UNIT 3
QUANTITIES IN CHEMICAL REACTIONS
Goals
ACCURACY AND PRECISION
Accuracy – a measure of how closely a result (or a set of results approaches the true or accepted value.
Precision – a measure of how closely the results within a set of results agree with each other.
Good accuracy good accuracy poor accuracy
Good precision poor precision good precision
Significant Digits (figures)
the measuring tool used when making a measurement will determine the precision of a measurement
in science we use significant digits to show the precision of a measurement Example: 4 kg – mass between 3.5 kg and 4.5 kg
– 1 significant digit
4.000 kg – mass between 3.9995 kg and 4.0005 kg – 4 significant digits
Rules to determine the number of significant digits
1. All non zero digits are significant
2. Zeros:
b) Zeros between significant digits are significant Example: 0.1047 4 significant digits
Significant
c) Zeros after the number = uncertain Example: 2000 1, 2, 3, or 4 SD
Write in scientific notation to indicate number of significant digits.
2 x 103 1 SD
2.0 x 103 2 SD 2.000 x 103 4 SD
d) Zeros after the number after the decimal are significant Example: 1.400 4 SD
3. Adding/Subtracting
The value with the fewest number of decimal places (ie: SD after the decimal place), determines the number you round your answer to.
Example:
4. Multiplication/Division
The answer must have the same number of significant digit as the measurement with the fewest significant digit.
HOMEWORK: Worksheets
6.1and 6.2 Qualitative and Quantitative Analysis
Goals
Complete the note outline by referring to pages 260-263 in the textbook
Qualitative Analysis:
Example
Qualitative Analysis:
Example
Quantitative Analysis and the Environment:
Indoor Carbon Dioxide Levels
Why are CO2 levels monitored in the workplace? Pollutants and Health
What occurred in 1976 that caused scientists to suspect a link between the environmental conditions and the incidence of diabetes? Describe a study that was conducted to test the suspected link.
Quantitative Analysis in Sport
What role does quantitative analysis play in sports?
Doping in Cycling
1) What is EPO?
3) How else could athletes increase their HCT and what is a normal level for an elite cyclist? What is the side effect to increasing the HCT beyond the normal range?
4) How do athletes avoid detection?
6.3The Mole and Molar Mass Goals
Strategy:
1. Write the word “mole” on the chalkboard. Ask the students what it means. Compare the chemistry mole to a dozen. Explain that it represents a set group of things, 6.02 x 1023, just as a dozen is 12 things. The number 6.02 x1023 is called Avogadro’s Number. It is the number of 12C atoms needed to have a mass of 12.0 g.
2. Ask the students, how many students are in one mole? How many pencils in one mole? How many atoms in one mole? Remind the class that a mole is a fixed number of things just as a dozen is.
3. Give each pair of students a bag of sixty paper clips. One student is to count the paper clips as quickly as possible, while the other records the time in seconds. Ask the students how long it would take them to count out a mole of paper clips, helping them (if need be) set up the proportions:
and then solve for x. Have them convert this answer into minutes, hours, days, and finally years. Is it reasonable to count out a mole of paper clips? Of anything? No!
4. Hold up a test tube containing one mole of aluminum wire and ask the class how many atoms of Al are in the tube. Hold up a test tube with one mole of H2O and ask how many molecules of H2O are in the tube. Repeat with the other test tubes. Hold up the test tube containing a mole of aluminum atoms and tell the class that you only needed five minutes to put the mole of atoms into the tube. Repeat with the other tubes. How could this be done? Remind the class that the average atomic mass of an atom of each element is listed in the Periodic Table
Avogadro’s constant (NA)
*the number of 12C atoms needed to have a mass of 12 g *6.022x1023 units or particles
Examples: 1 dozen eggs = 1 gross of pencils = 1 mole of atoms =
Example: How many atoms are in 2 moles of potassium?
# of K atoms = (2 moles)(6.022 x 1023 atoms/mole) =1.204 x 1024 atoms of K
On your calculator: 2 x 6 . 0 2 2 exp 2 3 =
Molecular Mass, Formula Mass and Molar Mass
*Molecule: a cluster of bonded atoms. The smallest unit of a covalent compound that retains all of its chemical properties.
*Molar mass: *Molecular mass:
*Formula mass: mass of the atoms in the smallest grouping of an ionic compounds expressed in amu. The molar mass (in g/mol) of a substance is numerically equal to the formula mass.
Element 1 atom 1 mole of atoms
Mg 24.32 24.32 g
Neon 20.2 20.2 g
Example 1:
What is the molar mass of water?
Example 2:
What is the molar mass of NaCl?
Find the molar mass of each of the compounds shown below:
1. H3PO4
2. CuSO45H2O
3. Sn(OH)4
4. (NH4)3PO4
5. KAl(SO4)212H2O
6. The _______________ is the mass in grams of Avogadro’s number of molecules.
7. Determine the mass of one mole of KCl.
Therefore the mass of one mole of KCl is 74.6g.
NUMBER OF MOLECULES IN A MOUTHFUL OF WATER
Purpose: To determine the number of molecules in a mouthful of water. You will not actually drink the water.
Materials Water
250mL beaker
Scale Balance
Procedure: Fill a beaker with water and determine the mass. Pour approximately
one mouthful of water down the drain. Re-mass the beaker and water.
Include all the units involved. Take your final answers to 3 significant figures.
Observations
1. Mass of beaker and water initially __________ g
2. Mass of beaker and water after pouring a mouthful down the drain __________ g
3. Mass of water poured ________________________ g
4. Molar mass of water
5. Moles of water poured
6. Number of molecules/mole
7. Number of molecules swallowed
Conclusion
1. Describe in point form, on the back of the handout, how you determined the # of molecules in a mouthful of water.
6.5 Calculations Involving the Mole Concept
MOLES of Substance
GRAMS of substance NUMBER of particles
where n = # of moles, m = mass (g) and M = molar mass (g/mol)
(1) How many moles are in 32.0 g of potassium?
The number of moles of potassium is 0.818 mol.
(2) How many molecules are there in 40.0 g of water? 1) MH2O = 2(1.01) + 16.00
= 18.02 g/mol
2)
3)
The number of molecules of water is
g/mol
x g/mol
x 6.022x1023 particles/mol
(3) How many molecules are present in 7.50 mol of H2SO4?
(4) How many atoms are present in 0.45 mol of H2SO4
(5) How many moles are present in 139 g of HC2H3O2?
(6) What is the mass of 0.763 mol of HNO3
MOLE PRACTICE QUESTIONS
1. How many moles are present in 205.6 g of HC2H3O2?
2. What is the mass of 0.625 mol of HNO3?
3. How many atoms are present in 63 g of H2SO4? In 29 g of Br2?
4. What is the mass of 4.803 x 1022 molecules of C
3H7NO2?
5. If a sample consists of 1.41 x 1023 molecules, how many moles are present?
6. How many atoms are present in 0.27 mol of C2H4?
7. How many sodium atoms are present in a 71.2 g sample of sodium?
8. How many atoms are present in 33.0 g of Ca(HCO3)2? In 256 g of Sr(NO3)2?
9. What is the mass of 7.45 x 1022 molecules of C
6H12O6? Of 3.11 x 1024molecules of
C4H9SH?
10. a) Calculate the number of moles of aluminum oxide, Al2O3 in 3.01 x 1022 molecules
of the compound.
b) Calculate the number of atoms of aluminum present in the same sample of Al2O3.
6.6 Percentage Composition
Goals
Percent composition
*USE THE LAW OF DEFINITE PROPORTIONS TO INTRODUCE THE SUBJECT
Any sample of a pure compound always consists of the same elements combined in the same proportion. It is possible to express the molecular composition of a compound in various ways.
1. Every molecule of water has the formula H2O. One molecule contains one atom of O and two atoms of H.
2. The molar mass of H2O, 18.02 g/mol always contains 16.00 g of O and 2 g of H.
3. Percent composition – the percentage of the mass of a compound represented by each of it constituent elements. This means the percentage of the total mass of the compound contributed by each element.
Example:
2. Calculate the percentage by mass of each element in Ba(IO4)2.
6.7 Empirical vs Molecular Formulas Goals
Molecular Formula is "a formula indicating the actual number of atoms of each element making up a molecule." In other words, the molecular formula must accurately state the exact number of atoms of all of the elements in one molecule of the substance.
Other Important Information:
The formulas for Ionic Compounds are empirical formulas.
A molecular formula is "a whole number multiple of an empirical formula."
Classify the following formula as “molecular” or “empirical”.
BH3 B4H10 B5H7 B5H11 B6H12
Convert Molecular Formula to Empirical Formula Glucose
Step 1: Molecular Formula =
Step 2: All subscripts are divisible by Step 3: Empirical Formula =
Acetic Acid
Step 1: Molecular Formula =
-- Step 1b: Standard Molecular Formula = Step 2: All subscripts are divisible by Step 3: Empirical Formula =
6.7 and 6.9 Calculating Chemical Formulas Goals
1. When you are given percentage composition of elements in a compound, you can find the empirical formula of that compound. When solving this type of problems, you can use the following rhyme to help you through the steps of the procedure to solve this type of problems.
Percent to mass Mass to mole Divide by small Multiply ‘til whole
2. Here’s an example of how it works. A compound consists of 72.2% Mg and 27.8% N by mass. What is the empirical formula?
ELEMENT MASS Number of
Moles
Ratio of
Moles
Whol-Number Ratio Mg
N
And the formula of the compound is
Calculation Molecular Formula
How do we determine the molecular formula of a compound? We need both the empirical formula and the molecular mass of the compound.
This is an example of how to find the molecular formula.
The empirical formula of a compound of carbon and hydrogen was found to be CH. Its molar mass was also found by experiment to be 78 g/mol. What is the molecular formula of the compound?
Any molecular formula having a 1:1 carbon-to-hydrogen ratio is possible, so the molecular formula could be the same as the empirical formula, twice the empirical formula, three times the empirical formula and so on. You can find the number of empirical formula units in one molecular unit by dividing the molar mass by the mass of the empirical formula unit.
Example: Molecular Formula
A reactive compound contains 13.2% B and 86.8% Cl by weight. If the molar mass of the compound is known from a separate experiment to be 163 g/mol, what is the molecular formula?
Element Mass (g) # of moles Ratio of moles
Whole-Number Ratio B
Cl
Therefore, the empirical formula is
Molecular formula is found by dividing the molar mass of the compound by the molar mass of the empirical formula.
Thus, the molecular formula is two times the empirical formula – two times ,
COMPLETE: Lab – Analysis of a Hydrate (Handout) Hydrates to use:
Lab: Analysis of a Hydrate
Purpose: To determine the number of moles of water molecules in one mole of hydrate.
Introduction: Salts which contain water as part of their crystal structure are called
hydrates and the water is called water of hydration. When a hydrate is heated, water is given off:
MaXbxH2O(s) MaXb(s) + xH2O(g)
Materials: 100 mL beaker glass stirring rod
approximately 3g of hydrate hot plate
wire mesh
Procedures:
1. Determine the mass of the 100 mL beaker
2. Place approximately 3g of the unknown hydrate in the crucible. Determine the
mass of the beaker and hydrate to 3 significant digits.
3. Place the beaker on the hot plate and turn on the hot plate to the maximum
temperature. Heat gently at 5 on the hot plate for about 5 minutes. (Some of the hydrates may spatter). Increase the heat and heat strongly at 10 for about 5
minutes.
Cool
for 5 minutes and then determine the mass of the beaker andcontents.
Observations:
Mass of 100 mL beaker =
Mass of beaker and hydrate =
Mass of hydrate =
Mass of beaker and dehydrated solid =
Mass of dehydrated solid =
Mass of water lost =
1. Determine the number of moles of dehydrated solid MaXb (using the molar mass
given by the teacher).
2. Determine the number of moles of water lost.
3. Determine the ratio of number of moles of dehydrated solid: moles of water (MaXbxH2O) (simplify the ratio of dehydrate: water and express it as 1:x).
Conclusion
Determining the Chemical Formula of a Hydrate
Question
What is the molecular formula of the hydrate of ________________________.
Materials 100 mL beaker hot plate
electronic balance glass rod
Approximately 3 g of hydrate
Procedure
1. Measure the mass of an empty beaker. Record the mass in the table. 2. Add 3 g of hydrate to the beaker.
3. Measure the mass of the beaker and hydrate. Record the mass in your table. 4. Set the hot plate to 5 and heat the container with the hydrate for approximately 5
minutes . Heat for another 5 minutes at 10. Allow to cool for approximately 5 minutes.
5. Find the mass of the container and the anhydrous substance. Record the mass in your table.
6. Place the anhydrous substance in the fume hood. Do not put in the sink or in the garbage.
Observations Mass of beaker (g)
Mass of beaker and hydrate (g)
Mass of beaker and anhydrous substance (g) Mass of hydrate (g)
Mass of anhydrous substance (g) Mass of water(g)
Analysis
11. Determine the percent by mass of water in your sample of hydrate. Show your calculations clearly. 2 marks
13. Calculate the percentage composition of the hydrate. 4 marks
14. Suppose you did not completely convert the hydrate to the anhydrous compound. Explain how this would affect: (2 marks)
a) The calculated percent by mass of water in the compound.
b) The molecular formula that you determined.
15. Suggest a source of error that would result in a value of x that is (2 marks) a) higher than the actual value.
7.1 Mole Ratios in Chemical Equations
Goals
*Balancing an equation: multiplying the reactants and products in a reaction equation by coefficients to ensure that equal number of each atom is on both sides of the equation.
*Steps in balancing an equation
(1) Balance all atoms except O and H first. (2) Then balance O
(3) Finally, balance H
(4) Make sure that the coefficients are in their lowest terms
Remember, only change the coefficients when balancing. Never change the subscripts within the chemical formula!
Examples
1. KOH + SO2 K2SO3 + H2O
Counting atoms:
Left side of equation Right side of equation
1K 2K
3O 4O
1S 1S
1H 2H
The reaction is not balanced.
put coefficients into the equation that will make the number and type of atoms on the left exactly the same as on the right.
2KOH + SO2 K2SO3 + H2O
2. MnO2 + HCl MnCl2 + H2O + Cl2
How to interpret a balanced chemical equation
N2(g) + 3H2(g) 2NH3(g) 1 molecule 3 molecule 2 molecule 1 mole 3 mole 2 mole
1( ) 3( ) 2( ) # of molecules ( ) ( ) ( )
*mole ratio: ratio of the amount in moles of reactants and/or products in a chemical reaction ( )
7.2 Calculating Masses of Reactants and Products
Goals
*quantitative calculations based on chemical equation = stoichiometry
Example
How many moles of oxygen are required to produce 242.0 g of magnesium oxide? What mass of O2 is required?
i) Write the balanced chemical equation for the reaction.
ii) Determine the number of moles of the substance whose volume or mass is known.
iii) Using the balanced equation, determine the number of moles of the substance you’re looking for.
Example
How many grams of oxygen will be produced from the decomposition of 346 g of potassium chlorate?
i) Write the balanced chemical equation for the reaction. 2KClO3 2KCl + 3O2
ii) Determine the number of moles of the substance whose volume or mass is known.
iii) Using the balanced equation, determine the number of moles of the substance you’re looking for.
iv) Convert the number of moles to the units requested in the problem.
Therefore, of oxygen will be produced.
7.3and 7.4 Calculating Limiting and Excess Reagents
*stoichiometric quantities: molar quantities of reagents which, when reacted together, are completely consumed.
Example
Pb + S PbS 1 mol 1mol 1 mol 207.2 g 32.1 g 239.3g
BUT
Pb + S PbS 1 mol 2mol 1 mol 207.2 g 64.2 g 239.3g
limiting excess reagent
*limiting reagent: -all of it is consumes in a reaction.
-determine the maximum amount of product that can be formed.
*reagent in excess: - some is left over at the end of the reaction.
Example
What mass of HCl is produced when 4.50 g of H2(g) and 140.0 g of Cl2(g) are reacted according to the following equation?
H2(g) + Cl2(g) 2HCl(g) 4.50 g 140.0g i) Determine the number of moles of each reactant.
ii) Calculate the number of moles of HCl expected from each reagent.
From H2
From Cl2
iii) Which is the limiting reagent?
The one producing the least number of moles.
Therefore, Cl2 is the limiting reagent and H2 is in excess.
Therefore, 144 g of HCl is produced.
Example
What mass of S is produced when 15.6 g of KMnO4, 7.95 g of H2S, and 15.3 g of H2SO4 are reacted according to the following equation?
2KMnO4 + 5H2S + 3H2SO4 K2SO4 + 2MnSO4 + 8H2O + 5S Mass
(g)
15.6 7.95 15.3 0 0 0 0
Moles (n/M)
0.0987 0.233 0.156 -- -- -- 0.233
Therefore,
HOMEWORK: p.330, 1-7 day 1
p.335, 1-11 day 2
Complete Investigation: Limiting and Excess Reactants
Limiting and Excess Reactants
In this investigation, you will predict and observe a limiting reactant. You will use the single displacement reaction of aluminum with aqueous copper(II) chloride.
2Al(s) + 3CuCl2(aq) 3Cu(s) + 2AlCl3(aq)
Note that copper(II) chloride, CuCl2, is light blue in aqueous solution. This is due to the
Cu2+(aq) ion. Aluminum chloride is colourless in aqueous solution.
Question
How can observations tell you which is the limiting reactant in the reaction of aluminum with aqueous copper(II) chloride?
Materials 250 mL beaker stirring rod 0.51 g CuCl2
0.25 g Al foil
Safety precautions
The reaction mixture may get hot. Do not hold the beaker as the reaction proceed.
Procedure
1. To begin the reaction, add about 150 mL of water to the beaker that contains the aluminum foil and copper(II) chloride.
2. Record the colour of solution and any metal that is present at the beginning of the reaction.
3. Record any colour changes as the reaction proceeds. Stir occasionally with the stirring rod.
4. When the reaction is complete, return the beaker, with its contents, to your teacher for proper disposal. Do not pour anything down the drain.
Observations
Beginning of the reaction:
During the reaction:
Analysis
1. According to your observations, which reactant was present in excess? Which reactant was the limiting reactant?
2. Do stoichiometric calculations to support your observations of the limiting reactant.
3. Magnesium and hydrogen chloride react according to the following skeleton equation:
Mg(s) + HCl(aq) MgCl2(aq) + H2(g)
a) Balance the skeleton equation.
b) Examine the equation carefully. What evidence would you have that a reaction was taking place between the hydrochloric acid and the magnesium?
Which Reagent is Limiting and How Much Precipitate is Formed?
The purpose of this Investigation is to test gravimetric stoichiometry by predicting and determining the mass of precipitate produced by the reaction of aqueous strontium nitrate and aqueous copper(II) sulfate pentahydrate.
Question
What is the mass of precipitate produced by the reaction of 2.00 g of strontium chloride with 2.00 g of copper(II) sulfate pentahydrate?
Materials 3 beakers
2.00 g strontium chloride
2.00 g copper(II) sulfate pentahydrate eyedropper hot plate funnel filter paper Erlenmeyer flask Safety Precautions
Careful with hot liquids. Never leave hot plate unattended.
Procedure
1. Measure 100 mL of water in a beaker. Heat the water using a hot plate. 2. In a beaker, measure 2.00 g of strontium chloride using an electronic balance. 3. In a second beaker, measure 2.00 g of copper(II) sulfate hydrate.
4. Once the water is warm, transfer water into the beaker containing strontium chloride using an eyedropper. Transfer enough water to totally dissolve the solid. Do not add more water than necessary.
5. Repeat step 4 for the beaker containing copper(II) sulfate hydrate. 6. Turn off the hot plate.
7. Mix the two solutions together by transferring the copper(II) sulfate hydrate solution to the strontium chloride solution.
8. Wash the walls of the beaker that contained the copper(II) sulfate hydrate solution using warm water and the eyedropper. Transfer to the beaker containing the two solutions.
9. Put the reaction beaker in an ice bath for 5 minutes to maximize the recovery of crystals.
10. Measure out the mass of your filter paper. Record the mass.
11. Filter your solution, washing the beaker with water to recover most of the crystals. 12. Let the filter paper dry overnight.
Observation
Mass of Strontium chloride (g) Mass of copper(II) sulfate hydrate (g) Mass of filter paper (g)
Mass of filter paper + precipitate (g)
Analysis
1. Write the chemical equation for the reaction studied in this experiment.
2. Do stoichiometric calculations to determine which reactant is limiting. What mass of the other reactant will remain after the reaction is complete?
3. Based on the limiting reactant, predict the mass of precipitate expected.
7.5 The yield of a chemical reaction
*Most reactions do not produce the exact amount of product that is predicted by the balanced chemical equation
1) Theoretical Yield :
2) Actual Yield :
3) Percentage Yield :
Example
Bromine was made according to the following reaction: HBrO3 + 5HBr 3Br2 + 3H2O
If 10.0 g of HBrO3 was reacted with an excess of HBr and 26.3 g of Br2 was produced, what was the percentage yield of the reaction?
i) Determine number of moles of reactant used.
ii) Use the molar ratio to compare the limiting reagent to the product.
mass of Br2 expected iv) Determine % Yield
QUANTITATIVE CAREERS
*stoichiometric calculations are used on an everyday basis by chemists *used as a “tool” to provide information about the experiment conducted.
Complete Part I and Part II and submit.
Part I (10 marks)
Using your device, select and describe 3 of the following careers in science.
Choose 1 of the careers and describe the education and training necessary for the career.
Examples
1. Art Conservationist
2. Chemical Laboratory Technician 3. Forensic Chemist
4. Pharmacy Technician 5. Veterinarian Technician 6. Respiratory Therapist
Part II (12 marks)
Contributions of notable Canadian Scientists
Describe a scientific contribution made by each of the following scientists. Jed Harrison
Louis Slotin Paul Kebarle
UNIT 2 – Performance Task
Quantitative Analysis of A Reaction
The active ingredient in baking soda is sodium hydrogen carbonate, NaHCO3(s). Upon
heating, this ingredient releases bubbles of a gas that give baked goods a light, spongy consistency. Your task is to use your knowledge and skills of quantitative analysis to identify the decomposition reaction of sodium hydrogen carbonate from five possibilities. In each of the possible reactions listed below, one or more gases are produced and a solid remains:
1. sodium hydrogen carbonate sodium (solid), water vapour, carbon monoxide
gas, oxygen gas
2. sodium hydrogen carbonate sodium carbonate (solid), water vapour, carbon
dioxide gas
3. sodium hydrogen carbonate sodium hydroxide (solid), carbon dioxide gas
4. sodium hydrogen carbonate sodium oxalate, Na2C2O4(s), water vapour, oxygen
gas
5. sodium hydrogen carbonate sodium oxide (solid), water vapour, carbon
dioxide gas
You will quantitatively analyze the decomposition of sodium hydrogen carbonate and then identify the chemical reaction that has taken place form the five reactions given above. In order to do so, you will design and perform an experiment to determine the mass of solid product formed when a known mass of sodium hydrogen carbonate is completely decomposed by heating.
Your report should include completed Introduction, Purpose, Materials, Procedure,
Data and Observations, Analysis, Discussion, Conclusion, Questions
Introduction: - one paragraph which discuss the background information needed to do the lab.
Purpose: - a statement of purpose
- “The purpose of this lab is to …”
Materials: - NaHCO3 3.0 g
- Other equipment you may need
Data and Observations: In the form of a table.
Analysis: Provide ONE sample of EACH different calculation performed in order to arrive at the final results.
The remainder of the results may be tabulated.
Discussion: Discussion of - Analyzed results
- error – human and instrumental - unexpected results
- changes which will improve procedure
- explain differences in theoretical and experimental results
Conclusion Statement that relates back to purpose.
Questions
1. From your knowledge of the properties of the two oxides of carbon, which of the
five possible reactions is not likely a chemical reaction designed for use in the home? Explain.
2. Describe other evidence or diagnostic tests that may be performed on the products
of the reaction to rule out one or more of the reactions under consideration.
3. It is often recommended that we keep a box of baking soda on hand in the kitchen
for extinguishing small grease fires. Relate this application to the properties of the products formed in the reaction you identified.
4. Sodium hydrogen carbonate is a weak base that reacts with acids to produce
