Download Mole Ratio and Stoichiometry

Mole Ratio and Stoichiometry
Background information
In Chemistry, the mole is used as a base unit to measure amount of substances needed for chemical
reactions. Chemists have found that chemicals react with one another in proportion based on mole
ratios, and generally speaking, do not react proportionately when measure in other quantities such
as mass or volume.
The mole
Definition: The mole is the amount of substance which contains the same number of particles as the
number of atoms in exactly 12.0g of carbon 12.
The number of particles contained in one mole of substance = 6.0 x 1023 number of particles. This is
also known as avogardro constant
Avogardro constant
1 mole = 6.0 x 1023 number of particles
To calculate the number of moles of a substance:
No. of moles of substance =
𝒎𝒂𝒔𝒔 𝒐𝒇 𝒔𝒖𝒃𝒔𝒕𝒂𝒏𝒄𝒆 𝒊𝒏 𝒈𝒓𝒂𝒎𝒔
𝑴𝒓 𝒐𝒇 𝒔𝒖𝒃𝒔𝒕𝒂𝒏𝒄𝒆
Definition of Mr/Ar: Relative molecular mass (Mr) of an substance is the average mass of its
molecule compared to
th the mass of a carbon-12 atom.
To understand Mr and mole further, see the following illustration, not needed in exam.
Mass of
th of a carbon 12 atom : mass of 1 sodium atom : mass of 1 chlorine atom
1.66 x 10
1
-24
g
:
3.82 x 10
:
23
-23
g
:
5.89 x 10
:
-23
35.5
g
 Mr of substance
Mr of carbon = 1 x 12 = 12.
One carbon atom has Mr of 12, 12g of carbon will have 1 mole of substance by definition, which is
experimentally found to have 6 x 1023 number of carbon particles.
Practice 1
Calculate the following:
Moles of molecules in 20g of H2SO4:
Moles of H+ ions in 20g of H2SO4:
Number of molecules 100g MgSO4:
Number of Mg2+ ions in 100g of MgSO4:
0.25 moles of element X was found to have a mass of 28g. Calculate its atomic mass. Identify
element X.
Using Mole ratio
Mole ratio is extremely useful in predicting outcomes of chemical reactions. Since chemicals react in
proportion to mole ratios, given any chemical equation, one can calculate
i)
ii)
Amount of reactants needed
Amount of products produced
Example 1
Sodium chloride and hydrochloric acid reacts in the following reaction:
NaOH + HCl  NaCl + H2O
It is given that there is 20g of NaOH. Find the mass of HCl required to react with NaOH.
Calculate the number of H2O molecules produced in this reaction.
Example 2
What mass and number of moles of magnesium chloride is formed when 5g of magnesium oxide is
dissolved in excess hydrochloric acid?
Practice 1
What mass and moles of sodium chloride is formed when 21.2g of sodium carbonate is reacted with
excess dilute hydrochloric acid?
Practice 2
Given the equation
Na2CO3 (s) + 2HCl (aq)  2NaCl (aq) +H2O (l) +CO2 (g)
Calculate the mass of hydrochloric acid needed to make 2.3g of Carbon dioxide.
Percentage Composition
The percentage composition of a substance is the proportion by mass of a particular element in a
compound.
Percentage composition =
𝑴𝒓 𝒐𝒇 𝒔𝒖𝒃𝒔𝒕𝒂𝒏𝒄𝒆 𝒊𝒏 𝑪𝒐𝒎𝒑𝒐𝒖𝒏𝒅
𝑴𝒓 𝒐𝒇 𝑪𝒐𝒎𝒑𝒐𝒖𝒏𝒅
Or
Percentage composition =
𝑴𝒂𝒔𝒔 𝒐𝒇 𝒔𝒖𝒃𝒔𝒕𝒂𝒏𝒄𝒆 𝒊𝒏 𝑪𝒐𝒎𝒑𝒐𝒖𝒏𝒅
𝑴𝒂𝒔𝒔 𝒐𝒇 𝒄𝒐𝒎𝒑𝒐𝒖𝒏𝒅
Example 1
Calculate the percentage composition of water in hydrated magnesium sulphate MgSO4.7H2O.
Mr of MgSO4.7H2O = 24 + 32 + (4 x 16) + [7 x (1 + 1 + 16)] =246
Mr of water = 7 x 18 = 126
% of water = 126 / 246 x 100 %= 51.2%
Example 2
Calculate the mass of sulfur in 136g of CuSO4.
Percentage mass of sulfur =
= 20%
Mass of sulfur = 136 x 20%
= 27.2g
x 100%
Practice 1
Calculate the percentage composition of chlorine in CuCl2.
Practice 2
Calculate the percentage by mass of oxygen in C4H10O.
Practice 3
What is the mass of potassium in 286g of K3Fe(CN)6?
Empirical Formula and Molecular formula
Empirical Formula of a compound is the simplest whole number ratio of the different atoms in the
compound.
Steps to determine empirical formula:
1) Divide percentage mass or actual mass by its Ar to find the number of moles
2) Divide each value by the smallest value and round off to get a whole number
3) If the numbers obtained are no whole number, multiply by an integer to get whole numbers
Example 1
A molecule contains 88.89% oxygen and 11.11% hydrogen. What is its empirical formula?
% mass
Ar
Number of mole
Divide by smallest number
Empirical Formula
H
11.11%
1
11.11/1 = 11.11
11.11/5.55 = 2
O
88.89%
16
88.89/16 =5.55
5.55/5.55 = 1
H2O
Example 2
A chlorinated hydrocarbon compound when analysed, consisted of 24.24% carbon, 4.04% hydrogen,
71.72% chlorine. Find the empirical formula and write the molecular mass of this molecule
ii)
It is found that the molecular mass is 99 from another experiment, find the actual
formula of the molecule.
The above formula is the molecular formula.
Molar Volume
Definition: One mole of any gas occupies a volume of 24 dm3 at room temperature and pressure
The volume of 24dm3 is called the molar volume of gas at room temperature
To find the volume of gas given the number of moles at room temperature pressure:
Volume of gas (dm3) = no. moles of gas X 24dm3
Example 1
What is the volume of 3.5 g of hydrogen at room temperature?
Example 2
Given the equation
MgCO3 (s) + H2SO4 (aq)  MgSO4 (aq) + H2O (l) + CO2 (g)
Calculate the mass of magnesium carbonate needed to make 6 dm3 of carbon dioxide at room
temperature.
Practice 1
A small teasppon of sodium hydrogencarbonate (baking soda) weighs 4.2g. Calculate the moles,
mass and volume of carbon dioxide formed when it is thermally decomposed in the oven.
NaHCO3 (s)  Na2CO3 (s) + H2O (g) + CO2 (g)
Practice 2
Hydrogen peroxide decomposes to form water and oxygen gas. Form a balanced equation and
calculate the volume of oxygen gas produced when 17g of hydrogen peroxide fully decomposes.
Molar Solution
Concentration of a solution shows the amount of solute dissolved in 1 dm3 of solvent. There are 2
ways to measure concentration mainly concentration and molarity.
Concentration =
Molarity (M) =
𝑴𝒂𝒔𝒔 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒆 𝒈𝒓𝒂𝒎𝒔
𝑽𝒐𝒍𝒖𝒎𝒆 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒊𝒐𝒏 𝒅𝒎𝟑
𝑴𝒐𝒍𝒆 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒆 𝒎𝒐𝒍𝒆
𝑽𝒐𝒍𝒖𝒎𝒆 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒊𝒐𝒏 𝒅𝒎𝟑
Example 1
250 cm3 of solution contains 5.95g of potassium bromide. Calculate the concentration of the
solution in (a) g/dm3 and (b) mol/dm3
Example 2
Calculate the number of moles of
(a) HCl in 4 dm3 of 0.3 mol/dm3 HCl
(b) NaOH in 30cm3 of 0.5 mol/dm3 NaOH
If (a) and (b) were to mixed together in a beaker, will the resulting solution be acidic, alkaline or
neutral?
Practice 1
Given a solution of NaOH with a concentration of 0.5mol/dm3, find the volume of NaOH required to
completely react with 30cm3 of HCl with a concentration of 0.35mol/dm3.
Practice 2
A chemist starts with 50cm3 of a 0.40M NaCl solution and dilutes it to 1000cm3 of water. What is the
concentration of NaCl in the new solution?
Percentage Yield
There is 2 forms of yield.
1) theoretical yield which is the amount of product obtained at the end of a reaction by means
of calculation from chemical equation.
2) Experimental or actual yield which is the amount of product obtained at the end of a
reaction by means of measurement from an actual experiment.
In general, these two yields give a different number due to several factors such as purity of products,
incomplete reaction, or environmental factors.
Actual yield will always be lesser than theoretical yield.
Percentage yield =
𝑨𝒄𝒕𝒖𝒂𝒍 𝒀𝒊𝒆𝒍𝒅
𝑻𝒉𝒆𝒐𝒓𝒆𝒕𝒊𝒄𝒂𝒍 𝒀𝒊𝒆𝒍𝒅
x 100%
Example
Magnesium reacts with dilute hydrochloric acid to produce magnesium chloride.
a) What is the maximum theoretical mass of MgCl2 that can be produced from 12.0 g of
magnesium?
b) If only 42.0g of purified MgCl2 was obtained after crystallizing the salt from the solution,
what is the percentage yield of the salt preparation?
Percentage purity
Percentage purity is useful for identifying the amount of impurity present in a chemical substance
Percentage Purity =
𝑴𝒂𝒔𝒔 𝒐𝒇 𝒑𝒖𝒓𝒆 𝒔𝒂𝒎𝒑𝒍𝒆
𝑴𝒂𝒔𝒔 𝒐𝒇 𝒔𝒂𝒎𝒑𝒍𝒆
x 100%
Example
When 3.05 g of zinc was heated in excess oxygen, 3.5 g of zinc oxide was obtained. Calculate the
percentage purity of zinc.
Practice
In the reaction given below, 5.40 kg of impure CaF2 was treated with 8.80kg of H2SO4 to yield 1.54kg
of HF.
CaF2 (s) + H2SO4 (l)  CaSO4 (s) + 2HF (l)
a) Calculate the mass of CaF2 that reacted in the reaction
b) Percentage purity of CaF2