Unit 7 – Chemical Quantities & The Mole

Day 1: Mole Math We use a unit called the, or, to measure the amount of a substance. The mole can represent or molemol mass (grams), volume (liters), amount (particles).

**You must ALWAYS go through the MOLE!** 1 mole = Avogadro's number: and is also called a. One dozen = _________ items. 3 dozen cookies = _______ cookies 0.5 dozen doughnuts = _________ doughnuts A “dozen” is a counting unit equal to of any object. A “Mole” is a counting unit equal to of any object, even really small ones like molar mass or atomic mass 22.4 Liters 6.02 x particles Mole x atoms, molecules, or formula units.

What is so special about 6.02 x ? Why do scientists use that number? Allows us to manipulate many small particles, as if they were one whole part. Example Conversion Factors: 1 mole = 6.02 x particles (atoms, molecules, or formula units) 1 mole of copper = 6.02 x of copper. 1 mole of CuCl 2 = formula units of CuCl 2 1 mole of CO 2 = molecules of CO 2 The amount of a substance containing Avogadro’s number of any kind of chemical unit is called a of that substance. One mole of K contains atoms. One mole of NaOH contains formula units. Conversion factor atoms 6.02 x mole 6.02 x 10 23

Molar mass is the mass of of a substance. Other names for molar mass include… Molar Mass one mole *formula mass *gram formula mass

Molar Mass One _____________of any element will have a mass in grams corresponding to the value of its ____________________. – 1 mol carbon = __________ g/mol – 1 mol nitrogen = ___________ g/mol – (its atomic mass from the periodic table = 1 mole) mole atomic mass

Diatomic Elements never exist alone in nature!!! *Trick for remembering the 7 diatomics *

Molar Mass of Compounds One ________ of any molecule/compound will have a mass in grams corresponding to the value of its molar mass (the sum of the masses of the elements that compose it). mol

Example: Water, H 2 O ElementMolar Mass # of Atoms Total HOHO XXXX g g g/mol This is the mass of 1 mol of water!

Ex. Ca(NO 3 ) 2 Ca: x 1 = N: x 2 = O: x 6 = g/mole

Sample Problems Potassium carbonate  ___________ K 2 CO 3 # of K atoms: x = + # of C atoms: x = + # of O atoms: x = ___________________ g/mol K 2 CO 3

Sample Problems Ammonium sulfate  _____________ (NH 4 ) 2 SO 4 # of N atoms: x = + # of H atoms: x = + # of S atoms: x = + # of O atoms: x = _________________________ g/mol (NH 4 ) 2 SO

Practice 1. What is the atomic mass of sodium? 2. Calculate the molar mass of Al 2 (SO 4 ) 3 3. Calculate the molar mass of nitrogen (hint: diatomic!?!?).

Mole Highway NOTES: To convert between units, follow the highway. Notice, there is no shortcut from grams to liters or between any of the three units surrounding the mole. This means you have to convert to moles before converting to another unit!

Draw the Mole Road Map!!

Molar mass: called gram atomic mass when single element is used. called gram formula unit when ionic compound is used. called gram molecular unit when molecular compound or diatomic molecules used. Diatomic molecules are atoms that bond with themselves. There are SEVEN of these that you need to remember: Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2 (Remember this by the name “BRINClHOF”)

Determine the number of moles in…. 25 g sodium 25 g Na 1 mol Na g Na = 1.1 mol Na Molar mass of Sodium = Na = 1 x = g

Determine the number of moles in…. 85 g H 2 SO 4 1 mol H 2 SO g H 2 SO 4 = 0.87 mol H 2 SO 4 Mass in grams of H 2 SO 4 = H = 2 x = S = 1 x = O = 4 x = Add to get molar mass of g

Determine the number of grams in…. 2.5 moles of sodium 2.5 mol Na g Na I mol Na = 57 g Na Molar mass of Sodium = Na = 1 x = g

Determine the number of grams in… moles of H 2 SO mol H 2 SO g H 2 SO 4 I mol H 2 SO 4 = 49 g H 2 SO 4 Mass in grams of H 2 SO 4 = H = 2 x = S = 1 x = O = 4 x = Add to get molar mass of g

How many moles are there in 27 g of ethanol (C 2 H 5 OH)? 27 g 1 mol g = 0.59 mol Mass of C 2 H 5 OH: C = 2 x = H = 6 x 1.008= O = 1 x = Add these all up for mass in grams = g

Homework page 7 Show all work and units to receive full credit.

Representative Particles  Moles

It’s a quantity, and it’s a BIG one. Avogadro’s Number = ______________ and is also called a _______. Like a “pair” or “dozen,” a “mole” represents a set number of things; in chemistry, those “things” are particles. For example: 1 dozen=______ items 3 dozen=______ cookies.5 dozen=______ doughnuts 6.02 x Mole

It’s a quantity, and it’s a BIG one. So, a “dozen” is a counting unit equal to __12__of any object. Likewise, a “Mole” is a counting unit equal to _____________ of an object, even really small ones like ____________, ____________, or _____________ x atomsmolecules formula units

What’s so special about 6.02 x ? Why do scientists use that number? – A conversion factor. – Allows us to manipulate many small particles, as if they were one whole part. – For elements on the periodic table, there is a 1 to 1 relationship between the mass of a single atom (in amu) and the mass of 1 Mole of the same species of atom (in grams). – Remember, Mole is a quantity…6.02 x particles.

What’s so special about 6.02 x ? Using the periodic table, we’ve learned that an Argon atom has a mass of amu. Unfortunately, manipulating a single atom of any element isn’t reasonable, so taking its individual mass isn’t possible. However, we do have the Mole… Atomic Number (number of protons) Atomic Mass (average number of protons and neutrons) Ar 1 Atom of Argon1 Mole of Argon g amu So, a g sample of Argon contains 6.02 x atoms…that’s Avogadro’s Number, the MOLE! It’s a 1 to 1 relationship between a single atom in amu and a Mole of atoms in grams. This is the MOLAR MASS, or mass of a mole of a substance.

Representative Particles A representative particle is the smallest unit of a substance.

Types of Particles Monatomic elements = _______________ Diatomic elements = _________________ Ionic compounds = _________________ Covalent compounds = ______________ Ions = _________________ Acids = _________________ atoms molecules formula units molecules atoms molecules

Avogadro’s number, which is ________________, represents the number of “chemical units” in one mole of any substance. For the monatomic elements, the “chemical unit” is an __________. 1 mol of any chemical = ______________ particles x atom 6.02 x 10 23

Examples 1 mol CaCl 2 = 1 mol Ca 2+ = 1 mol HCl (aq) = 1 mol P 2 O 5 = 1 mol Ca = 1 mol Cl 2 = 6.02 x formula units 6.02 x atoms 6.02 x molecules 6.02 x atoms

1.How many moles are in 4.50 x 10 5 atoms of manganese? 4.50  atoms Mn 1 mol Mn 6.02  atoms Mn = 74.8 moles Mn Sample problems

2. How many atoms are found in 3.27 mol of magnesium? 3.27 mol 6.02  atoms 1 mol = 1.97  atoms Sample problems

3. Chalk is composed primarily of calcium carbonate. How many particles are in 3.4 moles of calcium carbonate? 3.4 mol 6.02  formula units 1 mol = 2.0  formula units Sample problems

Moles of Chalk Lab

Homework page 10 Show all work and units to receive full credit.

Notes: Volume Moles Warm-up: 1. How many atoms are in 0.75 mol of zinc?

Sample Problems (SHOW ALL WORK & UNITS to receive full credit.) 1. The average lung capacity of a male is 6.0L. The average lung capacity of a female is 4.7L. Assume the following: If your TEACHER’s lungs are completely filled with oxygen, determine the number of moles of oxygen gas in the lungs of your chemistry teacher at STP. 2. At STP, how many moles are found in 54L of neon gas? 3. How many liters are found in 3.02mol of helium at STP?

Homework page 12

Empirical vs. molecular

Empirical formula Is the simplest form of a formula Is written in the lowest possible ratio E.g., CH 2 O or CH Will be reduced. May or may not exist in this form in the real world.

Molecular formula Is the true formula or actual ratio of the atoms in a compound. E.g., C 2 H 4 O 2 or C 6 H 6 Will not be reduced. Formula describes a substance as it actually exists.

Empirical or Molecular? Na 2 O___________________ C 3 H 6 ___________________ K 2 SO 4 ___________________ C 6 H 12 O 6 ___________________ empirical & molecular molecular empirical & molecular molecular Keep in mind that for some compounds, its empirical formula can also be its molecular formula.

Which pair has the same empirical formula? Na 2 O and Na 2 O 2 C 6 H 12 O 6 and CH 2 O C 3 H 6 and C 5 H 12 C 6 H 6 and C 5 H 5

Calculating an Empirical Formula 1.Determine the mass of each element. 2.Convert the mass of each to moles. 3.Find the mole to mole ratios of each element by dividing the number of moles of each element by the smallest number of moles. 4.If the ratio is not a whole number, multiply each ratio by a factor to make them all whole numbers. 5.Write the formula using the mole ratio as the subscript for the formula.

Find the empirical formula of 69.5% O & 30.5% N. Step 1: Divide % or grams by its atomic mass to get moles of each element g N 1 mol N g N = mole N 69.5 g O 1 mol O g O = mole O

Find the empirical formula of 69.5% O & 30.5% N. Step 2: Divide smallest mole number in each element to get ratio of that element mole O mole N mole = 2 O = 1 N

Find the empirical formula of 69.5% O & 30.5% N. *This answer becomes the subscript for that element; round to nearest whole number if.8 or higher or lower than.2. Answer = NO 2

Independent Practice: 1. Analysis of a compound shows that it contains 10.88g of calcium and 19.07g of chlorine. Determine the empirical formula of this compound.

Independent Practice: 2. One of the most commonly used white pigments in paint is a compound of titanium and oxygen that contains 5.99 g titanium by mass and 4.01 g oxygen by mass. Determine the empirical formula and name for this compound.

Independent Practice: 3. Used in the production of nylon, adipic acid is an organic compound composed of 49.31% C, 43.79% O, and the rest is hydrogen. Determine the empirical formula of adipic acid.

Calculating Molecular Formulas The molecular formula will have the same ratio as the empirical formula. To determine a molecular formula, we will multiply the empirical formula by a whole number factor (WNF).

Calculating Molecular Formulas

1. Empirical Formula = P 2 O 5 Molar Mass= g/mol What is the molecular formula of this compound? g/mol g/mol = 2 empirical mass = g/mol (P 2 O 5 ) 2  P 4 O 10 WNF =

2. Nitrogen and oxygen form multiple molecular compounds together. One of these compounds is used to fuel space shuttles and has the empirical formula NO 2. If the molar mass of this compound is g/ mol, what is the molecular formula?

3.Butane is commonly used in lighters. It is composed of 17.37% hydrogen and 82.63% carbon. It has a molar mass of g/mol. What is the molecular formula of butane? **Fix typo in your packet, please.

4. Vitamin C is 40.91% C, 4.587% H, and the remaining is oxygen. If the molar mass of Vitamin C is about 180 g/mol, determine the empirical and molecular formula.

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