molar ratios from empirical formulas

Molar Ratios From Empirical Formulas

Empirical formula the smallest whole number ratio of

atoms (or moles of atoms) of each element in a substance

H2O 2 H atoms for every 1 O atom 2 moles of H atoms for every mole of O

atoms


The relative number of moles of each element in a substance can be used as a conversion factor called the molar ratio.

Molar ratio = moles element A mole of substance

Molar ratio = moles element Amoles element B

or


Fe2O3:

Molar Ratio = 2 moles of Femole Fe2O3

Molar Ratio = 3 moles Omole Fe2O3

Molar Ratio = 2 moles Fe3 moles O


Molar ratios can be used to determine the number of moles of a particular element in a given substance.

MolesA

MolesB

Molar Ratio


Example: How many moles of Na+ ions are present in 2.5 moles of Na2SO4 ?


Remember, once you find the number of moles of a substance present, you can use: Molar mass to find the number of grams Avogadro’s number to find the number of

atoms, ions, or molecules

Moles B

Grams B

Atoms B

Molar

mass

N

Moles A

MolarRatio


Example: What is the mass of iron present in 4.00 moles of Fe2O3?

Quantitative Information from Chemical Equations

Coefficients in a balanced equation number of molecules (formula units, etc) number of moles

2 H2 + O2 2 H2O

2 molecules 2 molecules1 molecules2(6.02x1023) molecules 6.02x1023 molecules 2(6.02x1023) molecules

2 moles 1 mole 2 moles


In industry, aspirin is prepared by:

C7H6O3 + C4H6O3 C9H8O4 + HC2H3O2

Salicylicacid

Aceticanhydride

Acetyl-salicylic

acid

Aceticacid

Chemists want to use the right amount of reactants in order to minimize “left over” reactants that contaminate the product.


The coefficients in a balanced chemical equation can be used to relate the number of moles of each substance involved in a reaction.

Molar ratios

mol reactant mol reactant mol productmol reactant mol product mol product

3 different molar ratios (and their inverses) can be written:

1 mol N2 1 mol N2 3 mol H2

3 mol H2 2 mol NH3 2 mole NH3

Molar Ratios from Chemical Equations

For the reaction:

N2 + 3 H2 2 NH3

Molar Ratios from Chemical Equations

MOLAR RATIOS = CONVERSION FACTORS

Moles of product that can be formed from a certain number of moles of reactant(s)

Moles of reactants needed to form a certain number of moles of product

Moles of reactant 2 needed to completely react with reactant 1


Example: If you have 1.0 mole of H2, how many moles of NH3 can you produce?

N2 + 3 H2 2 NH3

Note: Make sure your equation is balanced!


Example: If you have 1.0 mole of H2, how many moles of N2 will be required to completely react all of the H2?

N2 + 3 H2 2 NH3


Example: How many moles of N2 are needed to produce 0.50 moles of NH3?

N2 + 3 H2 2 NH3


Finding # of moles is great BUT

You don’t measure out moles in the lab!

Chemists use a balance to measure the mass of a substance used or produced in a reaction.

How can you determine the mass of reactants or products?


Use the molar mass to convert from moles to grams The number of grams of a substance

per mole


Mass (g) Compound A

Moles Compound A

Moles Compound B

Mass (g) Compound B

Molar

mass

Molar

mass

Molar ratio

“The MAP”


Example: How many grams of water will be produced by the complete combustion of 10.0 g of propane?

Combustion Reactions – Revisited!

You should be able to write a balanced equation for the combustion of an organic compound or a metal.

Organic compounds:

Metals:

CnHm + O2 CO2 + H2O

CxHYOn + O2 CO2 + H2O

Metal + O2 Metal oxide

Not bal.


Example: Hydrofluoric acid can’t be stored in glass because it attacks the silicates in the glass:

Na2SiO3 (s) + 8 HF (aq) H2SiF6 (aq) + 2 NaF (aq) + 3 H2O

How many grams of HF are needed to dissolve 55.0 g of Na2SiO3?

Making Bologna Sandwiches

Suppose you were going to make bologna sandwiches:

+ +

2 Bread + 1 Bologna 1 sandwich

Making Bologna Sandwiches

2 Bread + 1 Bologna 1 sandwich

+

We can only make 4 sandwiches because we don’t have enough bologna!

Bologna = limiting reagent or limiting reactant

Limiting Reagent or Limiting Reactants

Similar situations occur in chemical reactions when one of the reactants is used up before the others.

No further reaction can occur

The excess reactant(s) are “leftovers.”


2 H2 (g) + O2 (g) 2 H2O (l)

If we react 10 moles of H2 with 7 moles of O2, not all of the O2 will react because we will run out of H2 first!


10 H2 7 O2 10 H2O + 2 O2


Limiting reagent (limiting reactant): the reactant that is completely

consumed in a reaction

determines or limits the amount of product formed.


Three approaches to identifying the limiting reactant:

Compare the number of moles of each reactant needed with the number of moles of each reactant available

Calculate the number of grams of product that each reactant could formReactant that forms the least amount of product will be the limiting reagent.

OR


Three approaches to identifying the limiting reactant (cont.)

Pick one of the reactants and calculate the grams of the other reactant needed to exactly react with the one you picked

OR


First Method:

1) Convert the mass of each reactant to moles.

2) Pick one of the reactants (it doesn’t matter which one) and calculate the number of moles of the other reagent needed to completely react with the one chosen.

3) Compare the # moles needed vs. # moles available


Second Method:

1) Calculate the number of grams of product that each reactant could form.

2) Limiting reagent = reactant that forms the smallest mass of product.


Third Method:

1. Pick one reactant.

2. Calculate the grams of the other reactant needed.

3. Compare the grams needed with the grams available.


Example: If 10.0 grams of H2 are mixed with 75.0 grams of O2, which reactant is the limiting reagent?

2 H2 (g) + O2 (g) 2 H2O (l)


Method 1


Method 2


Method 3


Once you find the limiting reagent, you can also find the amount of product that can be formed in the reaction. If you used the second method for

identifying the limiting reagent, you’ve already done this!

Moleslimitingreagent

Molesproduct

gramsproduct

Molar

ratio

Molar

mass


Example: How many grams of water are produced by burning 5.00 g of methane in the presence of 5.00 g of oxygen?

CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (l)

Example

Since the mass of both reactants is given, you need to check to see if one of them is a limiting reagent.

Example

Yield

Theoretical Yield: the quantity of product that is

calculated to form when all of the limiting reagent reacts

Actual yield: the amount of product actually

obtained in a reaction

Yield

Often, the actual yield is less than the theoretical yield: reactants may not react completely

i.e. the reaction does not go to completion

“by-products” may formunwanted side reactions (competing reactions)

difficulty isolating and purifying the desired product

Yield

Percent Yield: relates the actual yield and the

theoretical yield

% Yield = actual yield x 100%theoretical yield

Yield

Example: In a certain reaction between H2

and CO to form methanol, the theoretical yield is 83.3 g of CH3OH. If the actual yield of the reaction was 81.5 g, what was the % yield?

molar ratios from empirical formulas

Documents

moles of n2

moles of atoms

moles of fe2o3

moles of nh3

moles fe3 moles omolar

moles of na2so4

certain number of moles

moles of na ions