What makes a fizzy drink fizzy? Do we have to have an expensive gadget to do it, or can we make a drink fizzy without one?

Watch the video below and answer the question that follows.

Watch your instructor perform a demonstration and answer the questions below.

Your class will discuss your answers to how both the SodaStream and your instructor were able to add bubbles to a beverage.

Imagine you have a two identical bottles of water, both are 1 liter bottles. You plan to use the SodaStream to carbonate one of the bottles, and the method your instructor showed your class to carbonate the other. What type of data might you acquire before and after carbonation in order to compare methods?

You will perform two different experiments in order to determine what happens to matter (mass, because all matter has mass) during a chemical reaction.

Mg (s)	+	HCl (aq)		MgCl2 (aq)	+	H2 (g)
Reactants				Products

1. Place a weighing boat on your balance and zero it out.
2. Measure out 0.2 g of magnesium in the weighing boat.
3. Record the mass of a 125 mL Erlenmeyer flask.
4. Measure out 15 mL of 0.5 M HCl in a 25 mL graduated cylinder.
5. Pour the 15 mL of 0.5 M HCl into the flask and record its mass.
6. Subtract the mass of the flask itself to determine the mass of the HCl only.
7. Add the mass of the Mg and the HCl, record this total mass below. (This is the total mass of your reactants.)
8. Separate a small square of parafilm from its backing paper and record its mass.
9. Pour the Mg from the weighing boat into the flask full of HCl (still on the balance), then immediately cover the flask with parafilm so no gas is allowed to escape.
10. Place the flask (still covered with parafilm) back onto the balance.
11. Subtract the mass of the flask + the parafilm from the mass on your balance, then record that mass. (This is the total mass of your products.)
12. Record any observations of what occurred during the chemical reaction below. Observations include describing what happened during the reaction and noting differences between the reactants and products.
13. Dump the contents of the flask down the appropriate drain, rinse it out, and place it on the drying rack.

NaHCO3 (s)	+	CH3COOH (aq)		NaC2H3O2  (aq)	+	H2O (l)	+	CO2 (g)
Reactants				Products

1. Record the mass of a 125 mL Erlenmeyer flask and record the mass of your balloon.
2. Measure 20 mL of 1.0 M acetic acid (CH₃COOH) in a 25 mL graduated cylinder. Pour this into the empty 125 mL Erlenmeyer flask.
3. Place a weighing boat on your balance and zero it out.
4. Measure out 0.6 g of baking soda (NaHCO₃) in the weighing boat, then pour it carefully into your balloon.
3. Place the baking soda balloon and the acetic acid flask on the balance. Subtract the mass of the flask and the balloon from the mass on your balance and record that mass. (This is the total mass of your reactants.)
4. Carefully attach the balloon to the flask making sure not to empty any of the baking soda into the flask in the process!
5. Once the balloon is secure on the flask, shake the baking soda into the flask so it reacts with the acetic acid. Wait until all of the material seems done reacting.
6. Record any observations of what occurred during the chemical reaction below. Observations include describing what happened during the reaction and noting differences between the reactants and products.
7. Record the total mass after the reaction by again subtracting the mass of the flask and the balloon from the mass on your balance. (This is the total mass of your products.)
8. Dump the contents of the flask down the appropriate drain, rinse it out, and place it on the drying rack. Return the balloon to your instructor.

The overall reaction for experiment 1 was expressed as: Mg (s) + HCl (aq) → MgCl₂ (aq) + H₂ (g)

The overall reaction for experiment 2 was expressed as: NaHCO₃ (s) + CH₃COOH (aq) → NaC₂H₃O₂ (aq) + H₂O (l) + CO₂ (g)

Look carefully at how both of these chemical reactions are written and then think back to the actual experiments you completed in order to answer the questions below.

Grab a large whiteboard and make a table of your results from both experiments. Set up your table as shown below.

What happens to mass during a chemical reaction? Review your class data from the two experiments and explain what happens to the overall mass of the reactants and products.

The law of conservation of mass states that matter (mass) cannot be created or destroyed. Therefore the total mass of the reactants in a chemical reaction must be equal to the total mass of the products. How then is it possible to"lose" weight?

Watch the video below and answer the question that follows.

In the last lesson, you watched while your instructor carbonated a beverage in front of you without a SodaStream.

You will need the following resources to complete this assignment.

• Calculating Molar Mass

When biologists talk about moles, they’re usually referring to the tiny, grey rodents that dig underground to find tasty earthworms. When chemists talk about moles they’re usually referring to something entirely different. The term ‘mole’ represents a number in chemistry the same way the word ‘dozen’ represents 12 of something to a baker. In this case, one mole represents 6.02 x 10²³, a humongous number!

To help you and any bakers or hungry biologists reading this get a sense of just how many things are in one mole, we can use an analogy with something that is often found in dozens (Figure 1). One yeast-raised doughnut weighs approximately 40 grams. One dozen yeast-raised doughnuts weigh about 500 grams. One mole of yeast-raised doughnuts weighs about 301,000,000,000,000,000,000,000,000 grams—more than four times the mass of the moon. Imagine the whole moon made of doughnuts, yummy...

Figure 1: A dozen yeast-raised donuts

Obviously, the mole is not a term we need for most things in daily life. Instead of being used for things we encounter in daily life, the mole is used by scientists when talking about enormous numbers of particles like atoms or molecules. Atoms and molecules are very tiny things. A drop of water the size of the period at the end of this sentence would contain 10 trillion water molecules. Instead of talking about trillions and quadrillions of molecules (and more), it's much simpler to use the mole. However, the mole does more than represent a big number: it provides a key link for converting between the number (amount) of a substance, and its mass.

The International Committee for Weights and Measures defines one mole as the number of atoms in exactly 12 grams of carbon-12 (Figure 2). Experiments counting the number of carbon-12 atoms in a 12-gram sample have determined that this number is 6.02214076 x 10²³. Regardless of whether the substance is carbon-12 or doughnuts, one mole represents the same number of each of these things.

Figure 2: Carbon-12

Scientists have then defined the molar mass of a substance as the mass of 6.02214076 x 10²³ units of that substance. So, the molar mass of yeast-raised doughnuts is 301,000,000,000,000,000,000,000,000 grams. With doughnuts, this is not very useful. However, it is quite useful if we apply it to other substances, especially elements. By standardizing the number of atoms in a sample of an element, we also get a standardized mass for that element that can be used to compare different elements and compounds to one another. Carbon's molar mass is 12.01 grams, which represents the combined mass of 6.02 x 10²³ carbon atoms. However, other elements have different molar masses; for example, 6.02 x 10²³ sulfur atoms have a mass together of 32.06 grams, which is sulfur's molar mass.

Along with telling us the mass of one mole of an element, molar mass also acts as a conversion factor between the mass of a sample and the number moles in that sample. For example, 24 grams of carbon atoms would be equal to two moles since 24 grams divided by the mass of one mole (12.01) equals 2. Further, Avogadro’s number acts as the conversion factor for converting between the number of moles in a sample and the actual number of atoms or molecules in that sample. Extending our example, two moles of carbon atoms contains 2 times 6.02 x 10²³ atoms, which equals 12.04 x 10²³ atoms, which can be written as 1.204 x 10²⁴ atoms.

Look on a periodic table in order to answer the following questions.

Your instructor will provide you with a handout so you can practice calculating molar mass.

When are grams an appropriate measurement, when are moles necessary, and how can we convert between the two units?

You will need the following resources to complete this assignment.

• HW Gram Mole Conversions

In the following activity, you will measure out a mole, in grams, of several elements in order to see how much a mole really is. Make sure your group has a periodic table, you're going to need it.

Create a 1 mole sample of aluminum using aluminum pop cans. It should be close to a mole, but it does not need to be exact.

Create a 1 mole sample of iron using paperclips. It should be close to a mole, but it need not be exact.

Create a 1 mole sample of water. It should be close to a mole, but it not be exact. Measure out the amount of water needed in the graduated cylinder. Place the small 50 mL beaker on the scale and zero it. Then pour the water in the small beaker and record the mass of the water. (Hint: the density of water is 1 g/mL. Density is mass/volume. Get the molar mass and solve for the proper volume).

Create a 1 mole sample of salt (salt is sodium chloride = NaCl). Do not place salt directly on the scale, measure it inside of a weighing boat provided. It should be as close to a mole as possible, but it need not be exact.

In the conservation of mass labs we performed in class, all measurements were taken in grams.

Chemical equations can be used to represent what happens on either the microscopic or macroscopic scale.

C (s) + O₂ (g)  CO₂ (g)

This equation can be read in either of the following ways.

1. If, or when, solid carbon reacts with gaseous oxygen, one atom of carbon and one molecule of oxygen are consumed for every one molecule of gaseous carbon dioxide produced.
2. If, or when, solid carbon reacts with gaseous oxygen, one mole of carbon and one mole of oxygen are consumed for every one mole of gaseous carbon dioxide produced.

Your instructor will provide you with a handout so you can practice converting between grams and moles.

How do we know how many moles of something we have if the substance is dissolved in water? You have used many solutions in chemistry this year. The concentration of those solutions is always indicated with an "M". The "M" stands for something called molarity.

In Lesson 2: What Happens to Mass During a Chemical Reaction, you used both hydrochloric acid and acetic acid. Both of these chemical compounds were followed by (aq), meaning in an "aqueous" solution. This means that the hydrochloric acid and the acetic acid were in a solution of water.

You will need the following resources to complete this assignment.

• Molarity Calculations

Chemical reactions are typically described as so many moles of compound A reacting with so many moles of compound B to form so many moles of compound C. When we determine how much of a reactant to use, we need to know the number of moles in a given volume of the reactant. Percent solutions only tell us the number of grams, not moles. A 100 mL solution of 2% NaCl will have a very different number of moles than a 2% solution of CsCl. So we need another way to talk about numbers of moles.

Chemists need the concentration of solutions to be expressed in a way that accounts for the number of particles that react according to a balanced chemical equation. Since percentage measurements are based on either mass or volume, they are generally not useful for chemical reactions. A concentration unit based on moles is preferred. The molarity (M) of a solution is the number of moles of solute dissolved in one liter of solution. To calculate the molarity of a solution, you divide the moles of solute by the volume of the solution expressed in liters.

Note that the volume is in liters of solution and not liters of solvent. When a molarity is reported, the unit is the symbol M and is read as “molar”. For example a solution labeled as 1.5 M NH₃ is read as “1.5 molar ammonia solution”.

Your instructor will provide you with a handout so you can practice a variety of molarity calculations.

In this lesson, you will be asked to focus on the amount of one product in a reaction and how the amount of that product changes relative to the amounts of each reactant used. The goal of the lesson is to understand the best possible combination of reactants to use in a reaction so none of the reactants go to waste.

Observe as your instructor demonstrates the reaction of magnesium metal and hydrochloric acid to produce hydrogen gas (magnesium chloride is also produced in solution). Your instructor will use varying amounts of magnesium metal and hydrochloric acid in each flask. These amounts have been measured in moles and are represented in the table below.

Experiment with the simulation by varying amounts of the two reactants. Be sure to hit "setup" after adjusting amounts or the change will not be reflected in the simulation. After you feel you have learned all the features of the simulation move on to answer the questions below:

The simulation below models the same reaction your instructor demonstrated, but now you have control over the amounts of both reactants being used.

Directions:

1. Adjust the slider bars to vary the amounts of each reactant (Mg and HCl) as instructed in the table below.
2. Then press "setup". The values you've selected for each reactant will now appear in the box at right. Moles of magnesium are represented by the grey circles and moles of hydrochloric acid (HCl) are represented by the green circles in the blue rectangle.
3. Now press "go" and observe how much H₂ gas is produced with each combination of reactants. All amounts are again measured in moles.

Now you will use the same simulation, but this time you have control over the amounts of each reactant you will use.

Directions:

1. Adjust the slider bars to vary the amounts of each reactant (Mg and HCl) and record those amounts in the table below.
2. Then press "setup". Again, the values you've selected for each reactant will now appear in the box at right. Moles of magnesium are represented by the grey circles and moles of hydrochloric acid (HCl) are represented by the green circles in the blue rectangle.
3. Now press "go" and observe how much H₂ gas is produced with each combination of reactants. All amounts are again measured in moles.

This simulation will help you to visualize the process of balancing equations. You will begin with the Introduction where you will go through the process of balancing three different equations. After you finish the Introduction, you will proceed to the Game which is broken down into three levels. Each level has five equations to balance.

In the previous lesson, you observed a chemical reaction with magnesium (Mg) and hydrochloric acid (HCl) that produced magnesium chloride (MgCl₂) and hydrogen gas (H₂). This chemical reaction is represented as a chemical equation below.

Mg	+	HCl		MgCl2	+	H2
Reactants				Products

In order to use all of the starting magnesium and produce the maximum amount of hydrogen gas possible, a 1:2 ratio of magnesium to hydrochloric acid is needed. This is represented in the chemical equation below by adding a coefficient in front of the HCl.

Coefficients tell you how many moles of each chemical are needed for the reaction to take place most efficiently. If there is no coefficient present, then it is assumed to be 1. For the reaction above, you will need 1 mole of Mg with 2 moles of HCl to yield 1 mole of MgCl₂ and 1 mole of H₂.

In Lesson 2, "What Happens to Mass in a Chemical Reaction?" you learned about the law of conservation of mass. Recall that the law of conservation of mass states that matter (mass) cannot be created or destroyed. Therefore the total mass of the reactants in a chemical reaction must be equal to the total mass of the products. Chemical reactions represent different quantities or amounts of reactants and products using coefficients. Adding the coefficients insures equal masses on both sides of the equation; this process is known as balancing an equation.

The chemical reaction shown below is unbalanced, that is, it does not follow the law of conservation of mass.

Mg

+

HCl

MgCl2

+

H2

 

We can determine the reaction is unbalanced by performing an atom inventory. An atom inventory is counting different types of atoms on both sides of the equation to see if they are equal. An atom inventory for the equation is shown below.

Mg	+	HCl		MgCl2	+	H2
1			Mg	1
1			H	2
1			Cl	2

 

In order to balance the equation, we added a coefficient of 2 in front of the hydrochloric acid (HCl). The updated atom inventory is shown below.

Mg	+	2 HCl		MgCl2	+	H2
1			Mg	1
2     1			H	2
2     1			Cl	2

On the next screen, you will use a simulation in order to balance several chemical reactions. All reactions will be presented with no coefficients (understood to be a coefficient of 1). You will be able to adjust coefficient numbers up and down as you observe scales that represent the atom inventory for all atoms present in the reaction. The reaction is balanced when all the scales are balanced.

1. Select Game
2. Choose Level 1 to start
3. Each level has five equations for you to balance. When you think you have balanced the equation correctly, click the "Check" button" in the bottom center of the screen. You will be awarded points if you get the equation correct. You only get two attempts per equation.
4. Continue on to complete both Levels 2 and 3.
5. After you complete each level, enter your score in the appropriate question box for that level below.

The Law of Conservation of Mass states: that mass is neither created nor destroyed in any chemical reaction. Therefore balancing equations requires the same number of atoms on both sides of a chemical reaction. The number of atoms in the reactants must equal the number of atoms in the products.

You will use molecular model kits to help you balance equations. You will begin by building reactants then disassembling them to build the products. If you have any atoms of the reactants left, you must go back and adjust the number of reactant molecules in such a way that none of the atoms are left over after you build the products.

After completing the activity you will be provided with basic notes on a preferred element order to balance equations.

You will need the following resources to complete this assignment.

• Balancing Chemical Equations With Molecular Model Kits
• HW Balancing Practice
• Notes on Balancing Equations

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