results in a nonconductive solution.
5.5 Redox Reactions
• In redox reactions, some species is always oxidized by an oxidizing agent, which is itself reduced in the process. The substance being oxidized brings about the reduction of the oxidizing agent and is itself referred to as a reducing agent.
• The reducing agent (species being oxidized) loses electrons and becomes more positively charged during the reaction.
• The oxidizing agent (species being reduced) gains electrons and becomes more negatively charged.
• Important redox reactions include bleaching, corrosion, combustion, and metallurgy as well as processes involved in respiration and recovery of metals from ores.
Chemistry in Action Batteries
• The charging and discharge of batteries typically involves oxidation and reduction.
• Batteries perform thousands of applications in our lives. Some may be implanted inside the human body.
5.6 Recognizing Redox Reactions
• Each atom has a certain oxidation number or state.
- Atoms of free elements have oxidation numbers of zero.
- Monoatomic ions have oxidation numbers equal to the charge of the ion.
- Atoms in molecular compounds usually have the same oxidation number as they would if they were present in monoatomic ions.
- The sum of the oxidation numbers in a neutral compound is zero.
5.7 Net Ionic Equations
• Acid-base, redox, and precipitation reactions may be represented by net ionic equations.
• Some ions appear as spectator ions present, but having no role in the reaction.
Lecture Demonstrations
• Chemical reactions provide a wealth of intriguing demonstrations. Start with a piece of magnesium ribbon. Light it with a burner and note evidence of a chemical reaction. It’s a memorable reaction and gives a simple equation to balance. (Magnesium reacts with oxygen to form magnesium oxide.) Care! Magnesium burns at a high temperature and can ignite other flammables. The light is intense, but probably not harmful if one doesn’t look directly at it. There’s another, hidden, reaction: at the same time, magnesium reacts with nitrogen in air to form magnesium nitride.
• A series of reactions can be used as follows:
- Add 10 mL of 0.1 M NaCl to 10 mL of 0.2 M AgNO3 in a beaker. Note the formation of silver chloride precipitate. Ask students to write both chemical equations and net ionic equations for the reaction.
- Filter the solution, collecting the AgCl precipitate on filter paper. (Here’s a chance to introduce the terminology of filtration.) Using a stirring rod or rubber policeman, spread the precipitate into a thin layer. Cover half the precipitate with the base of a ring stand. Ignite a piece of magnesium ribbon and hold it over the precipitate. After the ribbon has burned, remove the ring stand and note that the exposed precipitate has darkened (decomposed into silver grains and chlorine gas). Note the similarity to reactions in photographic film when a photo is taken. Have students write balanced equations.
- Inform students that chemists seldom use exact portions of reactants. Ask which reactant (NaCl or AgNO3) was in excess in the original reaction. Test small portions of the filtrate to find the answer (AgNO3). Place the filtrate (with some distilled water added for dilution) in a beaker and add a length of copper wire. After a few minutes, note that the copper color has been replaced by a grey color on the wire, which becomes thicker and finally forms beautiful crystals of silver.
- Following this demonstration in one class, a student searched the city for sources of silver solutions and found enough silver in X-ray and photo labs to pay the remainder of her tuition throughout college. Silver ion is a dangerous pollutant, and most photo labs are required to prevent discharge of silver or other heavy metals into the local sewage system.
- If you start with a weighed amount of magnesium ribbon, students can calculate the expected yield of magnesium oxide.
• For acid-base reactions, use dilute solutions of NaOH and HCl.
- Note the use of indicators.
- Note the heat of reaction which is very small in this case but may be considerable in industrial applications.
- Use some carbonate or bicarbonates either in solid form or in solutions. Allow HCl to react with them. Note the production of CO2.
- Display a box of baking soda and ask why we place it in refrigerators.
• Illustrate redox reactions by burning, as with the Mg ribbon above. Also burn wood splints and note which substances are being oxidized or reduced, and identify the oxidizing agent.
• Add a small amount of Cl2 water to a solution of KBr or KI. Shake the resulting solution with a small amount of methylene chloride or trichloromethane. Note the characteristic colors of elemental bromine and iodine. Write equations and identify oxidizing agents.
• Cut an apple and note the browning due to polyphenol oxidase action on the tissue. This is a protective action, producing substances which help prevent insect larvae from growing in the apple tissue. Place a slice in ascorbic acid solution to see if browning is retarded.
• Warn students to cover their ears, and ignite a balloonful of hydrogen gas (use a meter stick with a burning splint on the end). For a much louder explosion, use a balloon containing a stoichiometric mixture of hydrogen and oxygen. Don’t do this in a small lecture hall, and do warn the students before doing this! They’ll want to see the second explosion, even though they expect it to be louder.
• Some instructors use the thermite reaction to illustrate redox. A safer method is to strike together two very rusty iron balls (such as cannonballs) wrapped with aluminum foil. Thermite reactions occasionally go awry with nasty results, but their use in industry allows the welding of heavy duty amounts of steel under adverse conditions.
• If you start with a weighed amount of magnesium ribbon, students can calculate the expected yield of magnesium oxide.
- For acid-base reactions, use dilute solutions of NaOH and HCl.
- Note the use of indicators.
- Note the heat of reaction that is very small in this case but may be considerable in industrial applications.
- Use some carbonate or bicarbonates either in solid form or in solutions. Allow HCl to react with them. Note the evolution of CO2.
- Display a box of baking soda and ask why we place it in refrigerators.
- Illustrate redox reactions by burning, as with the Mg ribbon above. Also burn wood splints and note which substances are being oxidized or reduced, and identify the oxidizing agent.
Teaching Tips
5.2 It may help to emphasize the fact that matter can’t be created or destroyed in an ordinary reaction, so the numbers of atoms must be equal on both sides of an equation.
5.2 Some students may have been taught to use fractional coefficients in balancing equations.
5.2 Many errors in balancing can be avoided if students recognize that spectator polyatomic ions often appear unchanged on both sides of an equation.
5.2 Remind students that they can change the coefficients in balancing, but not the subscripts.
5.2 Monotonous as it seems, the key to understanding Avogadro’s number is working lots of problems.
5.2 It may help students understand mole ratios to point out that gases react in discrete molar volumes such as two parts hydrogen and one part oxygen.
5.2 Students usually expect chemical reactions to occur between exact molar quantities of reactants. They seldom realize that in real life, one reagent is usually limiting.
5.2 Students often don’t realize that reactions are seldom 100% complete and that the actual yield must be measured for each reaction.
5.3 It may help to list several general types of insoluble combinations of compounds.
5.4 When first balancing acid-base neutralizations, some students prefer to consider water as HOH and to view the –H as coming from the acid and the –OH as coming from the base.
5.5 It may help to point out that oxidation and reduction always occur together.
5.5 The simplest way to recognize redox reactions is to look for a change in charge or oxidation state of an atom.
5.5 It may help students to label each of the parts of a redox reaction such as the substance being oxidized, that being reduced, and the oxidizing and reducing agents.
5.5 Students often don’t understand that charges on ions in ionic formulas can be shown for convenience, but need not be shown unless called for.
5.5 It will probably help to emphasize that some ions are spectators and do not change oxidation states during reactions. The start of successful balancing is to assign oxidation states.
5.6 In balancing redox reactions, it’s helpful to underline each substance changing in oxidation state so that other substances don’t confuse the issue.
5.7 If a student learns to use net ionic equations, they can balance very, very difficult equations.
Group Problems
5.73 (a) Nitrogen commonly forms nitrogen monoxide (NO) and nitrogen dioxide (NO2) during incomplete combustion in power plants. N2 + O2 → 2 NO N2 + 2 O2 → 2 NO2 Reaction with sulfur commonly produces sulfur dioxide (SO2). S + O2 → SO2
(As the oxides of nitrogen vary with combustion conditions, they are often represented as NOx.)
(b) NO2 + OH– → HNO3, H2O + SO2 → H2SO3, H2O + SO3 → H2SO4
(c) The reactions in part (a) are redox because uncombined elements with a zero charge become incorporated into compounds and have a new oxidation state when they do so. For example, N is oxidized as it goes from 0 to +4 and oxygen is reduced as it goes from 0 to
2. The reactions in part (b) both appear to be Lewis acid-base reactions.
5.74 (a) Mg(OH)2 (b) Mg(OH)2(s) + 2HCl(aq)→ MgCl2(aq) + H2O(l)
5.75 The “HCl” means that the compound is presented as a hydrochloride salt. It’s formed by reacting the neutral compound with HCl and makes the product more soluble in water or body fluids.