Calculation of Empirical Formulae



Calculation of Empirical Formulae.

Example

1. A compound contains 75% carbon and 25% hydrogen. What is its empirical formula?

|Elements |Carbon |Hydrogen |

|Amount in question (% or mass) |75 |25 |

|Atomic mass (periodic table) |12 |1 |

|Number of moles = Amount in question |75 |25 |

|Atomic mass |12 |1 |

| |=6.25 |=25 |

|Mole ratio (divide each number of moles by the smallest number of moles) |6.25 |25 |

| |6.25 |6.25 |

| |=1 |= 4 |

| Empirical formula | C |H4 |

2. An oxide of carbon contains 27% carbon.

What is its empirical formula?

|Elements |Carbon |Oxygen |

|Amount in question (% or mass) |27 | |

|Atomic mass (periodic table) |12 | |

|Number of moles = Amount in question |27 | |

|Atomic mass |12 | |

| |= 2.25 | |

|Mole ratio (divide each number of moles by the smallest number of moles) |=2.25 | |

| |2.25 |2.25 |

| |=1 |= |

| Empirical formula | | |

3. Fluorspar is made of calcium and fluorine. If 51% is calcium, calculate the empirical formula.

|Elements |Calcium |fluorine |

|Amount in question (% or mass) | | |

|Atomic mass (periodic table) | | |

|Number of moles = Amount in question | | |

|Atomic mass | | |

| | | |

|Mole ratio (divide each number of moles by the smallest number of moles) | | |

| | | |

| Empirical formula | | |

4. 2.70g of aluminium is combined with 10.65g of chlorine. What is the empirical formula of the new substance?

|Elements | | |

|Amount in question (% or mass) | | |

|Atomic mass (periodic table) | | |

|Number of moles = Amount in question | | |

|Atomic mass | | |

| | | |

|Mole ratio (divide each number of moles by the smallest number of moles) | | |

| | | |

| Empirical formula | | |

5. 1.68g of iron is combined with 0.48g of oxygen. What is the empirical formula of the new compound ?

|Elements | | |

|Amount in question (% or mass) | | |

|Atomic mass (periodic table) | | |

|Number of moles = Amount in question | | |

|Atomic mass | | |

| | | |

|Mole ratio (divide each number of moles by the smallest number of moles) | | |

| | | |

| Empirical formula | | |

6. Saltpetre is a potassium salt. It is 13.9% nitrogen ,38.6% potassium and 47.5% oxygen. What is its empirical formula? (Hint- the table needs an extra column but you work out everything else in the same way)

|Elements | | | |

|Amount in question (% or mass) | | | |

|Atomic mass (periodic table) | | | |

|Number of moles = Amount in question | | | |

|Atomic mass | | | |

| | | | |

|Mole ratio (divide each number of moles by the smallest number of moles) | | | |

| | | | |

| Empirical formula | | | |

Extension- This one has an extra twist to it!.

7.A hydrocarbon has 80% carbon and 20 %hydrogen.

a) Calculate its empirical formula.

b) When the mass of a single molecule of this compound was analysed, it was

found to have a mass of 30g. What must the true (molecular) formula of this

substance be?

Calculation of Reacting Masses.

Example 1

Calculate the mass of iron sulphide produced when 5.6g of iron reacts completely with excess sulphur.

|Equation |Fe + S FeS |

|This means; |1mole of Fe reacts with 1mole of S to make 1mole of FeS |

|How many moles are we |We are using 5.6g Fe. |

|using? |Number of moles of iron being used is 5.6 = 0.1 Moles |

| |56 |

|How many moles of |Since 1mole of Fe reacts with 1mole of S to make 1mole of FeS |

|product will be made ?|Then 0.1 moles of Fe will produce 0.1moles of FeS |

|Amount of product |1 mole of FeS = (56+32) = 88g |

|produced. |So, 0.1 mole of FeS will have a mass of 0.1x 88 =8.8g |

Example 2

How much carbon would be needed to make 8.8g of carbon dioxide?

|Equation |C + O2 CO2 |

|This means; |__mole of C makes 1mole of CO2 |

|How many moles are we |We have made __g of CO2 |

|making? |1 mole of CO2 has a mass of _________g |

| |Number of moles of CO2 that have been made =Mass of CO2 made |

| |Mass of 1mole of CO2 |

| | |

| |= |

| | |

|How many moles of |Since 1mole of CO2 is made from 1mole of C, |

|product will be made ?|This means we must have used ______ moles of C. |

|Amount of reactant |Mass of carbon =Relative atomic mass of carbon x number of moles used |

|used |= |

| |= |

3.Calcium carbonate (CaCO3) decomposes when heated to give calcium oxide (CaO)and carbon dioxide (CO2)

a) Write a balanced equation for the reaction.

b)Calculate the mass of calcium carbonate needed to make 560g of calcium oxide.

4. Copper carbonate decomposes to give copper oxide and carbon dioxide.

a) Write a balanced equation for the reaction.

b)What mass of copper carbonate is needed to produce 8g of copper oxide?

5.How much calcium oxide is produced by heating 25 tonnes of calcium carbonate?

6.What mass of water is produced by completely burning 15kg of butane?

2C4H10 + 13O2 8CO2 + 10H2O.

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Empirical formula

CH4

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