• Provide feedback on the completed practice worksheets and on student responses to the questions posed during the demonstration and instruction.

• Provide feedback and corrections as students work through the example problems during the guided instruction portion of the lesson and the class discussion.

Have students take out their science notebooks. Tell them that you are going to show two pictures and perform one demonstration. They are to answer one simple question regarding all three things at the end. Show students the following pictures, or similar pictures, and ask the corresponding questions (S-C-3-1_Gas Tank and Rust Pictures.doc).

Pour 10 ml of vinegar into a resealable plastic bag. Add 1 spoonful of baking soda to the bag. Close the bag immediately. Tell students what the reactants are as you add them to the bag. Wait for the bag to expand. Ask students, “Where did the gas in the bag come from? Where did the baking soda go?”

Allow students a few minutes to respond to the questions by writing in their notebooks. When students have completed their responses, ask for a few sample responses to each question. Then respond by saying, “The gasoline that you put in your car seems to disappear. The rust seems to appear from nowhere. The baking soda disappears, while a gas develops. These are all examples of chemical reactions. In the demonstration, the baking soda chemically reacts with the vinegar to produce carbon dioxide gas. That is why we saw the bag inflate. The atoms that make up the baking soda were not lost, nor were the atoms that make up the carbon dioxide gas created. In fact the atoms in baking soda (NaHCO₃) are rearranged to produce carbon dioxide (CO₂) and other products.” Explain that this is a closed system, in which the reactants and products are all within the sealed bag. If we measure the total mass of the reactants and the products, they will be the same. In an open system (i.e., open container), the CO₂ gas would have dispersed in the surrounding environment.

The reaction (two-step) is as follows:

(1) CH₃COOH + NaHCO₃  H₂CO₃ + NaC₂H₃O₂

(2) H₂CO₃  CO₂ + H₂O

Note: The last product of the first reaction is carbonic acid, which quickly decomposes into carbon dioxide and water. The CO₂ is what you see foaming and bubbling in this reaction. Vinegar is an aqueous solution of acetic acid.

Continue by saying, “As you can see, the atoms are not lost as the reaction proceeds nor are atoms created. They are merely rearranged. Let’s investigate why this is true and how chemists express this as they write chemical equations.”

“Let’s take a closer look at the reaction we just talked about. Look at the first reaction between vinegar and baking soda (sodium bicarbonate). Count the number of atoms of each element on the reactant side and then on the product side.” You may need to remind them that the reactants are on the left side of the arrow and that the products are on the right side. Guide them in forming lists like these:

Tell them to do the same thing for the second reaction. Their atom inventory should look like this:

At this point, say, “What do you notice about the number of atoms on the reactant and product sides of both reactions?” Students should recognize that the number of atoms of each element is the same. If students are not getting their atom inventory to match yours, remind them that the subscripts in the equation represent the number of atoms for the element before it only. For example, HCO₃ has 1 hydrogen atom, 1 carbon atom and 3 oxygen atoms. Chemists do not write 1 as a coefficient.. Say, “The fact that atoms are neither created nor destroyed in a chemical reaction is called the law of conservation of matter. There are similar laws as well that you may be familiar with such as the law of conservation of energy and the law of conservation of mass.” If there are still questions, put an example on the board that has 1:1 molar ratios, such as:

C + O₂  CO₂ or NaCl + K  KCl + Na

Optionally, give students colored jelly beans in a small cup with toothpicks. Designate a specific color for an element, such as red jelly beans for carbon and green jelly beans for oxygen. Have them construct models of the examples you did from the lesson. Remind them that they cannot add individual jellybeans, as that would be analogous to adding a subscript. They may only add complete sets (molecules/compounds) to satisfy the law of conservation of matter.

After you have formatively assessed that students are ready to move on (through interaction with the class and examples on the board), give them a more difficult problem. Say, “I want you to complete the same task but for a slightly more challenging reaction. Create an atom inventory or a list of the number of each atom on both sides of the equation.” Use the following reaction:

AlCl₃ + Na  Al + NaCl

Ask students what they notice about the atom inventory. Was the law of conservation of matter satisfied? (No). Tell students, “According to the equation as it is currently written, 2 chlorine atoms seemed to have disappeared as the reaction progressed. Remember that matter cannot be destroyed, according to the law of conservation of matter. Chemists use something called coefficients to address this discrepancy. Coefficients are large numbers that are placed in front of a molecule, element, or compound to satisfy the conservation law. After coefficients are used and there are equal numbers of atoms of each element on both sides of the equation, it is said to be a balanced equation.”

Say, “If we put a large 3 (a coefficient) in front of the NaCl, the equation would look like this” (write on the board):

AlCl₃ + Na  Al + 3NaCl

“If you were to create an atom inventory (count of the number of atoms of each element for both the reactant and product side), you will see that it now looks like this” (write on the board):

“Notice that the coefficient of 3 is placed in front of the NaCl. The 3 is distributed to all atoms in the compound (both the Na and Cl), which fixes the chlorine discrepancy, but creates a new discrepancy for sodium. It would seem easier to put the coefficient in the middle of the NaCl like this: Na3Cl, but that is not allowed. The ionic compound NaCl must be written that way due to the ionic bond that sodium and chlorine make. The coefficients cannot be sandwiched between two atoms in a compound. The fact that chlorine was ‘fixed’ but now sodium is wrong is not a problem. Another coefficient will be placed like this” (write on the board):

AlCl₃ + 3Na  Al + 3NaCl

“If you complete an atom inventory of the reaction, it should look like this” (write on the board):

“The law of conservation of matter has been satisfied and the equation is said to be balanced.” At this point, some students may ask what the coefficients mean and why chemists can just add them. It may help to compare a chemical equation to a recipe. Say, “Suppose you had the following recipe to make a s’more” (write on the board):

1 s’more (S)

2 graham crackers (G)

1 marshmallow (M)

1 piece of chocolate (C)

“If you were to write this recipe out as if it were a chemical reaction it may look something like this” (write on the board):

2 G + 1M + 1C  1S

“Notice how you would need two crackers to make one s’more? If you had only one cracker and you were following the recipe, you would not be able to make a complete s’more. Balancing a chemical reaction is similar in that reactions need specific ratios of reactants in order to occur.” Do a few more examples on the board.

Example 1

Unbalanced: K + Zn(NO₃)₂  KNO₃ + Zn

“The 2 coefficient is applied to both the potassium and the nitrate ion.”

Example 2

Unbalanced: Mg(OH)₂ + NaCl  MgCl₂ + NaOH

Remind students that they cannot add subscripts. You may need to review how binary ionic compounds are formed. Some students may want to add a subscript of 2 to the sodium hydroxide on the product side. Remind them that sodium hydroxide forms as follows:

Na⁺¹ + OH^-1  NaOH

“If you were to change that to Na(OH)₂, you would be indicating that sodium forms a +2 ion, which it does not.” Therefore, the balanced equation is written as:

Balanced: Mg(OH)₂ + 2NaCl  MgCl₂ + 2NaOH

Example 3

Unbalanced: CH₄ + O₂  CO₂ + H₂O

Tell students, “Note that there is an odd number of oxygen atoms on the product side. Additionally, the oxygen atoms on the product side came from two different molecules. You need to add them together when making your atom inventory.” Write on the board:

Balanced: CH₄ + 2O₂  CO₂ + 2H₂O

At this time, you may do more example problems on the board and/or hand out the Balancing Practice Worksheet (S-C-3-1_Balancing Practice Worksheet and KEY.doc).

Extension:

• Students who might need additional practice with balancing equations can use interactive online activities such as those listed in Related Resources.

• Students who may be going beyond the standards can balance more complex reactions such as the combustion of octane (a component of gasoline):

Unbalanced: C₈H₁₈ + O₂  CO₂ + H₂O

(Note: Balanced: 2C₈H₁₈ + 25O₂  16CO₂ + 18H₂O)