Measurements of Matter Mass, volume, length, density… all these are measurements of matter. What is matter made of though? atoms If you measured the density of an object to be 2.04 g/mL, you have a certain number of atoms all joined together in a certain amount of space. So how many atoms are in there?

The MOLE It’s a beauty mark… It’s a small furry garden pest… No, wait… its how we count ATOMS!

What is a Mole and how is it used in Chemistry? A mole is a quantity of atoms. Since atoms are so tiny they are hard to individually measure, so scientist measure them in large quantities. Problem: How do you count something as small as an atom?

Solution: the mole Specific quantity (amount) of atoms in a substance. Basically, a word that represents a specific number (ex: pair = 2) 1 mole = 6.02 x 1023 representative particles (a lot of very tiny particles) This value is called Avogadro’s Number

What are representative particles? smallest particle that retains chemical and physical properties 3 types depending on the compound: Atoms: Single element Molecules: covalent compound Formula Units: Ionic Compounds or ions

What is a mole? Try this

How do we use it? Used to count groups of small particles for the purpose of mixing them with other particles in the lab. Ex: If 2.5 moles of sodium were used in an experiment, how many atoms of sodium were there? 1 mole = 6.02x1023 2.5 mole = ___________

Ex: Exactly 9.673 x 1019 molecules of water are needed to safely react in with sodium. How many moles of water are needed?

The Mole Used to count groups of small particles for the purpose of mixing them with other particles in the lab. Problem – no way to count a mole of anything. (atoms are too small to see) If we can’t count them – must be another way to measure them Solution – measure the mass of a mole with a balance (molar mass)

Using Molar Mass as a conversion factor The periodic table contains the molar mass of all elements. Molar Mass = mass of 1 mole (6.02x1023)of its atoms. 6.02x1023 atoms Carbon = 1 mole carbon = 12.01 grams Carbon If 2.5 moles of carbon were used in an experiment, how many grams of carbon must you weigh out?

Using molar mass How many grams are in 15.7 mole MgCl2

The Mole of a gas Why would it be difficult to measure the mass of helium? (how would it stay on the balance) Scientist have discovered that 1 mole of any gas has a constant volume at STP. (standard temperature & pressure) 1 mole = 22.4 L of any gas This value is called Molar Volume

Using molar volume as a conversion factor How many moles are in 12.95 L of oxygen gas? How many liters are in 0.758 moles of nitrogen gas?

Number of particles – Avagadro’s number What does a mole tell us? Number of particles – Avagadro’s number Mass particles in solid or liquid – Molar mass Volume of particles in a gas – Molar Volume

Molar Mass The mass of a substance.

The subscript is the number at the bottom of a formula. FORMULAS - review MgCl2 The subscript is the number at the bottom of a formula. There is 1- Mg & 2 - Cl

Calculate the molar mass Calculate the molar mass of 1 mole of magnesium chloride. (first you need a formula)

How to calculate molar mass Identify the # of atoms of each element Multiply # atoms by the atomic mass of that element. (round to 2 #’s after the decimal) Add them all together Grams (g) is the unit

MgCl2 Mg – 1 (24.31) = 24.31 Cl – 2 (35.45) = 70.90 or 95.21 g/mol Molar Mass MgCl2 Mg – 1 (24.31) = 24.31 Cl – 2 (35.45) = 70.90 or 95.21 g/mol 1 mole MgCl2 = 95.21 g MgCl2

1 mole of Ammonium phosphide Calculate the molar mass: (you MUST write the formula correctly before answering) 1 mole of Ammonium phosphide 1 mole of Trinitrogen pentachloride

What is a percentage mean? How is a percent calculated? What is the percentage of girls in the class? What is the percentage of students in Mrs. Muchnick’s class if she has 10 girls and 12 boys?

Percent Composition The percent BY MASS of each element in a compound – divide the element’s total mass (part) by the molar mass (whole) then multiple by 100 to get the percent. Ex: % composition of MgCl2 Mg – 1 (24.31) = 24.31 / 95.21 x 100 = 25.53% Mg Cl – 2 (35.45) = 70.90 / 95.21 x 100 = 74.47% Cl Molar mass = 95.21 g/mol (PART) (WHOLE) (PART) (WHOLE)

Practice - Calculate the % comp of KMnO4: K – 1 (39.10) = 39.10 / 158.04 x 100 = 24.74% K Mn – 1 (54.94) = 54.94 / 158.04 x 100 = 34.76% Mn O – 4 (16.00) = 64.00 / 158.04 x 100 = 40.50% O molar mass KMnO4 = 158.04

Calculating the amount of an element in a sample Find the % comp of the element in the compound Change the % to a decimal (move decimal 2 times to the left or divide by 100) Multiply that decimal by the amount (g) of the sample. Ex: Calculate amount of chlorine in 203.5 grams of MgCl2. (use the % we found earlier) 74.47% Cl = .7447 x 203.5 = 151.5 grams Cl

Practice: Calculate the amount of oxygen in 15.75 grams of water. Molar mass H2O = 18.02 / 18.02 x100 = 88.79% O 88.79% O = .8879 x 15.75 = 13.98 grams O Complete % comp worksheet

Using Molar Mass Remember – Molar mass is the mass (grams) of 1 mole 1 mole Fe = _________grams Fe 2.5 mole Fe = ________ grams Fe 113.5 grams Fe = _______ moles Fe Mass to mole = divide by molar mass Mole to mass = moletiply by mole mass 

Using molar mass In the lab, Mrs. Mathieson needs 2.57 moles of NaCl to do an experiment. How many grams would be needed to equal 2.57 moles of NaCl? After doing the experiment, Mrs. Mathieson has 1.02 moles of NaCl remaining – how many grams does that equal?

Empirical Formula The lowest whole number ratio (subscripts) of elements in a compound. Cannot be reduced!!! not empirical empirical Ex: C6H12O6  CH2O

Molecular Formula Actual number of atoms in a chemical compound EX: C12H24O12 Molecular Formulas can be reduced to Empirical Formulas molecular empirical EX: C12H24O12  CH2O Different molecular formulas can have similar empirical formulas molecular empirical EX: N3O9  N12O36  Molecular formula: C76H52O46 Empirical formula: ___________ NO3

C2H4 CH2 - empirical NO3 S9Cl12 C3Cl9 N4S9 S3Cl4 - empirical PRACTICE: 1. Identify each as empirical (can’t be reduced) or molecular (can be reduced) 2. If its molecular – write the empirical C2H4 NO3 S9Cl12 C3Cl9 N4S9 CH2 - empirical molecular empirical S3Cl4 - empirical molecular CCl3 - empirical molecular empirical

Finding Empirical Formula from Percent Composition Ex: A compound was found to be 54.53% Carbon, 9.15% Hydrogen, and 36.32% Oxygen. Find its Empirical Formula. Steps: Assume a 100g sample (change %  g) ÷ by the molar mass of that element to find moles (sig fig it!) Identify the lowest # of moles and ÷ them all by that number Round each to the nearest whole # The resulting whole #are the subscripts for that element in the empirical formula

Calculating Empirical Formula 63.5% Silver 8.2% Nitrogen 28.3% Oxygen 63.5 g Ag 8.2 g N 28.3 g O 107.87 14.01 16.00 .589 mole Ag .59 mole N 1.77 mole O .589 .589 .589 1 1 3 AgNO3

Calculating Empirical Formula (special) 60.00%C 4.48%H 35.53%O 60.00g C 4.48g H 35.53g O 12.01 1.01 16.00 4.996 mole C 4.44 mole H 2.221 mole N 2.221 2.221 2.221 2.249 2 1 x4 x4 x4 9 8 4 C9H8O4

Calculating Molecular Formula Find the empirical formula Calculate the molar mass of your empirical formula Identify the molar mass of your molecular (GIVEN in the problem every time!) Divide the molecular mass / empirical mass Round to the nearest whole # Multiply the whole # by the subscripts in the Empirical formula

Practice If a compound has an empirical formula of NO3 and a molecular mass of 186g – what is the molecular formula? Empirical formula: NO3 molar mass: 62.01g Molecular mass (given) 186g empirical mass 62.01 3 x NO3 = N3O9

How many f.u. are in 2.45 moles of NaCl? How many moles are in 5.90 x 1026 molecules of CO?