Final Exam Project - Crime Scene Mystery

For the final exam project, pairs of students will work to perform a specific analysis of evidence collected from the crime scene and compare the results with established knowns.  The project assignments are:

Analysis                      12C                          12A                                        12B
1. DNA Analysis          Max & Anthony         Andrea, Gabe, & Natalia        Anna & Yasmynn
2. Blood Typing           Ryan & David          Stacey & Melissa                   Katie &Jenna
3. Fingerprinting          Holly & Annie           Shannon & Pooja                   Melissa & Ashton
4. Chromatography    Claire and Gabbi      Joe & Mikah                            Domo & Risha
5. ID of Unknown        Lizzy & Travis          Daniel & Kevin                       Sara & Breanna
6. Shoe Impression    Megan & Kayla         Gus & Olivia                           Alex & Michael
7. Accelerant              Thom & Rocky         Brice & Dionysus                   Lukas & Philip

You are instructed to become an expert in your field.  You should thoroughly understand your analysis and be ready to run experiments by Thursday, May 12th.  Experiments will be run from Thursday, May 12th to the following Tuesday, May 17th.  Your team should be working on a Powerpoint presentation and a handout regarding your methodology and experimental results (this is major part of your final exam grade).  Team presentations will begin Thursday, May 19th and be completed by Monday, May 23rd.  Final exams are May 24th through May 26th.

Finding the Ka and Kb of a Weak Acid and Weak Base - Lab Report

Lab Report Instructions (Due Monday, May 16th):

Purpose: Discover the equivalence point and derive the equilibrium constants of a weak acid (Ka) and a weak base (Kb).

Materials - This is different from the handout given our available supplies; draw a picture of the apparatus.

Procedure - This is different from the handout given our available supplies.

Data - This is your raw data in data tables.

Analysis - This will include your graphs (volume of strong acid or base added v. pH); identify the equivalence point from the graph; calculate the Ka and Kb using the equivalence point.

Conclusion - Restate the purpose; summarize the analysis and draw conclusions (address the characteristics of weak acids and bases in comparison to the behaviors of strong acids and bases); possible sources of error; possible extensions for the experiment.

Adding a Strong Base to a Weak Acid

Problem: Calculate the pH of a solution when 45.0mL of 0.100M NaOH (a strong base) is added to 50.0mL of 0.100M CH3COOH (a weak acid).

HINTS:
We found that the number of mols of NaOH is 4.50 x 10^-3 and the number of mols of CH3COOH is 5.00 x 10^-3.

Since each mol of NaOH causes the weak acid to dissociate, the number of mols of CH3COO- is equal to 4.50 x 10^-3 and the remaining number of mols of CH3COOH is 5.00 x 10^-4.

We can find the Molarity of CH3COO- and CH3COOH by dividing the moles by the total volume, 0.095L.

We discussed in class that as the concentration of the conjugate base, CH3COO-, increases, it will react with water and release OH- ions.  This results in the solution becoming more basic.  Please note that the Ka accounts for this!!  So, we simply use the equation for Ka to solve for [H+].

Extra Credit Assignment - Powerpoint about Article!

Choose an article to read from a credible science journal or magazine.  Create a Powerpoint presentation to summarize the information presented in the article.  The slides should address the following questions:

1. What was researched/discovered?
2. What method(s) was used in the investigation?
3. What evidence was found to support/refute their hypothesis?
4. Include a citation for your article.

INSTRUCTIONS FOR USING GOOGLE DOCS:

• You must have a Google account in order to use Google docs.  

• You may create your presentation in MS Powerpoint or in Google Docs (under the drop-down labeled "Create").  

• If you create your presentation in MS Powerpoint, you must "Upload" your *.ppt file to your Google account.  

• Then, you need to select (click on it in your list of document) your presentation within Google Docs.  Once selected, look over to the right side of the screen to see the options for this file.  Next to "Share" click on "Settings".  In the new window, "Add" the following - tpachemistry@gmail.com.  This will allow me to pull up your presentation in the classroom through my account.

Chemical Equilibrium - Reactions and Solubility

In preparation for our exploration into acids and bases, it is imperitive that you have a strong understanding of chemical equilibrium.  We have investigated reversible reactions involving ions in solution as well as solids dissolving into solution.  Please review the labs, your notes, and the reading in the textbook regarding chemical equilibrium.

HW:
(1) Complete the BONUS question from the Chemical Equilibrium Quiz (due Monday); and
(2) Complete a write-up (minimum of two-pages including drawings) regarding the following (due Tuesday):

• What is a reversible reaction? What does it mean for a reaction to reach equilibrium?  How is the reaction rate related to concentration?
• Phosgene is the chemical compound with the formula COCl2. This colorless gas gained infamy as a chemical weapon during World War I. It is also a valued industrial reagent and building block in synthesis of pharmaceuticals and other organic compounds.Consider the following endothermic reaction:
  
  COCl2  <---> CO(g) + Cl2(g)          Keq = 170
  
  Write the formula for the equlibrium constant for this reaction.  If the concentrations of CO and Cl2 are each 0.15M and the concentration of COCl2 is 1.1 x 10^-3M, has the reaction reached equilibrium?  How do you know? If not, in which direction will the reaction proceed?
• What is Le Chatelier's Principle?
• For the above reaction, explain how the following changes would shift the equilibrium position of the reaction:
• Increasing and decreasing pressure - explain in text and draw particle diagrams
• Increasing and decreasing reactant concentrations - explain in text and draw partile diagrams
• Increasing and decreasing temperature - explain in text
• Explain the processes of dissolution and precipitation.  What does it mean for a solution to be unsaturated vs. saturated?  What does it mean for a solid to reach solubility equilibrium?
• Consider the following chemical equation:
  
  Fe(OH)3 (s) <---> Fe3+(aq) + 3 OH-(aq)    Ksp = 6.3 x 10^-38
  
  Write the formula for the equlibrium constant for this reaction.
• For the above reaction, explain how the following changes would shift the equilibrium position of the reaction:
• Increasing and decreasing reactant concentrations - explain in text
• Increasing and decreasing temperature - explain in text
• Adding HCl to the solution - explain in text
• Adding NaOH to the solution - explain in text

Practice Problems - Dissolving More with Acid

Today, we calculated the volume of acid needed to dissolve given masses of Mg(OH)2(s) and CaF2(s).  The problems are listed below.  Be sure that you are comfortable with the steps involved in solving them.

1. Ten (10.0) grams of Mg(OH)2(s) is combined with distilled water to yield a volume of 50.0ml of solution.  What volume of 3.00M HCl must be added to dissolve all of the
Mg(OH)2(s)?

2. Twelve (12.0) grams of CaF2(s) is combined with distilled water to yield a volume of 60.0ml of solution.  What volume of 2.00M HCl must be added to dissolve all of the CaF2(s)?

We will have a quiz on solubility equilibrium on Thursday.  You should be able to:

1. Given a list of Ksp values, order substances from least to most soluble.
2. Calculate solubility (g/L) from Ksp.
3. Predict the effect of acid on solubility.
4. Calculate the volume of acid needed to completely dissolve a given mass of solid (assuming that lowering the pH will increase solubility).
5. Predict the effect of a common ion on solubility.
6. Calculate the effect of a common ion on solubility.

Math with Solubility Equilibria

On Friday, we observed a demonstration of milk of magnesia being combined with HCl.  Complete the following in terms of your observations during the demonstration and understanding of solubility equilibrium:

1. Describe the demonstration performed on Friday.
2. What does the Ksp = 1.8 * 10^-11 tell you about Mg(OH)2?
3. Explain the processes which resulted in the observed color/transparency changes.
4. Did the addition of 1.0M HCl change the Ksp of Mg(OH)2?
5. Calculate the solubility (g/L) of Mg(OH)2 in distilled water.

Early Release Tomorrow - April 8, 2011

In memorial of Rachel Ewer, TPA will be observing a half-day schedule tomorrow, April 8th.  All students will be released at 12:15pm. 

Information for Rachel's Memorial Service:

Saturday, April 9th - 2:00pm

The Church of Jesus Christ of Latter-day Saints
1100 N. Cooper Rd
Gilbert, AZ  85233

The Ewer family is requesting that cookies be brought to the memorial service to be shared at a time of fellowship following the services.  If you are able to bring cookies, please notify Mrs. Moffitt, kmoffitt@tempeprep.org.  Thank you.

Solubility Product Constant

The solubility-product constant (or simply solubility product) is the equilibrium constant expression for the equilibrium between a solid and an aqueous solution of its component ions.  Solids do not appear in the equilibrium constant expression for heterogeneous equilibria.  Therefore, Ksp is equal to the product of the concentrations of the ions raised to the power of their stoichiometric coefficients.

CxAy  <-->  xC  +  yA

Ksp = [C]^x [A]^y

HW: On a separate sheet of paper, complete the following problem.

Problem -
(a) Write the chemical equation for the addition of CaF2 to distilled water.
(b) Write the equation for Ksp for the addition of CaF2 to distilled water.
(c) The Ksp for CaF2 at 25 degrees C is 3.9 x 10^-11.  What is the solubility of CaF2 in grams per liter?
(d) What affect would the addition of Ca(NO3)2 have on the reaction?
(e) What affect would the addition of HCl have on the reaction?

Solubility Equilibria

The dissolution of solids is a reversible reaction.  We can calculate an equilibrium constant (Ksp) for solubility reactions based upon solubility measurements.  We can also calculate solubilities using Ksp.  What factors can influence the solubility of a substance?

HW: Read pages 561-569, 580 in the textbook.

Exploring Equilibrium - It Works Both Ways

Today, we began our investigation into factors which affect equilibrium in the "Exploring Equilibrium - It Works Both Ways" Lab.  We will complete the discussion of the lab tomorrow.  I will be collecting the journals on Monday.  In addition, we will be having a quiz on equilibrium on Wednesday.

HW: Complete the following journal entry -
Write a summary lab report for the "Determining the Keq of a Reaction" lab (purpose, materials, procedure - part 1 for the reference solutions and part 2 for the experimental solutions, data, analysis, and conclusion).  Include an explanation of the Beer-Lambert law which was used to determine the concentration of iron thiocyanate in solution as well as an explanation of the concentrations used to produce the reference solutions.

Scientific Theories

So, where do you think that you fall in the spectrum of Super-realism and Realism?  Or do you see yourself more as an Instrumentalist or Descriptivist?  In reviewing your notes, write down some pros and cons of each viewpoint (from your point of view).  Then, complete your journal entry in which you describe your stance on the formation of scientific theories.

HW: Journal entry due Friday - (1) What is a scientific theory? (2) What is an empirical law? (3) How are they related? (4) What are observables and unobservables? (5) Summarize the four philosophical viewpoints discussed in class and in the reading?  Include a discussion of how each views observables/unobservables, theoretical statements, and empirical statements. (6) Write a paragraph describing your position on scientific theories in relation to the four philosophical viewpoints. (7) Explain the process of theory formation. (8) Explain the two ways in which empirical laws are generated.

Introducing ........ Acids and Bases!

Acids and bases are a type of reversible chemical reaction.  Instead of using absorbance to decect changes in concentrations, indicators made up of organic molecules are added to acid/base solutions to produce a color change based upon H+ and OH- concentrations.

Tomorrow, we will complete our discussion of scientific theories.  Please be sure to have the reading completed for a lively discussion.

HW: Complete the Chemical Equilibrium Notes in your journal by Wednesday. 
1. What is a reversible reaction?
2. How do the rates of the forward and reverse reactions change throughout a reaction?
3. Write the equilibrium constant equation.  How does it relate to chemical equilibrium?
4. What can you conclude about the magnitude of Keq?
5. What is the reaction quotient (Q)? How does it relate to equilibrium?
6. State Le Chatlier's principle.
7. Predict the shift in equilibrium for the following reaction if (a) CO is added, (b) As4 is removed, (c) pressure is increased:

As4O6 (s) + 6C (s) ---> As4 (g) + 6CO (g)

8. What is the Haber process?  Why is an understanding of chemical equilibrium essential to the industrial production of ammonia?
9. Summarize how changes in concentration, pressure, and temperature affect reversible reactions at equilibrium.  Give examples of each.

I will post an additional journal entry for (1) the "Determining the Keq of a Reaction" lab and (2) scientific theories by mid-week which will be due by Friday.

Chemical Equilibrium

Today, we focused on the Keq constant.  No new information was covered during class due to many absent students.

HW: Read Chapter 16 in the textbook.

Determine Keq of a Reaction

Over the past two days, we completed the lab to determine the Keq of a reaction.  Be sure to complete the handout provided for the lab.  Based upon the equilibrium constant equation, what do you think the value of Keq reveals about the reaction?

HW: Complete the Keq of a Reaction handout.

What is a Theory?

What is a theory? How do we deal with observables and unobservables in science and philosophy of science? Explain the four different philosophical views of science with regard to theories?

HW: Read the excerpt from Introductory Readings in the Philosophy of Science pp. 309-315.

Solubility Lab and Dissolved Gases

Today, we completed the solubility lab.  In addition, we observed the affect of increasing temperature on dissolved gases.

HW: Complete the following journal entry (Due Monday) -

1. What is solubility?
2. How does a solution form (three step - solvation and hydration)? What is the role of energy in this process? Draw diagrams to illustrate both endothermic and exothermic processes.
3. How does an increase in temperature affect the solubility of a solid? and a gas? Explain using the Kinetic Molecular Theory.
4. Solubility Lab - summarize the materials, procedure, observations, analyses (solubility curves plotted on a graph) and conclusions.

Problems using Solubility Curves

After reviewing the "Solubility Curve Practice Problems" handout, it is apparent that solubility increases with increasing temperature, in most cases.  Be sure to develop a hypothesis regarding the varying solubilities of substances with the same cation or anion.

HW: Review your hypotheses for the last two questions in the packet.  Adjust your explanations based upon our discussion in class.

Concentrations of Solutions

Solutions involve the dissolution of a solute in a solvent.  In a saturated solution, the maximum concentration of a solute in a solvent is reached.  No more solute is able to dissolve.  We have been collecting data to determine the saturation of different in water at various temperatures.  Using this data, we will construct solubility curves.

HW: Review the article entitled, "The Dissolution Process" - we will have a reading quiz.  Complete the following problems on a separate sheet of paper (write question and answer).

1. Calculate the molarity (M) of a solution containing 30.0g of NaCl in 1500 mL of solution.
2. Calculate the mole fraction composition of a solution containing 500g C2H5OH and 500g of water.
3. How many grams of KBr could be obtained by evaporating 50 mL of a 0.50M solution of the KBr in water?
4. Calculate the molarity of a solution prepared by adding 500 mL of water to 100 mL of 0.60M solution.
5. To what volume must 100 mL of 6.0M HCl be diluted in order that the resulting solution be 1.0M?

Answers: 1. 0.34M NaCl; 2. C2H5OH = 0.28, H2O = 0.72; 3. 2.98g KBr; 4. 0.10M; 5. 600mL

Bond Energy and Bond Length/Introducing Enthalpy

Today, we compared the bond energies of several diatomic molecules and related the strength of the bonds to the bond length, number of bonds, and position of the elements on the Periodic Table (with regard to the properties of the element based upon its atomic structure).  You were provided with three readings to assist with your understanding of bonding which will be useful in predicting reactions.

HW: Read the three handouts provided in class and make notes and/or write down questions about the readings.

LOL Diagrams

We are using the LOL diagrams to track energy involved in chemical reactions.  Unit 6 Worksheet 4 provided examples of both exothermic and endothermic reactions.  Some of the reactions were spontaneous while others were not spontaneous (energy was deliberately added).  Use this worksheet as a guide for completing LOL diagrams for the five experiments performed in class today.

HW: The following items will be collected in class tomorrow:
1. The homework assigned on Friday.
2. Predicted (balanced) reactions and LOL diagrams for all 10 reactions observed in the labs.

Experimenting with HCl - Part 2

Today, we combined HCl with 5 different solids and recorded observations - dextrose, potassium chloride, sodium hydroxide, sodium carbonate, and silver nitrate.  What were some of the conditions that you observed as evidence that a chemical reaction took place (or did not take place)?

HW: Write out five predicted (balanced) reactions observed in the lab.

Chemical Potential Energy

Prepare for conducting 5 more experiments tomorrow!  We will be discussing our observations and working to predict the products of the reactions (if any).  Chemical potential energy should be a part of this discussion and used to explain the breaking of bonds and the formation of new bonds.

HW: Complete Unit 6 Worksheet 4

Quiz Tomorrow

In preparation for the quiz tomorrow, review Unit 6 Worksheets 1 and 2.  You should be able to fill in coefficients to balance chemical equations, write chemical equations from word equations, write word equations from chemical equations, and identify the type of reaction.

HW: Study for the quiz.

More Practice Balancing Equations

In preparation for the quiz on Tuesday, please complete Unit 6 Worksheet 2 which will give you more practice in balancing equations.  In addition, we are starting to think about why some substances react while others do not.  Given the 5 solids being combined with HCl in the test tubes, did all of the substances react with HCl?

HW: Complete Unit 6 Worksheet 2; Write balanced equations for the 5 reactions from the lab (you will need to predict the products) and write a short paragraph for each reaction in which you explain why the reaction took place based upon some of the trends of the Periodic Table (i.e. atomic radius, ionic radius, ionization energy, electron affinity, electronegativity).

Experimenting with HCl - Part 1

Today, we combined HCl with 5 different solids and recorded observations - sodium bicarbonate, magnesium sulfate, sodium acetate, calcium carbonate, and zinc.  What were some of the conditions that you observed as evidence that a chemical reaction took place (or did not take place)?

HW: Review Unit 6 Worksheet 1

Five Types of Reactions in ACTION!

Demonstrations - (1) Combustion of isopropyl alcohol, (2) Magnesium and hydrochloric acid, (3) Combustion of hydrogen, (4) Decomposition of sodium bicarbonate, and (5) Sodium bicarbonate and hydrochloric acid. 

HW: Write balanced equations for the five reactions that were demonstrated in class, and label each with the type of reaction; complete Unit 6 Worksheet 1 on balancing equations and identify the type of reaction for each.

Don't Get Carried Away - Gases

Before we delve into the next unit, we need to revisit our model so far (M5). 

M5 - What is a substance made of?

• Using the C-12 isotope, we assigned the mass of 1 proton or 1 neutron to be 1 amu.
• This determination allowed the masses of all elements to be assigned in amu.
• Avogadro discovered a relationship between volume and number of gas particles.
• Chemists determined the number of particles required to have the mass of the substance be equal to the atomic mass in grams = Avogadro's number (6.02 x 10^23)
• One mole = 6.02 x 10^23
• Convert moles to grams and grams to moles
• Calculate % composition
• Steps to find the empirical formula of a substance
• Find the mass of each element present in the compound (decomposition or single replacement reaction)
• AB -> A + B    (measure masses of A and B)
• AB + C -> AC + B   (measure mass of AB, mass of B, and subtract mass of AB from the mass of B to get the mass of A)
• Convert mass to moles
• Divide by the smallest number to calculate the ratio of elements
• Assign these whole number ratios as subscripts to write the empirical formula
• Steps to find the molecular formula of a substance
• Obtain a known mass of the substance
• Convert the substance to a gas
• Record the T, P, and V
• Use PV = nRT to calculate the number of moles (n)
• Calculate mass (g)/number of moles = molar mass
• Calculate the (molar mass of the substance)/(molar mass of empirical formula) = multiplier
• Multiply the subscripts of the empirical formula by the multiplier to write the molecular formula
• Determine the limiting reactant - the reactant that "runs out" in a reaction.
• Determine the % yield
         % Yield = (Actual mass of product)/(Theoretical mass of product) x 100

HW: Complete the following journal entry - (1) Define Dalton's Law of Partial Pressures, (2) Complete problems from Unit 8 Worksheet 1 - #5, 8, and 9, and (3) Complete problems from Unit 8 Worksheet 2 - #4, 6, and 8.

Describing Chemical Reactions

Today, we discussed the 5 types of chemical reactions: combination, decomposition, single replacement, double replacement, and combustion reactions. 

HW: Complete the definitions on page 1 and the analysis questions on page 3 of the handout.  The topics that we are currently discussing are found in Chapter 9 in the textbook.

Balancing Chemical Reactions

As a segway into the next unit, test your ability to balance chemical reactions.  Don't worry if you struggle with some of them - you will have many opportunities to practice.

HW: Complete page 2 of the handout by balancing the reactions and drawing particle diagrams.

Chemical Reactions - Ways that Elements Rearrange and Why!

Our next unit will cover the types of chemical reactions which take place between substances.  Last semester, we discussed thermal and phase energy.  We will now begin to explore chemical energy with the understanding the energy can be stored as thermal, phase, and chemical energy and be transferred via heating, radiating, and working.

HW: None.

Review for Exam!

Study! Study! Study!  Students should focus on understanding the concepts and problems in the exam review packet.

Review for Unit 5/8 Exam

Conservation of mass - this unit has been about keeping track of amounts of substances before and after chemical reactions.  The record-keeping allows us to determine empirical and molecular formulas of substances.  In addition, we can compare theoretical yields of products to actual measured masses to calculate a percent yield of product in an experiment.  This is important for industries which track costs of expensive reactants!  They will want to maximize the yield of a product that they sell while minimizing the use of an expensive reactant.

HW: Review for the exam scheduled for Wednesday.

Review Packet Answers:
Unit 5 - Moles, Empirical Formulas, Molecular Formulas
2. a. 101.087g;    b. 96.0862g;     c. 331.728g;     d. 31.998g;     e. 164.09g; f. 331.2g
4. a. 0.215 moles;     b. 0.00361 moles;     c. 13g;     d. 1.29g;     e. 85g
5. a. 9.03 X 10^20 atoms Zn;     b. 80.0g O2;      c. 0.0271 mol Na;     d. 0.306 mol Cl;
    e. 3.28 x 10^-14g Au
6. N2O5
7. MgO
8. C6H6; N2O4
9. C6H12O6
10. 88.8%
11. 77.8% Fe; 22.2% O
Unit 8 -Molar Volume and Gas Stoichiometry
1. 12.4 L; O2 would exert the most pressure.
2. 24 L
3. 3.02 L
6. a. 2NO + O2 => 2NO2

Molar Mass and Molar Volume

Given the established relationship between volume and the number of particles, we should be able to calculate the molar mass of a substance in the gas state.  By measuring the mass of the gas and the volume occupied by the gas, the molar mass can be determined by:

STEP 1 - Convert the volume to number of moles using 22.4L / mole.

                                        Mass of gas
STEP 2 - Calculate        ------------------    which provides g/mol for the substance.
                                        Moles of gas


HW: Unit 5/8 Exam Review; Journal entry - complete the following:

1. What is the difference between an empirical formula and a molecular formula? 
2. Explain how we can find the molar mass of a substance in the gas state. 
3. Give an example problem for finding the molar mass of a substance.
4. Given the empirical formula and the molar mass, how can we determine the molecular formula of a substance (#10 on the first journal entry of the semester)?

Volume and Number of Moles

Avogadro discovered the relationship between the number of atoms/molecules of a gas and the volume occupied by the gas.  Of course, as more gas particles are added, the volume increases.  By experiment, chemists have been able to determine the volume occupied by one mole of any gaseous substance: 22.4 L at STP.  This is called the molar volume. 

HW: Journal entry (assigned 1/19/11); Unit 8 Worksheet 2.

Unit 8 Worksheet 2 Answers:
(assume all zeros to be significant)
1. 11.2 L of O2
2. 3.50 L + 4.00 L + 2.54 L + 5.55 L = 15.59 L
3. 0.00995 mol H2
4. 1.06 L of CO2
5. 422 mL of Cl2
6. 3.46 g P4
7. 4.96 L of NH3
8. 27.1 L of H2

Gases Revisited - Dalton's Law of Partial Pressures

According to Dalton's work with gases, the total pressure of a gaseous mixture is the sum of the partial pressures exerted by the different types of gases.  Each gas would exert a pressure if it were in the container by itself, and it would exert the same pressure if it were in the container along with other gases (assuming no interaction between the particles).  Based upon the relationship between P and n in the Ideal Gas Law, the mole ratio in a mixture of gases determines the partial pressure of each gas.

HW: Read "Dalton's Law of Partial Pressures" article; finish Worksheet 1 (Unit 8), journal entry due Tuesday - complete the analysis and conclusion questions from the "Chemistry in a Bag" handout using the data provided.

Unit 8 Worksheet 1 Answers:
(assume all zeros to be significant)
1. 3.9 atm
2. 95.1 mL
3. 14 mL
4. 750 mmHg
5. %nCO2 = 33.3%; %nN2 = 50.0%; %nHe = 16.7%
6. PN2 = 600 mmHg; PO2 = 150 mmHg; PAr = 8 mmHg
7. 20%
8. 710 mmHg; 97%
9. 713 mmHg; 23 mL

Unit 5 – Determining an Empirical Formula (Journal Entry)
Introduction
In 1778, Lavoisier concluded that combustion was a reaction of oxygen in the air with a sample of matter. He realized that as the substance burned gained mass, the same mass was lost from the surrounding air. A great deal of chemical knowledge has been amassed by using simple combustion experiments conducted with crucibles, burners, and balances.

An empirical formula gives the simplest whole number ratio of the different atoms in a compound. The empirical formula does not necessarily indicate the exact number of atoms in a single molecule. This information is given by the molecular formula, which is always a simple multiple of the empirical formula.

In this experiment, you will determine the empirical formula of a magnesium-oxygen product, a compound that is formed when magnesium metal reacts with oxygen gas. According to the law of conservation of mass, the total mass of the products must equal the total mass of the reactants in a chemical reaction. Therefore,

mass Mg + mass O2 = mass MgxOy

Since you will measure the mass of magnesium and the magnesium-oxygen product, you will be able to calculate the mass of oxygen consumed during the reaction.

Data Table
Items                      Mass (g)
-------------------------------
Magnesium              2.54
Magnesium oxide     3.95

Calculations:
1. Assuming all of the magnesium has reacted, what is the mass of magnesium reacted?
2. What is the mass of the magnesium oxide produced?
3. Determine the mass of oxide in the magnesium oxide.
4. Determine the number of moles of magnesium, then the number of moles of oxide.
5. Determine the ratio:

Conclusion:
1. Since you believe that atoms combine in simple, whole-number ratios, what do you think is the likely ratio?
2. How does your value compare to the accepted value?
3. What is the empirical formula of magnesium oxide?
4. Write a balanced equation for the reaction.
5. What are some possible errors in your experiment?
6. What was the limiting reactant in the experiment?
7. What was the % yield of magnesium oxide?

Empirical Formula, Limiting Reactant, and Percent Yield

Working through the analysis and conclusion questions, we are working to determine the empirical formula, limiting reactant, and percent yield for the experiment involving calcium chloride and sodium bicarbonate. 

HW: Complete the analysis and conclusion questions for the experiment.

Below are some hints to assist with your work:

Analysis:
1. You measured the mass of calcium chloride that I put into your baggie.  Use percent composition to determine the mass of calcium present in your baggie.

40.078g Ca
------------------
110.978g CaCl2

2. This is the mass of material in your filter.
3. This is the mass of material in your filter minus the mass of Ca in your bag.
4. Using answers from #1 and #3, convert grams Ca and CO3 to moles.
5. Divide larger number by the smaller number to get the mole ratio of Ca to CO3.

Conclusion:
4. The reaction is: CaCl2 + 2NaHCO3 -> CaCO3(s) + 2NaCl(aq) + H2O(l) + CO2(g)

6. What is the limiting reactant (all of the substance is used up in the reaction)?

• From your mass in grams, determine the number of moles of calcium chloride that you had in your baggie. 
• Now, looking at the reaction, how many moles of sodium bicarbonate react with one mole of calcium chloride? (2)
• Multiply your number of moles of calcium chloride by 2 to find the number of moles of sodium bicarbonate needed to react with all of your calcium chloride.
• Convert the number of moles of sodium bicarbonate to grams of sodium bicarbonate.
• Compare your calculated mass of sodium bicarbonate to your measured mass (from the lab) of sodium bicarbonate.  If your calculated mass is greater than your measured mass, then you did not have enough sodium bicarbonate to react with your calcium chloride (sodium bicarbonate is your limiting reactant).  If your calculated mass is less than your measured mass, then you had more than enough sodium bicarbonate to react with all of your calcium chloride (calcium chloride is your limiting reactant).

 7. What was the percent yield (actual mass of product/theoretical mass of product) from the experiment?

• Given a measured amount of your limiting reactant (assuming that it all reacted), you should be able to produce a certain amount of product.
• Convert the measured mass of your limiting reactant to moles.
• Now, looking at the reaction, how many moles of calcium carbonate are produced one mole of your limiting reactant?
• Using this ratio, convert the moles of limiting reactant to moles of calcium carbonate.
• Convert the moles of calcium carbonate to grams of calcium carbonate.  This is your theoretical mass of calcium carbonate.
• Calculate the % yield by

       actual measured mass of calcium carbonate

% yield = ------------------------------------------------------------ x 100

theoretical mass of calcium carbonate

How much Calcium Carbonate?

After drying overnight, all groups obtained the mass of the filter paper and solid calcium carbonate in class.  The necessary data has been collected to determine the empirical formula for calcium carbonate.  We will be working through the analysis and conclusion questions to find the empirical formula as well as determine the limiting reactant and the percent yield.

HW: None - R&R weekend! Enjoy!

Chemistry in a Bag!

Calcium chloride reacts with sodium bicarbonate to produce calcium carbonate (a solid) along with some other substances.  Today, you performed the experiment in the baggie.  Did you notice any temperature changes?  Did you notice anything about the volume of the contents of the baggie? How did the mass of the material before the reaction compare to the mass after the reaction?

HW: Unit 5 Worksheet 3 - More Practice on Empirical Formulas (Naming compounds)

Moles to Grams & Grams to Moles

Everyone should be working to master the problems on Worksheet #1 (Unit 5) to convert grams to moles and moles to grams.  In addition, moles can be converted to numbers of atoms/molecules by multiplying by Avogadro's number - 6.02 * 10^23!

HW: Prepare for the quiz - 4 problems to convert grams to moles and moles to grams!

Percent Composition of Oreo Cookies

What is the percent composition of crunchy cookie and creamy filling for regular Oreo cookies and Double Stuf Oreo cookie?  Are Double Stuf Oreo cookies truly double-stuffed?  Based on the data collected in class, we will determine the percent composition of oreo cookies.  Bring your data to class tomorrow.

HW: Complete the journal entry assigned last week; prepare for quiz regarding mass/mole/number of particles conversions.

Determining Empirical Formulas from Composition

Because you can calculate the composition given the formula, you should be able to work the other way around and get formulas from their compositions.  That's what we worked on today. 

Weights

Moles

Mole Ratio

Atom Ratio

Express as Formula

Students should review the example problems that we worked in class.  We will continue this work in class on Monday.

HW: Spend 10 to 15 minutes reviewing problems; Senior Thesis!

Mole Day!

DID YOU KNOW???  Mole Day is celebrated annually on October 23 from 6:02 a.m. to 6:02 p.m. to commemorate Avogadro's Number (6.02 x 10^23), which is a basic measuring unit in chemistry.

Practice, practice, practice.  That's what is required to master molar mass problems.  Keep working through the additional example problems from class.  For extra practice, you can work some problems in Chapter 10 - #1-9 on page 342, #30 - 39 and #45 on page 344.  Next week, we will be having a quiz.

HW: Complete a journal entry using the Unit 5 Notes sheet - include all of the information in your journal (due Tuesday).

Are there MOLES in the classroom?

We found that the mole is a very useful converstion between mass and number of particles.  Since substances combine at the atomic level, we must keep track of the number of particles of each type in reactions.  You should review the handout from class entitled "The Mole" which included example problems.  Make sure that you are able to work these types problems with ease.

HW: Read the handout and review the example problems from "The Mole" handout; Read "The History of 6.02 times 10 to the 23rd" article.

Welcome Back! - What is a Mole?

During the next week, we will be exploring and using a Mole - not the burrowing animal!  Today, we determined that we can the number of "particles" of a substance can be determine by measuring mass.  Now, do we know the mass of individual atoms?  Actually, we do.  But Dalton struggled with this...

In estimating atomic weights, Dalton was confronted with certain grave difficulties.  Since it is impossible to weigh single atoms, any system of atomic weights must be formulated on a comparative basis.  The atom of some element must be arbitrarily selected as the reference weight.  Dalton chose the hydrogen atom and assigned one as its weight.  The atomic weight of oxygen could then be found by either (1) comparing the weights of equal numbers of oxygen and hydrogen atoms or (2) finding by analysis the combining weights of oxygen and hydrogen in water.  Dalton considered the first approach but rejected it.  Since to him, atoms in a gas were analogous to a pile of shot, and since he believed that atoms of different gases varied in diameter, therefore equal volumes of gases could not contain equal numbers of atoms.  The second approach he considered valid.  However, reflection will show that it is valid only when the ratio of atomic combination is known.[1]

   [1] A. Ihde, The Development of Modern Chemistry, Harper & Row, 1964.

The breakthrough came when Avogadro devised an hypothesis to account for the observations of Gay-Lussac regarding the reacting volumes of gases.  Gay-Lussac had noted that gases appeared to react in simple integer ratios.

Avogadro’s paper of 1811, based on Gay-Lussac’s law and Dalton’s atomic theory, reconciled the … problems. Starting with the assumption that equal volumes of all gases contain equal numbers of molecules under similar conditions, Avogadro proceeded to analyze the facts of gaseous combination. Using the examples discussed by Gay-Lussac, he showed that the ambiguities disappeared if he assumed that the molecules involved in typical reactions might split into “half-molecules”; that is, he supposed the existence of molecules of elemental gases which contained more than a single atom. He did not use the term “atom”, but always used the term “half-molecule” as its equivalent.  (Ihde)

Once we accept Avogadro’s Hypothesis, we can compare the mass of various gases and deduce the relative mass of the molecules.  So, from the mass, can we determine the number of particles?  Yes.  In fact, the word chosen to represent the standard weighable amount of stuff, the mole, comes from the Latin: lump of stuff.

HW: Read the article "What is a Mole in Chemistry?"

Metals in Water - What's the Reaction?

We repeated our experiments from Friday.  What were your observations of the demonstration?  Based on our discussion, we determined that the metal atom gives up an electron to form a cation.  What observations support this idea? 

HW: Read the handout about hydrolysis and complete a journal entry for the demonstration - summarize the materials, procedure, observations (including the indicator), and conclusions.  Write the chemical reaction for the metal in water.  Write the reaction for the addition of acid to the solution.  Describe the products in terms of acids and bases.  Do you think that this reaction would take place with alkaline earth metals?

Metals and Water - what's in a reaction?

Today, we observed substances from the alkali metals group placed in water - lithium, sodium, and potassium.  What did you observe during the demonstration?  How were the interactions similar?  How were they different?  What do you think was happening?

HW: Have a restful weekend...finals are approaching.

We can name them, but how do the form?

Today, we reviewed the rules for naming ions, compounds, molecules, and acids.  Now that we can name these substances, we must begin to address the question: How do the form?  We will spend our remaining (few) days of the fall semester attempting to answer this question. 

HW:  Journal entry - Summarize the naming of ionic compounds and molecules.

Rutherford and the Gold Foil Experiment/Naming Compounds

Rutherford, a student of Thomson, continued designing experiments to study the interior structure of the atom.  Today, you received a handout discussing his gold foil experiment.  What findings did he make which altered the Plum Pudding Model?  What was his new model for the atom?

Naming compounds: Remeber the general rule - ionic bonds involve a metal (monoatomic cation) and a non-metal (monoatomic or polyatomic anions); covalent bonds are between non-metals.  Ammonium ion (NH4+) is the exception to the rule.

HW: Journal entry: Summarize the Gold Foil Experiment and Rutherford’s new model of the atom.

JJ Thomson and the Plum Pudding Model

We have discussed the work of JJ Thomson in studying cathode ray tubes.  Based upon his findings, Thomson put forth a model of the structure of an atom which included charged particles.

HW: Journal entries: (1) Summarize the Sticky Tape Lab; (2) Explain the discovery of the electron (cathode rays) and JJ Thomson’s model of the atom – describe the behavior of electrons.