Chemical Quantities The Mole.

Chemical Quantities The Mole

But first… CuCl2 Na3PO4 S2O5 PbSO4∙2H2O HNO2 Silver acetate
Gold (III) iodide Dinitrogen tetrabromide Cuprous nitrate tetrahydrate Hydrofluoric acid

Answers Copper (II) chloride or cupric chloride Sodium phosphate
Pick out the mistake and write it down in your notes…do not say it out loud, or in a whisper, or in sign language, or in any way that would let someone else know the answer before they could figure it out on their own. Copper (II) chloride or cupric chloride Sodium phosphate Dinitrogen pentoxide Lead (II) sulfate dihydrate or plumbous sulfate dihydrate Nitrous acid Ag2C2H3O2 AuI3 N2Br4 Cu(NO2)2 ∙ 4H2O HF

And second… The portfolio… Sections
Notes—date each day, chronological order and not kept in a spiral notebook. HW—dated and put in chronological order Quests—see HW Tests—see HW General—everything that doesn’t have a home in the first set of sections goes here. Daily Journal…

The Daily Journal A statement of what was done that day. (We took notes over moles and Avogadro’s number). ****A summary of your notes. Explain in detail in your own words what was covered in the notes for that day. Or explain what the purpose of the lab was, etc.**** A couple statements about what you do not clearly understand or what you could not clearly describe in part #2 above. What you are doing that night or what you have already done that night concerning this class.

Now we can begin with notes…
Ready?? Now we can begin with notes…

Goals for Learning To convert between moles and number of particles
To convert between mass and moles To convert between moles and gas volume at STP To use molarity in conversions involving solutions To calculate the percent composition of a compound To find an empirical formula from percent composition To use an empirical formula to find a molecular formula

Percent composition, empirical and molecular formulas
Organization of Unit Volume of a solution Number of Moles Number of particles Mass Percent composition, empirical and molecular formulas Volume of a gas at STP

Objectives Explain what a mole is
Convert between moles and number of particles Define molar mass Find the molar mass of an element using the periodic table Calculate the molar mass of a compound Use molar mass to convert between mass and moles

Key Terms Mole (mol) Avogadro’s number Atomic mass Molar mass

Measuring Matter All matter is made of different kinds of particles.
These particles can be molecules, atoms or ions. Elements such as helium and iron exist as single atoms. Other elements, like H, O and N are diatomic molecules. Molecular compounds, like CO2 and water, also exist as molecules.

Measuring Matter Ionic compounds, like ammonium carbonate, consist of cations and anions in formula units. Depending on the substance, different names for the particles are used. In this unit, all atoms, molecules and formula units are referred to as particles.

The Mole The particles in matter are very, very small—too small to see, even with a microscope. Counting the number of particles in a sample is not possible. Instead of counting them, chemists measure the number of particles with a unit called the mole. The abbreviation for mole is mol (I suppose there is a logical reason for this). The mole is a unit for measuring the amount of substance.

The Mole The word mole means a number, similar to the word dozen.
1 dozen eggs = 12 eggs 1 mole eggs = x 1023 eggs One mole of any substance contains x 1023 particles. These particles can be atoms, molecules or formula units.

The Mole The number 6.022 x 1023 is called Avogadro’s number.
It is a very, very large number. A mole of eggs is more eggs than have ever been eaten in the history of the world. If you stacked a x 1023 sheets of paper, the pile would reach to the sun and back more than a million times! However a mole of molecules, atoms and formula units is not very big.

Today… Get ready for practice…
Notes to fill in the flow chart from Friday Journals—in class—5 minutes before the bell—you must write the entire time…

Time to Practice NiCl2 K3PO4 S2Cl2 H2CO3 Cu(NO3)2 Sodium hypochlorite
Silver bromide Hydrochloric acid Strontium permanganate Sulfur trioxide

Time to Practice Nickel (II) chloride or nickelous chloride or #23 chloride Potassium phosphate Disulfur dichlorine Hydrocarbonic acid Copper (II) nitrate or Cupric nitrate NaClO AgBr HCl Sr(MnO4)2 SO3

Time to Recall What is the date of Mole Day?
What time does it begin? End? Write down the food you will be bringing in on Mole Day. If you had a mole of bananas, how many bananas would you have? What is this # called? (Name of dude?)

Mole-Particle Conversions
Volume of a solution Number of Moles Number of particles Mass Use Avogadro’s # of 1 mole = X 1023 particles Percent composition, empirical and molecular formulas Volume of a gas at STP

Chemists use Avogadro’s number to calculate the number of particles in a sample of matter. Suppose you are told that balloon contains 2.00 mol of the gas argon–and nothing else. You want to know how many particles are in the balloon. You want to change the measurement units from moles to particles. Moles Ar  particles Ar

Your conversion factor is 1 mol Ar = x 1023 particles Ar. Write the given number and unit. Include the chemical symbol or formula of the substance. Set up a factor label grid. Use the conversion factor in the grid—making sure to align units so that they will cancel. Give the answer with the correct unit and the correct number of significant digits.

You try… How many moles are in 4.35 x 1024 molecules of CO2? How many formula units are in 3.15 mol of sodium oxide?

Molar Mass We know that 1 mole of a substance is a certain number of particles—regardless of what the particles are. 1 mole of donuts is x 1023 donuts 1 mole of desks is x 1023 desks 1 mole of ¥ is x 1023 ¥

Molar Mass 1 mole of a substance also has a certain mass—but it is different for every substance. 1 mole of donuts does not have the same mass as 1 mole of desks, but they still contain the same number of each.

Molar Mass It is impossible to count the particles in a sample of matter but it is easy to find the sample’s mass. If you know the mass of a substance, you can determine the number of moles in it. To make a mass-to-mole conversion, you need to know about another piece of information from the periodic table. Atomic number—number of protons and electrons in an atom Atomic mass or molar mass is the mass of one atom of the substance or the mass of one mole of the substance.

Molar Mass If you want to know the atomic mass (mass of one atom) the number has units of amu (atomic mass units). If you want to know the molar mass (mass of one mole of atoms) the number has units of g/mol (grams of substance per mole of substance).

The Molar Mass of Elements
The molar mass of carbon is 12 g/mol. The molar mass of oxygen is 16 g/mol. This means that x 1023 atoms of carbon (1 mol) has a mass of 12g. This also means that x 1023 atoms of oxygen (1 mol) has a mass of 16g.

Review Mini What is a dozen? What is a mole?
How many atoms of H are in a dozen of H? How many atoms of H are in a mole of H? What is the mass of one H atom? What is the mass of 1 mole of H atoms? What is the mass of 1 O atom? What is the mass of 1 mole of O atoms? What is the mass of 1 mole of H2O?

Today… Short review time… Get boards
Notes—molar mass, mass to mole conversions, gas volume to mole conversions, solution volume to mole conversions, percent composition and empirical formula concept… HW—review 1 and 2 due tomorrow

Review time… CaI2 NaNO3 FeCl3 SO2 CoF2∙3H2O Lithium nitride
Aluminum hydroxide Carbon tetrabromide Cupric nitrite Magnesium acetate dihydrate

Review time… What number represents a mole?
How many moles of He is 6.35 x 1026 particles of He? When you see that magnesium has a mass number of 24.31; what are the units we will be dealing with?

Answers Calcium iodide Sodium nitrate
Iron (III) chloride or ferric chloride Sulfur dioxide Cobalt (II) fluoride trihydrate Li3N Al(OH)3 CBr4 Cu(NO2)2 Mg(C2H3O2)2∙2H2O

Answers 6.022 x 1023 1050 moles grams / mole

The Molar Mass of Compounds
To find the molar mass of a compound of two or more elements, add the masses of 1 mol of each atom in the compound's formula or formula unit. For example, to find the molar mass of CO2, count the number of C and O atoms in the formula.

Then locate the molar mass of carbon and oxygen on the periodic table. Add the molar masses of each carbon an oxygen that make up the formula. CO2  1 C and 2 O  C = 12.01g/mol, O = 16.0g/mol, O = 16.0g/mol = g/mol CO2  1C(12.01g/mol) + 2O(16.0g/mol) = g/mol

You try… Calculate the molar mass of the ionic compound, Ba(C2H3O2)2. Calculate the molar mass of ammonium sulfate.

Mole-Mass Conversions
Volume of a solution Use molar mass in g/mol Number of Moles Number of particles Mass Percent composition, empirical and molecular formulas Volume of a gas at STP

Molar mass can be used to convert between the mass of a sample and the number of moles in that sample. If you know a sample’s mass, you can find the number of moles in the sample. If you know how many moles are in a sample, you can find its mass. Write the given number and unit as well as the formula for the substance.

Set up a factor label grid. Set up the conversion factor of grams per mole or moles per gram depending on your initial given information. Use the atomic molar mass if converting atoms or use the molecular molar mass if converting compounds. Write the answer with the correct units and the correct number of significant digits.

Example: What is the molar mass of carbon dioxide? How many molecules are in 1 dozen carbon monoxide? How many molecules are in 1 mole of carbon dioxide? How many moles are in 26.52g of CO2? What is the mass in grams of 3.25 mol of NaBr?

Mini Quiz How many grams are in 1.00 mole of calcium phosphate?
How many grams are in 3.20 moles of calcium phosphate? How many molecules are in 4.12 moles of calcium phosphate? How many molecules are in g of calcium phosphate? How many grams in 3.24 x 1025 molecules of calcium phosphate? What is the molarity of a solution that has 153.6g of Ca(NO3)2 in 1.2L of total solution? How many liters does 402g of O2 gas take up at STP?

Objectives Define STP and standard molar volume
Convert between gas volume, moles, mass and number of particles Calculate the molarity of a solution Calculate the mass of solute Calculate the volume of a solution

Key Terms Atmosphere Standard temperature and pressure (STP)
Standard molar volume Concentrated Dilute Concentration Molarity

Use 1 mole of a gas at STP takes up 22.4L of space
Molar Volume Volume of a solution Number of Moles Number of particles Mass Percent composition, empirical and molecular formulas Use 1 mole of a gas at STP takes up 22.4L of space Volume of a gas at STP

Molar Volume Gases have a property that liquids and solids do not have. Under certain conditions, 1 mol of any gas has a volume of 22.4L. The two conditions that make this true are… A temperature of 0ºC A pressure of 1 atmosphere (atm)

Molar Volume Atmosphere (atm) is a unit for pressure.
The air pressure at sea level is a bout 1 atm. The two conditions previously mentioned are standard temperature and pressure (STP). For now it is important to know that 1 mol of any gas at STP has a volume of 22.4L. This value is called standard molar volume.

Molar Volume For gases at STP, standard molar volume, 22.4L/mol, is used to convert between gas volume and moles. This is a conversion factor that we can use in our factor label grid. IMPORTANT: This conversion factor only works at STP!

Molar Volume Example: What is the volume of 1.50 mol of CO2 @ STP?
You try… How many moles are in 75.3L of O2 at STP?

Use molarity (M) in moles / L
Molarity Conversions Volume of a solution Use molarity (M) in moles / L Number of Moles Number of particles Mass Percent composition, empirical and molecular formulas Volume of a gas at STP

The Molarity of Solutions
We discussed solutions in Unit 2. Solutions are homogenous mixtures. They consist of one or more solutes dissolved in a solvent.

Concentration and Molarity
Solutions can be concentrated or dilute, depending on the amount of solute and solvent. A concentrated solution has more solute compared to another solution with the same volume. A dilute solution has less solute compared to another solution. Think of making a pitcher of pink lemonade. Mmmmmmmmmmmmmmmmmmm

The directions state for you to add 4-1/3 cans of cold water to the concentrate. If you add 3 cans of water to the concentrate, then the lemonade will be very concentrated or strong. If you added 7 cans of water, the lemonade would be very dilute or weak. Strong and weak, concentrated or dilute are qualitative descriptions that are too vague for chemists to use so chemists use different units to describe concentration.

Chemists use units to measure concentration that make the measurement quantitative. The units of solution concentration is called molarity (M). Molarity is moles of solute / liters of entire solution The more moles of solute in a given amount of solution, the more concentrated it is and the higher the molarity.

Example The volume of an aqueous solution is 1.50L. It contains 12.5g of NaCl. What is the molarity of the solution? 12.5g NaCl is moles NaCl 0.214 moles NaCl in 1.50L of solution is… 0.143M NaCl Practice… The volume of a solution is 1.67L. It contains 39.0g of diatomic bromine. What is the molarity of the solution?

Today & This Week… If you would like extra credit, the next few lines will be extremely important. You can bring in a pumpkin and make it go “Boom!” It must be carved and cleaned out—but leave the eyes, nose & mouth pieces in place. The first 2 people per class to answer my riddle get to bring one in. Here is the deal… I will not accept an answer today—guesses should be turned in no earlier than tomorrow during brunch. Mr. Habs is the middle man. You must give him your guess with your name, hour, date, time and a spooky drawing. Only one guess per student per day. Oh yeah, quest is Thursday. Here is the riddle…What do you get when you cross a vampire with a snowman?

Mini Quiz How many grams are in 3.20 moles of calcium phosphate?
How many molecules are in 4.12 moles of calcium phosphate? How many molecules are in g of calcium phosphate? How many grams in 3.24 x 1025 molecules of calcium phosphate? What is the molarity of a solution that has 153.6g of Ca(NO3)2 in 1.2L of total solution? How many liters does 402g of O2 gas take up at STP?

More on Molarity The definition of molarity involves three factors, moles/L, moles and liters. If you know any two of the three you can solve for the third. Example… A 1.35M solution of KF has a volume of 1.33L. How many moles of solute does it contain? How many grams?

More on Molarity Another…
An ammonium chloride solution has a concentration of 0.573M. It contains mol of solute. What is the volume of the solution?

Use 22.4L molar volume at STP
Conversion Review Sol’n Volume Use molarity in FL Grid # of particles Molar Mass in FL Grid Avogadro’s # Moles Mass Use 22.4L molar volume at STP Gas STP

HW Problem How many moles of Cu are found in 155g of copper.
The density of copper is 8.92 g/cm3. How many atoms of copper are found in a sheet of copper that measures 8 inches by 12 inches. Assume the sheet is 2mm in thickness.

Today… Practice quiz… Notes % Composition and Empirical Formula—pick up from desk. HW sheet for tonight (U have) Lesson 3/4 + HW sheet due tomorrow Thursday Quest—big one over lots of stuff (20% old) - (80% new). Pumpkin Friday + lab? Extra Credit Winners—Tom,Pam & Alex

Objectives Get out your calculator, periodic table, sheet of paper and flow chart for conversions (page 1 of this unit) Convert 2.20 moles of sodium nitrate to grams—do not ask anyone for the formula for nitrate. Convert 350. grams of silver acetate to moles of silver acetate—on your own. How many molecules of sulfur dioxide are found in 11.7L of this gas at STP? How many moles of NaCl are contained in 255mL of 0.55M NaCl? How many grams?

Objectives Find the total molar mass for each element in a compound
Find the percent composition of a compound

Percent Composition Compounds have a definite composition, given by their formula and formula units. Water for example contains 2 atoms of H and 1 atom of O for every molecule of water. Each mole of water contains 2 moles of H and 2 moles of O for every mole of water. If we had exactly 1 mole of water there would be 2.02g of H and 16.0g of O. The total molar mass would be 18.02g.

Percent Composition This breakdown can be expressed as a percentage called percent composition. Percent composition tells the percentage by mass of each element in a compound. 2.02/18.02 for H = 11.2% 16.0/18.02 for O = 88.8% The two percentages add up to 100% or very close.

Percent Composition Try for carbon tetrafluodride
Try for calcium acetate. SF6 NH3 Plumbous carbonate Potassium nitrate A compound has a molar mass of 42.4g/mol. It contains elements X and Y. The total molar mass of element X is 6.9g/mol. What is the percentage of this compound? What is the percentage of the other?

Empirical and Molecular Formulas
An empirical formula shows the smallest whole number ratio of atoms in a compound. Sometimes an empirical formula is the compound’s chemical formula. If it is not, the subscripts in the empirical formula are important clues to the chemical formula.

Types of Formulas An example of an empirical formula is dextrose.
For every 1 mol of C there are 2 mol of H and 1 mol of O. A molecule of dextrose always has 6 atoms of C, 12 atoms of H and 6 atoms of O. C6H12O6 is the molecular formula of dextrose. A molecular formula gives the actual number of atoms in a molecule. Molecular formulas is synonymous with chemical formula.

Types of Formulas Both ionic and molecular formulas have empirical formulas. But only molecular compounds have molecular formulas. Ionic compounds have formula units.

Types of Formulas An empirical formula shows the simplest ratio of atoms in a compound. A molecular formula, or chemical formula, shows the actual number of atoms in one molecule. A formula unit shows the simplest ratio of cations to anions in an ionic compound.

Determining Empirical Formulas
A mole ratio is a ratio or fraction that compares the moles of one substance to the moles of another substance. Mole ratio for water is 2 mol H : 1 mol O

Steps: Assume you have a 100g sample of the given compound. If you know its % composition, you can change each percent symbol to grams. Keep the number the same, for example, 30% becomes 30g. (if you are given a mass for each element, skip the first step) For each element, convert grams to moles using molar mass. (treat all elements as monatomic and remember sig figs). Create a mole ratio using the mole amounts from step 2 by dividing all mole components by the smallest number of moles. The numbers in the simplest ratio of step 3 become the subscripts for the chemical formula. Write the empirical formula using these values.

BOO! We have the following to do for today…
Finish notes (man there are a lot of notes!) Lab Day is pushed until Monday  so that means that Quest day is now pushed back to Weds ! After notes we can say Whoa! With a fiery pumpkin face!

Example: What is the empirical formula of an unknown substance that contains 30.4% nitrogen and 69.6% oxygen by mass? What is the empirical formula of a compound that contains 80.0% carbon and 20.0% hydrogen by mass? What is the empirical formula of a molecular compound that contains 4.37g of phosphorus and 5.63g of oxygen? A compound contains 4.20g of carbon and 4.21 x 1023 atoms of hydrogen and 1.05 mole of oxygen. What is the empirical formula of this compound?

Molecular Formulas and Hydrates
A molecular formula is either the empirical formula or a multiple of the empirical formula. molecular formula = n (empirical formula) n must be a whole number. For the compound dextrose, n = 6. Empirical formula = CH2O Molecular formula = C6H12O6 = 6 (CH2O) Ex credit Q for the Quest…are dextrose and glucose the same thing? If not, how can they have the same molecular formula?

Determining Molecular Formulas
To find the molecular formula of a compound, you must know two things… the empirical formula the molar mass of the compound Think of dextrose again for example. Suppose you do not know its molecular formula, but you know its empirical formula and that the mass of 1 mol of dextrose is g/mol.

Determining Molecular Formulas
Calculate the mass of the empirical formula—this is not necessarily the molar mass of the compound. For dextrose… 12.01g/mol + 2(1.01 g/mol) g/mol = 30.02 g/mol Set up the following… Molar mass of compound = n (empirical formula mass) 180.0 g/mol = n (30.02 g/mol) n = 6 So multiply the subscripts of the empirical formula by 6 to get… C6H12O6

Examples The empirical formula of a compound is NO2. The molar mass of the compound is 92.0 g/mol. What is the molecular formula? The empirical formula of a compound is CH. The molar mass of the compound is 78.0 g/mol. What is the molecular formula? An unknown compound contains 2.17g of C, 0.362g of H and 0.966g of oxygen. Its molar mass is g/mol. What are its empirical and molecular formulas? An unknown compound contains 21.8g of phosphorus and 28.2g of O. Its molar mass is 284 g/mol. What are its empirical and molecular formulas?

Hydrates Some compounds have molecules of water trapped as part of the compound. Hydrates are compounds that are chemically combined with water in a specific ratio. The compound in a hydrate is usually on ionic compound. In a hydrate’s empirical formula, the water molecules are written a the end. For example, Ni(NO3)2· 6H2O shows that there are 6 water molecules with each formula unit of nickel (II) nitrate. The dot in the formula is NOT a multiplication sign.

Hydrates On a larger scale this tells us that there are 6 moles of water molecules for every mole of nickel (II) nitrate. When a hydrate is heated, the water molecules leave and mix with the surrounding air. What remains after heating is the ionic compound without the water. Its formula unit shows no water molecules. CuSO4·5H2O + heat  CuSO4

Hydrates Suppose you measure the mass of a sample of hydrate.
Then you heat it until the water is gone. You measure the mass of the compound that is left. You know the following: Mass of hydrate Mass of compound without water Mass of water in sample (hydrate mass – compound mass)

Hydrates From this information, you can find the empirical formula of the hydrate. It is based on the ratio of moles of the compound to moles of water.

Example A hydrate of magnesium sulfate is heated with the following results. What is the empirical formula of the hydrate? Mass of hydrate before heating = 5.65g Mass of compound after heating = 2.76g

Example A hydrate of sodium sulfide is heated with the following results. What is the empirical formula of the hydrate? Mass of hydrate before heating = 154g Mass of compound after heating = 50. g

Physics Sit down in quest seats and read the following…
Be silent during prayer and announcements—if you talk, you cannot work on the quest  Today’s Quest—you have the class period to finish it—but you cannot ask me any questions whatsoever. Calculator, equation sheet (not sheets) can be on your desk. This is the last grade for this 1st MP. There will be big changes on Monday…

Chemical Quantities The Mole.

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