THE MOLE

CHAPTER 10

I can:

State with 100% accuracy Avogadro's number as 6.02x1023

particles Apply Avogadro's number to convert

between particles of an element and grams of an element with increasing accuracy (mole conversions)

What the heck is a mole?

Is it:a. A Mexican sauce.b. A skin imperfection.c. A small, burrowing mammal. d. A unit for counting in chemistry.

All of the above!

BIG IDEA

Chemists use the mole to count the number of atoms, molecules, ions and compounds they work with.

Origin of “mole”

From the German word “mol” which is short for molekulargewicht (molecular weight) and the French word moléculaire (molecule)

Why do we use the mole?

Chemists need a more convenient method for accurately counting the number of atoms, ions, molecules and compounds in their work.

Other ways of naming various quantities

Century = Week =Gross =Ream = Dozen = Acre =

How much is a mole?

1 mole = 6.02 x 1023 particles OR

602,213,670,000,000,000,000,000

A particle may be an atom, ion, molecule, compound, etc. Ex: Na, Na+, NaCl, Cl2

How much is a mole?

6.02 x 1023 is called Avogadro’s number.

Amadeo Avogadro- Italian physicist

(1776-1856)In 1811 determined the

number ofatoms of 1 mole of a gas

to be 6.02 x 1023 atoms.

How much is a mole?

1f one mole of pennies were divided up and given to every person on the earth, each person would receive 1.5 x 1016 pennies.

At a spending rate of 1 million dollars per day, it would take each person 4 thousand years to spend a mole of pennies.

How much is a mole?

1 mole of basketballs would fill up a ball bag the size of the earth.

1 mole of marshmallows would cover the US to a depth of 600 miles.

Quantity Conversions

If you have 2 dozen roses, how many roses do you have?

If you have 3.6 dozen chocolates how many chocolates do you have?

If you have 82 roses, how many dozen roses do you have?

Atoms, Molecules, Formula Units

“Representative Particles” are either atoms, formula units or molecules

Atoms ElementsFormula Units (F.U.’s) Ionic

Compounds (is there a metal present?)

Molecules Covalent Compounds (nonmetals only)

Your conversion will always be either 1 mol/6.02 x 1023 or vice versa

Mole Conversions

How many sodium atoms are in a mole?

If you have 2 moles of sodium atoms, how many sodium atoms do you have?

If you have 3.5 moles of sodium atoms, how many sodium atoms do you have?

Mole Conversions

If you have 4.9 x 1023 chlorine atoms, how many moles of chlorine atoms do you have?

If you have 22.7 x 1027 chlorine atoms, how many moles of chlorine atoms do you have?

If you have 5.5 x 1021 H2O formula units, how many moles of H2O formula units do you have?

Why Use % composition?

BIG IDEA: All chemists who create or manufacture compounds/chemicals need to analyze the stuff/compounds they make to be sure it is what they say it is. Examples: pharmaceuticals, paints, food

additives % composition data is determined using a small

sample of the solid the chemist has synthesized. The chemist than takes the % composition data

and works BACKWARD to determine the molecular formula of the compound.

Thermal Gravimetric Analysis for % Composition Determination

Why Use % composition?

Example: You have a compound which was analyzed and found to have the following composition by mass: 27.37% Na 1.20% H 14.30 % C 57.14% O

Working backward a chemist can determine the formula of the compound to be NaHCO3 or baking soda.

Why Use % composition?

Example: You have 2 different ores which contain gold. You only have the time, money and equipment to mine for one of the ores. You need to decide which ore contains the most pure gold. 350 lbs. of calaverite containing 40% by

mass Au. 480 lbs. of petzite containing 25% by

mass Au. Which of the 2 ores do you mine and

process?

Why Use % composition?

Example: A food processing plant manufactures different sweeteners- glucose (C6H12O6 ) and sucrose (C12H22O11). Some numbskull in the lab got the two products mixed up. The % composition analysis shows the following:

% carbon = 40.0%% hydrogen = 6.7%% oxygen = 53.3 %

Is the product sucrose or glucose?

Empirical Formulas

A formula giving the proportions of the elements present in a compound but not the actual numbers or arrangement of atoms.

Finding Empirical Formulas

Please Write this down and use often!!!

Steps to finding an empirical formula Change all percents to grams Convert all grams to moles Divide by the smallest number of moles If all your numbers are whole, you are

done▪ If a number ends in .5, multiply all answers by

2▪ If a number ends in .33 or .66, multiply all

answers by 3

Finding the molecular formula from the empirical formula

To find the molecular formula from the empirical formula, they must give you the molar mass or the molecular formula.

Divide the molar mass of the molecular formula by the molar mass of the empirical formula.

Take that number and multiply by everything within the molecular formula.

Chemical Hydrates

BIG IDEA: Hydrates are solid ionic compounds in which molecules of water are trapped.

Dessicants are chemical compounds which absorb water to form a hydrate. Examples: SiO2, CaCO3 – these are

dessicants added to products such as shoes and electronics before shipment overseas.

Chemical Hydrates

Formulas of hydrates contain a chemical formula followed by a dot and the # of H2O molecules for the hydrate.

Example:Na2CO3•10H2O called sodium

carbonate decahydrate

NH4OH•H2O called ammonium hydroxide

monohydrateCaCl2•2H2O called calcium

chloride dihydrate

Review of the Prefixes

1 – Mono 2 – Di3 – Tri 4 – Tetra5 – Penta 6 – Hexa7 – Hepta 8 – Octa9 – Nona 10 - Deca

Chemical Hydrates

Practice CuSO4•5H2O FePO4•4H2O MgSO4•6H2O CaCl2 •2H2O MgSO4 •7H2O Cd(NO3)2 • 6H2O Calcium hydroxide octahydrate Calcium chloride

trihydrate barium hydroxide octahydrate lithium chloride

tetrahydrate barium chloride dihydrate sodium sulfate

pentahydrate

Finding Hydrate Formulas Weigh the hydrate (in problems, this is

given to you) Drive off the water and re-weigh what’s

left. (Once you’ve driven off the water, the compound is now called an “anhydrate” or “anhydrous”)

Calculate the weight of the compound and the weight of the water driven off.

Convert both to moles Come up with a ratio of water to

compound

Example 1

A 15.67 g sample of a hydrate of magnesium carbonate was heated, without decomposing the carbonate, to drive off the water. The mass was reduced to 7.58 g. What is the formula of the hydrate?

Example 2

A hydrate of Na2CO3 has a mass of 4.31 g before heating. After heating, the mass of the anhydrous compound is found to be 3.22 g. Determine the formula of the hydrate and then write out the name of the hydrate.

Example 3

A 5.00 g sample of hydrated barium chloride, BaCl2 · nH2O, is heated to drive off the water. After heating, 4.26 g of anhydrous barium chloride, BaCl2, remains. What is the value of n in the hydrate's formula?

STOICHIOMETRY

What Is Stoichiometry?

BIG IDEA: Stoichiometry is the study of quantitative relationships between the amounts of reactants used and the amounts of products formed by a chemical reaction.

The amount of each reactant present at the start of a chemical reaction determines how much product can form.

Based on the law of conservation of mass!

How is Stoichiometry useful?

Using stoichiometry allows chemists and engineers to calculate how much reactant or starting materials they must buy and how much product they will produce BEFORE the reaction even takes place.

Chemicals and materials are expensive! Knowing how much you will need and how much you will produce is more efficient and saves money!

Where Did We Get This Word From?

STOICHIOMETRY (or stoich for short) derives from the Greek word “stoikheion”, meaning “element” + “-metry”, meaning to measure.

Stoichiometry requires the use of balanced chemical equations and molar masses of chemical compounds!

Mole Ratio

A mole ratio is the ratio between the numbers of moles of any two of the substances in a balanced chemical equation.

The mole ratio comes from the coefficients in the balanced chemical equation.

The mole ratio relates the masses of reactants to products using the molar masses.

Mole Ratio

Example:4Fe (s) + 3O2 (g) ―> 2Fe2O3

(s)

4 moles of Fe and 3 moles of O2 are required to produce 2 moles of 2Fe2O3.

In other words- moles of Fe to O2 to Fe2O3 can be represented using the mole ratios or fractions- 4Fe 3O2

3O2 2Fe2O3

Why do reactions stop?

• Reactions proceed until one of the reactants is used up, leaving the other reactant(s) in excess.

The limiting reactant limits the extent of the reaction and, thereby, determines the amount of product formed.

The excess reactants are all the leftover unused reactants.

Steps in finding the limiting reagent Pick one of the products (if there’s only one, use

that one) – perform a gram-to-gram calculation from both reactants over to that product (two separate problems)

Whichever answer is smaller, that reactant that you started with is the limiting reagent (the other is the excess reagent)

To find the amount of excess reagent used, start with grams of the LR and do a gram-to-gram from the limiting to the excess reagent. If they want how much of the excess reagent is left over, you subtract the answer from the original amount.

LR and ER can also be called “limiting reactant” and “excess reactant”

Stoichiometry Calculations

Silicon nitride, Si3N4, is used in the manufacturing of high-temperature thermal insulation for heat engines and turbines. It is produced by the following chemical reaction:

__Si + __N2 ―> __ Si3N4

1. How many grams of Si3N4 can you produce if you start with 100 g of silicon, Si?

2. How many grams of Si3N4 can you produce if you start with 100 g of nitrogen, N2?

3. If you have 100 g of Si and 100g of N2, how much Si3N4 can you actually produce? _________________ What will the limiting reactant be? ________________

Percent Yield

% yield- a measure of the efficiency of a chemical reaction.

Theoretical yield- the use of stoichiometry calculations to predict the amount of product you can obtain from the reactant(s).

Actual yield- the amount of product produced when the chemical reaction is actually carried out in the lab.

% Yield = Actual Yield X 100 Theoretical Yield

Practice

Solid silver chromate (Ag2CrO4) forms when potassium chromate (K2CrO4) is added to a solution containing 0.500 g of silver nitrate (AgNO3). Calculate the amount of Ag2CrO4 produced.

__K2CrO4 + __AgNO3 ―> __Ag2CrO4 + __KNO3

Calculate the % yield for the reaction if you actually produce 0.455 g Ag2CrO4 in the lab.

If you start with 0.500 g of AgNO3 and 0.500 g of K2CrO4 how much Ag2CrO4 can you actually produce? Does this change your % yield you calculated above?