Ch. 6 Notes -- Chemical Composition What is a mole?

Ch 6 – Chemical Quantities The Mole!!! A counting unit Similar to a dozen, except instead of 12, it’s 602 billion trillion… (602,000,000,000,000,000,000,000) ___________ (in scientific notation) This number is named in honor of Amedeo _________ (1776 – 1856), who studied quantities of gases and discovered that no matter what the gas was, there were the same number of molecules present…6.02 x 1023 6.02 x 10 23 Avogadro

Just How Big is a Mole? Enough soft drink cans to cover the surface of the earth to a depth of over 200 miles. If you had Avogadro's number of un-popped popcorn kernels, and spread them across the United States of America, the country would be covered in popcorn to a depth of over 9 miles. If we were able to count atoms at the rate of 10 million per second, it would take about 2 billion years to count the atoms in one mole.

The Mole 1 dozen cookies = ___ cookies 12 1 dozen cookies = ___ cookies 1 mole of cookies = ___________ cookies 1 dozen cars = ___ cars 1 mole of cars = __________ cars 1 dozen Al atoms = ___ Al atoms 1 mole of Al atoms = __________ atoms Note that the NUMBER is always the same, but the ______ is very different! Mole is abbreviated ______ . 6.02 X 1023 12 6.02 X 1023 12 6.02 X 1023 MASS mol

The Mole and Mass Mass in grams of 1 mole equal to __________ of the atomic masses. Practice problem: Calculate the mass of 1 mole of CaCl2. Ca = 1 x ________ g/mol = Cl = 2 x ________ g/mol = + = __________ g/mol CaCl2 1 mole of CaCl2 = the sum 40.1 40.1 g/mol 35.5 71.0 g/mol 40.1 g/mol 71.0 g/mol 111.1 111.1 g/mol

Mole Conversion Factors that you will need to know! 1 mol = __________ atoms/molecules/etc. 1 mol = _____________ grams 1 mol = ________ Liters of gas at STP (STP is Standard Temp. and Pressure… we will talk about what this means later!) 6.02 x 1023 ? ( molar mass) 22.4

Ch. 6 Notes -- Chemical Composition Practice Problems: (1) How many atoms of hydrogen are there in each compound? a) Ca(OH)2 ___ b) C3H8O___ c) (NH4)2HPO4 ___ d) HC2H3O2 ___ (2) Calculate the formula mass of each compound. (Add up all the atomic masses for each atom from the Periodic Table.) a) CaCO3 b) (NH4)2SO4 c) C3H6O d) Br2 e) H3PO4 f) N2O5 2 8 9 4 2 N’s = 2 x 14.0 = 28.0 8 H’s = 8 x 1.0 = 8.0 S = 32.1 4 O’s = 4 x 16.0 = 64.0 Ca = 40.1 C = 12.0 3 O’s =3 x 16.0 = 48.0 Add them up! 132.1 g/mol Add them up! 100.1 g/mol 159.8 g/mol C = 3 x 12.0 = 36.0 H = 6 x 1.0 = 6.0 O =16.0 2 Br’s = 2 x 79.9 = Add them up! 58.0 g/mol 2 N’s = 2 x 14.0 = 28.0 5 O’s = 5 x 16.0 = 80.0 3 H’s = 3 x 1.0 = 3.0 P = 31.0 4 O’s = 4 x 16.0 =64.0 Add them up! 98.0 g/mol Add them up! 108.0 g/mol

3) Convert 835 grams of SO3 to moles. 4) How many molecules of CH4 are there in 18 moles? 5) How many grams of helium are there in 5.6 x 1023 atoms of helium? 6) How many molecules are there in 3.7 grams of H2O? 1 mole SO3 835 g SO3 10.4 moles of SO3 x = 80.1 g SO3 6.02 x 1023 molecules CH4 18 moles CH4 x = 1.08 x 1025 molecules CH4 1 mole CH4 4.0 grams He 5.6 x 1023 atoms He x 3.72 grams He = 6.02 x 1023 atoms He 6.02 x 1023 molecules H2O 3.7 grams H2O 1.24 x 1023 molecules H2O x = 18.0 grams H2O

Calculating Percent Composition by Mass Step 1: Find the formula mass of the compound by adding the individual masses of the elements together. Step 2: Divide each of the individual masses of the elements by the formula mass of the compound. Step 3: Convert the decimal to a % by multiplying by 100. Practice Problems: (1) Find the % composition of the elements in each compound. a) Na3PO4 b) SnCl4 = 0.421 = 42.1% 3 Na’s = 3 x 23.0 = 69.0 P = 31.0 4 O’s = 4 x 16.0 = 64.0 ÷ 164 Sn = 118.7 4 Cl’s = 4 x 35.5 = 142.0 ÷ 260.7 = 45.5% = 0.189 = 18.9% ÷ 164 + ÷ 260.7 = 54.5% 260.7 = 0.390 = 39.0% + ÷ 164 164

Elements in the Universe: % Composition by Mass

Earth’s Crust: % Composition by Mass

Entire Earth (Including Atmosphere): % Composition by Mass

Human Body: % Composition by Mass

Determining the Empirical Formula for a Compound The empirical formula for a compound is the simplest __________ number __________ of the atoms in the compound. Examples: H2O is the empirical formula for water. _______ is the empirical formula for glucose, C6H12O6. Practice Problems: What is the empirical formula for the following compounds? a) C6H6= ________ b) C8H14O2 = ________ c) C10H14O2 = _________ d) Ca5Br10 = ________ e) N3O9 = ________ whole ratio C1H2O1 CH C4H7O C5H7O CaBr2 NO3

Determining the Molecular Formula for a Compound The molecular formula for a compound is either the same as the empirical formula ratio or it is a “_________ _________ of this ratio. It represents the true # of atoms in the molecule. Examples: 1) H2O is the empirical & molecular formula for water. 2) CH2O is the empirical formula for sugar, ethanoic acid, and methanol. The molecular formula for glucose is C6H12O6, (___times the empirical ratio!) Practice Problems: (1) If the empirical formula for a compound is CH2, which of the following is a possible molecular formula for the compound? a) C8H16 b) C8H8 c) C4H2 d) C3H9 (2) If the empirical formula for a compound is C2H3, which of the following is a possible molecular formula for the compound? a) C2H6 b) C10H15 c) C6H12 d) C8H14 whole # multiple 6

Determining the Molecular Formula for a Compound Find the molecular formula for C2H7 if the molecular mass of the compound is 93.0 g/mol. Find the molecular formula for P2O5 if the molecular mass of the compound is 283.88 g/mol. C2H7 = 31.0 g/mol 93.0 g/mol = 3 31.0 g/mol C2H7 x 3 = C6H21 141.94 g/mol P2O5 = 141.94 g/mol = 2 283.88 g/mol P2O5 x 2 = P4O10