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4. Stoichiometry
Formulae of simple compounds
Formulae of simple compounds can be deduced from their ions but there
are some that you should know off by heart.
You must learn the following formulae:
Carbon dioxide
Water
Hydrogen
Sodium chloride
Ethanol
Sodium hydroxide
Nitric acid
Sulfuric acid
Hydrochloric acid
Ammonia
Naming Compounds
When the compound contains and metal and a non metal;
the name of the metal is given first then the non metal but ending in –ide.
E.g. Sodium chloride, Magnesium oxide, iron sulfide.
When the compound is made up of two non-metals;
If one is hydrogen, that is named first,
Otherwise the one with the lower group number comes first then name of the other non-metal with
ide on the end.
e.g. Hydrogen chloride, carbon dioxide
Valency
The valency of an element is the number of electrons its atoms lose, gain or share, to form a
compound.
Group
I
II
III
IV
V
VI
VII
0
Hydrogen
Transition e’s
Lose / Gain
electrons
Valency
Examples
Using Valencies to write formulae
A
Write the formula of aluminium oxide
Model answer:
Symbol
Valency
Al
3
O
2
“swap” the numbers and if possible, cancel.
AL2O3
B
Write the formula for aluminium phosphide
Model answer:
Symbol
Valency
Al
3
P
3
“swap” the numbers and if possible, cancel.
Al3P3, cancel numbers: Al1P1, written as AlP.
Write the formula of the following simple compounds.
1.
lithium oxide
2.
sodium oxide
3.
boron fluoride
4.
calcium nitride
5.
germanium chloride
6.
arsenic oxide
7.
aluminium sulphide
8.
silicon carbide
EC p49/51
Balancing charges to write formulae
!
!
You are expected to know the formulae of all the ions listed below, there will be a test
at some point. You have been warned!
!
The next few pages are worksheets that will test your skills on deducing formulae and balancing
chemical equations. Please do them yourself!
Writing Chemical Equations
Now that we can write chemical formulae, we must now begin to use them in
chemical equations. There are 3 types of equations that you will have to be
comfortable with:
- Word Equations
- Symbol Equations, just called “Equations”
- Ionic Equations
Word equations
Here we simply name the elements and compounds that are involved in the reaction…
Carbon + Oxygen
Carbon Dioxide
(Symbol) Equations
This type of equation is the most common and involves using the chemical formulae
of the substances involved…
HCl + NaOH
NaCl + H2O
We must learn to balance these equations, and there are some simple guidelines…
• These equations must have the same number and type of symbols on each side of
the equation (we cannot destroy the elements involved or mass!!)
• If we want to change the number of symbols on either side, we must put large
numbers in front of the formulae (we cannot change the ratios within the formulae
themselves).
EG.
CuO + HCl
EC p53 Q1-3
Ionic Equations
These are perhaps the most tricky but the most meaningful equations in Chemistry.
Ionic substances tend to break up when dissolved in water and ionic equations show
only the ions that take part in the chemical reaction…
EG
BaCl2 + MgSO4
NB. State symbols are important in these equations, and can be:
(s) Solid
(l) Liquid
(g) Gas
(aq) Aqueous
EC p55 Q1-2
End of chapter questions
Chemical Calculations
Relative atomic mass definition – Ar (sometimes written RAM)
Relative atomic mass and isotopes
Chlorine has a relative atomic mass of 35.5! How can an element contain half of a
neutron? The answer is due to the isotopes of chlorine and their relative abundance.
!
There are 3 times as many 35Cl atoms as there are 37Cl. So by doing a simple
calculation the average relative atomic mass is a decimal. Although you won’t be
asked to do this, the following method is used to determine the Ar.
Finding the masses of molecules and ions
Using Ar values, it is easy to work out the mass of any molecule or group of ions.
Relative Molecular Mass (RMM) = Mr
The combined mass of all the atoms present in the substance
Relative Formula Mass (RFM) = Mr
The combined mass of all the ions present in the substance
Ions have the same mass as atoms as the mass of an electron is negligible
Hydrogen (RMM) Mr :
Water (RMM) Mr :
Sodium Chloride (RFM) Mr :
Work out the Mr of the following substances:
1.
NH3
2.
CH4
3.
C6H12O6
4.
CuSO4
EC p59 Q1-3
Calculations involving reacting masses
If you know the actual amount of two substances that react you can:
- Predict other amounts that will react
- Say how much product will form.
Two laws of chemistry you must remember:
1
Elements always react in the same ratio to form a given compound.
This “recipe” is represented by the chemical equation for the reaction:
Eg. Carbon and Oxygen
12g of carbon combines with 32g of oxygen, and so on…….
6g of carbon combines with 16g of oxygen to form carbon dioxide, so……
2
The total mass does not change during a chemical reaction.
Total mass of reactants = Total mass of products
6g of carbon combines with 16g of oxygen to form 22g carbon dioxide,
12g of carbon combines with 32g of oxygen to form 44g of carbon dioxide
Calculating Quantities
Example 1: A student obtains 48g of magnesium sulfate from 9.6g of magnesium. What
mass of magnesium sulfate can the student get from 1.2g of magnesium?
9.6g of magnesium gives 48g of magnesium sulfate (write down the info!!!)
so.
1.2g of magnesium gives 1.2 x 48 = 6g magnesium sulfate
9.6
Example 2: In the reaction Mg + CuSO4 ➔ MgSO4 + Cu. 6.4g of copper are formed
from 2.4g of magnesium. What mass of magnesium is needed to get 32g of copper?
6.4g of copper is formed from 2.4g of magnesium (write down the info!!!)
so.
32g of copper requires 32 x 2.4 = 12g of magnesium
6.4
The Mole
An atom is impossible to weigh on a balance. 1000 atoms are impossible to
weigh on a balance, a million, a billion etc. So we scale it up to a specific
number of atoms, ions or molecules. An atom or molecule can be called a
unit, they count as 1 !
This number is 6.02x1023 . We call this number Avogadro’s constant.
In one mole of helium atoms there are 6.02x1023 helium atoms
In one mole of water there are 6.02x1023 water molecules
And so on..
but be careful..
In one mole of water there are 6.02x1023 water molecules but there are
6.02x1023 x 3 atoms!
That means there are 18.06x1024 atoms in a mole of water.
Draw below:
Molar Mass - The mass of 1 mole
You can find the mass of one mole of any substance by these steps:
1.
2.
3.
Write down the symbol or formula of the substance
Find its Ar or Mr
Express that mass in grams (g).
For this you will need to remember the formula of common compounds and
which elements exist naturally as molecules. Oxygen exists as O2 for
example.
Complete the table:
Substance
Symbol or
formula
Ar
Mr
Mass of 1 mole
Helium
He
He = 4
Exists as single
atoms
4g
Oxygen
O2
O = 16
2 x 16 = 32
32g
Ethanol
C2H5OH
C = 12
H=1
O = 16
Water
Carbon Dioxide
Copper
sulphate
Sodium
chloride
H20
5.
6.
2 x 12 = 24
6x1=6
1 x 16 =
16
46g
You will need to learn this equation and/or the calculation triangle
!
Number of moles (in a given mass) =
mass
mass of 1 mole
or
Mass of a given number of moles
= mass of 1 mole x number of moles
Eg.
EC p61 Q1-2
Moles Worksheets
1)
Define “mole”.
2)
How many moles are present in 34 grams of Cu(OH)2?
3)
How many moles are present in 2.45 x 1023 molecules of CH4?
4)
How many grams are there in 3.4 x 1024 molecules of NH3?
5)
How much does 4.2 moles of Ca(NO3)2 weigh?
6)
What is the molar mass of MgO?
1.
How many moles of Na are there in 42 g of Na?
2.
How many moles of O are there in 8.25 g of O?
3.
How many moles of O2 are in 8.25 g of O2
4.
What is the mass of 0.28 mol of Iron?
5.
How many atoms are in 7.2 mol of chlorine?
6.
How many atoms are in 36 g of bromine?
7.
How many moles of CO molecules are in 52 g of CO?
8.
How many moles of C2H6 are in 124 g?
Calculations involving equations
Most questions you will be asked to answer will have an equation. It is very
important you understand how to use moles and ratios.
Consider the following reaction:
Mg
+
2 HCl
➔
MgCl2
+
H2
IMPORTANT !!!
- Those big numbers in front on the formulae tell us in what proportions (in
moles) the substances react. That is why we balance equations.
So..
1 mole of magnesium reacts with 2 moles of hydrochloric acid to give 1
mole of magnesium chloride and 1 mole of hydrogen.
Question:
How many moles of magnesium chloride are produced when 5 moles of
magnesium reacts with an excess of hydrochloric acid?
Mg
+
2 HCl
➔
MgCl2
+
H2
Answer:
Look at the ratio of Mg to MgCl2 1 mole of Mg, gives 1 mole of MgCl2
So 5 moles of Mg must give us 5 moles of MgCl2
Question:
If we react 36g of Mg with an excess of hydrochloric acid, what mass of
MgCl2 is produced?
Remember the big number are for proportions in moles NOT MASS!
You must work out the number of moles first BEFORE you use the ratio
(copy the model answer from board)
Mg
+
2 HCl
➔
MgCl2
+
H2
200g of calcium carbonate reacts with an excess of hydrochloric acid.
1.
2.
3.
4.
Write the balanced equation
Calculate how many moles of calcium carbonate there are
Deduce how many moles of carbon dioxide are produced
Calculate the mass of the carbon dioxide produced.
123g of glucose reacts with an excess of oxygen.
1.
2.
Write the balanced equation for glucose reacting with oxygen.
(Respiration)
Calculate the mass of water produced.
“Limiting reagent” (reactant) and reagent “in Excess” (INXS)
When a chemical reaction occurs, unless we have the exact proportions of each
substance required for the reactants to be perfectly used up, one of the substances will
be used up before the other one(s) involved.
The reactant which is completely used up is called the “limiting reagent”
We say that the reagents that are left over once the limiting regent was used up we “in
excess”.
EC p63 Q1-2
Molar gas Volume
1 mole of any gas occupies 24dm3 at RTP.
That means that 1 mole of oxygen, carbon dioxide, hydrogen or methane
will 24dm3 of space at RTP. (room temperature and pressure)
Remember that the mass of 1 mole for different gases is not the same!
!
The volume is the same but the mass is different.
The equation you need to learn is:
!
Example Simple Calculation:
What volume does 0.25 moles of a gas occupy at rtp?
1 mole occupies 24dm3 so
0.25 moles occupies 0.25 X 24dm3 = 6dm3
1 dm3 = 1 litre.
In 1dm3 there are 1000cm3
Reacting Masses / Gas Calculations
The steps involved in a calculation are as follows :
(a) Convert the information given to moles of one substance.
(b) Use the chemical equation to find moles of other substance needed.
(c) Convert back from moles to mass (or concentration, volume etc.)
1. When calcium carbonate is heated, carbon dioxide is evolved:
CaCO3 ! CaO + CO2
What volume of carbon dioxide (at RTP) is produced from 500g
calcium carbonate? [3]
2.
500g calcium carbonate was treated with hydrochloric acid
at RTP:
CaCO3 + 2HCl ! CaCl2 + H2O + CO2
What volume of CO2 gas was produced? [3]
3.Ammonia is produced in the Haber process according to the
equation;
N2 + 3H2 ! 2NH3
What volume of (i) nitrogen and (ii) hydrogen is required to
produce 68g of ammonia, at RTP?
[6]
4. When carbon and carbon dioxide are heated together carbon
monoxide is produced:
C + CO2 ! 2CO. What volume of CO can be produced from 3g
carbon at RTP? [4]
5. What volume of oxygen is released when 1000g of sugar,
C6H12O6, is photosynthesis at RTP?
6CO2 + 6H2O ! C6H12O6 + 6O2
[4]
6. Calculate the volume of i) CO2 and ii) H2O produced when 150
cm3 of methane is combusted.
CH4(g) + O2(g) ! CO2(g) + 2H2O(g)
[4]
EC p65 Q1-3
Concentration
The concentration of a solution is defined as:
The units are
or
The formula you must learn is: please label the units you can use use on the
bottom row.
!
Practice Calculations:
1) Calculate the number of moles of potassium hydroxide that must be
dissolved to make the following solutions:
(i)
500cm3 of 1 mol/l
(ii)
200cm3 of 0.5 mol/l
(iii) 100cm3 of 0.1 mol/l
(iv) 2 litres of 0.25 mol/l
(v)
250cm3 of 2 mol/l
2) Calculate the concentration of each of the following solutions of
hydrochloric acid:
(i)
1 mol of HCl dissolved to make 100cm3 of solution
(ii)
2 mol of HCl dissolved to make 1 litre of solution
(iii) 0.1 mol of HCl dissolved to make 500cm3 of solution
(iv) 0.5 mol of HCl dissolved to make 250cm3 of solution
Part 2
1) Calculate the number of grams of substance needed to make each
of the following solutions
(i)
50cm3 of NaOH (aq), concentration 2 mol/l
(ii)
100cm3 of KOH (aq), concentration 0.5 mol/l
(iii) 1 litre of Na2CO3 (aq), concentration 0.1 mol/l
2. Calculate the concentration of each of the following solutions:
(i)
5.65g of NaCl dissolved to make 1 litre of solution
(ii)
2.5g of CaCO3 dissolved to make 100cm3 of solution
(iii) 8g of NaOH dissolved to make 250cm3 of solution
Titrations
We can find out the concentration of a solution of unknown concentration
using another solution of known concentration using this method.
An indicator is used to show the “end point” of the reaction.
Steps to follow…
1. Write the balanced equation for the reaction
2. Calculate the number of moles of substance used of known concentration
3. Use this to find the number of moles of the substance of unknown
concentration and hence the unknown concentration.
EC p71 Q1-2
Empirical Formula Calculation
The empirical formula of a compound is the simplest whole number ratio of
the elements present. The following exampleS illustrates how we calculate
empirical formula in practice:
Eg. 1.2g of Magnesium ribbon is completely burnt to produce 2.0g of
Magnesium Oxide. Calculate the empirical formula.
Magnesium
Oxygen
Mass of element (or percentage mass)
Divide by the Ar
Simplify the ratio
Express the simplest whole number ratio as
the EMPIRICAL FORMULA
NB. The formula of an IONIC COMPOUND is always the same as its
EMPIRICAL FORMULA
Eg2: A hydrocarbon is found to contain 80% carbon and 20% Hydrogen by
mass. Calculate its EMPIRICAL FORMULA…
Carbon
Hydrogen
Mass of element (or percentage mass)
Divide by the Ar
Simplify the ratio
Express the simplest whole number ratio as
the EMPIRICAL FORMULA
Molecular Formula Calculation
In order to calculate the actual (molecular) formula of the compound, we
need its relative molecular mass.
Take the previous questions. The empirical formula was _____________
If we are told that the Mr of the compound is 30, what is the molecular
formula?
EC p69 Q1-3
Empirical and Molecular Formula Questions
1. 6.9g of Sodium combine with 4.8g of Oxygen to form a single
compound of Mr = 78. Calculate the:
(a)empirical formula.
(b)molecular formula.
(3)
2. 9.6g of carbon combine with 2.0g of hydrogen to form a single
compound of Mr = 58. Calculate the:
(a)empirical formula.
(3)
(b)molecular formula.
(2)
3. A compound contains the following elements in the given proportions:
P 43.66%; O 56.40%. The molar mass of the compound is 284g.
Calculate the:
a. empirical formula.
b. molecular formula.
(3)
Experiment: Finding the formula of Magnesium Oxide MgxOy
Aims 4.2 Calculate empirical formulae and molecular formulae
Introduction
Magnesium reacts readily with oxygen. When magnesium ribbon is heated
in air it burns with an intense white light.
1. Write a word equation for the reaction (incl. state symbols)
Magnesium is so reactive that it even reacts with nitrogen in the air,
however the mount of magnesium nitride formed is less than 1 % of the
product of the reaction.
The surface of the magnesium ribbon may be dull due to a layer of
magnesium oxide (or nitride).
The rate of burning of the magnesium depends on the concentration of
oxygen that surrounds it: a good circulation of air is needed for magnesium
to remain burning at a reasonable rate without the reaction stopping.
Hypothesis
Predict the empirical formula of magnesium oxide.
Explain your prediction
Method
2. Plan an experiment to calculate the empirical formula of magnesium
oxide.
Safety Precautions
•
•
•
•
•
Do not look into the crucible when it is heating.
Eye protection is essential.
Open flame will be present.
Do not breathe the fumes generated.
Do not touch the crucible, lid, triangle, ring, or stand during or after
they have been heated.
Results
Record your results
a. Mass of crucible and lid
…………………………………..g
b. Mass of crucible and lid and magnesium strip before burning
…………………………………..g
c. Mass of magnesium
…………………………………..g
d. Mass of crucible and lid and magnesium oxide
…………………………………..g
e. Mass of oxygen
…………………………………..g
Conclusion
Calculate the empirical formula of magnesium oxide MgxOy by working
through the following steps
Magnesium
Mass of element (or percentage mass)
Divide by the Ar
Simplify the ratio
Express the simplest whole number ratio as
the EMPIRICAL FORMULA
Oxygen
Evaluation
7. How do your results compare with the theoretical formula you predicted
in Q2 ?
8. What sources of error could there have been in your experiment?
9. How could you improve on this method ?
Demonstration: The formula of Copper Oxide
Copper oxide can have its oxygen removed by the methane in this
experiment. We say that the copper oxide is “REDUCED”
RESULTS:
Mass of boiling tube = _______________________________
Mass of boiling tube + Copper Oxide = ___________________
Mass of boiling tube + residual copper at the end of the experiment
= ______________________________
Calculate the empirical formula of copper oxide using the following table:
Copper
Mass of element (or percentage mass)
Divide by the Ar
Simplify the ratio
Express the simplest whole number ratio as
the EMPIRICAL FORMULA
Conclusion:
Oxygen
Challenge: VERY HARD STOICHIOMETRY QUESTION
12cm3 of methane gas (CH4) reacted with 20cm3 oxygen to produce
carbon dioxide and water.
(i) Write a balanced equation for the reaction (2)
(ii)How many moles of methane reacted? (1)
(iii) What mass of methane reacted? (1)
(iv) How many moles of carbon dioxide were produced?
(v)What volume of carbon dioxide was produced?
(vi) How many moles of oxygen reacted?
(vii) What volume of oxygen reacted?
(viii) Assuming that the water produced in the reaction condensed to
form a liquid, what volume of gas remained at the end of the
experiment?
Percentage Yield
The percentage yield is a measure of how much of a particular product we
have obtained in an experiment. We can calculate this using the following
formula:
Eg: In an experiment to make Copper(II)Sulphate crystals, a pupil used 4g
of Copper(II)Oxide and XS sulphuric acid. 4.5g of the crystals were
obtained.
Calculate the theoretical yield of the crystals….
Calculate the percentage yield…..
Percentage Purity
This is very important in the world of foodstuffs and medicine. We need to
know how pure a particular substance is. The equation to use is:
% Purity = mass of pure substance X 100%
total mass of sample
Eg. A sample of 22g of ibuprofen was found to contain 18.5g of the pure
substance. Calculate the percentage purity….
EC p67 Q1-3
End of chapter questions
Percentage Purity of Limestone
Objective: to calculate the percentage purity of limestone with respect to
the carbonates that it contains.
Introduction
Limestone is a mixture of mainly calcium carbonate and a little magnesium
carbonate. As with all metal carbonates calcium carbonate reacts with acid
to form a salt, water and carbon dioxide.
1. Write a word equation for the reaction.
2. Write a chemical equation for the reaction (incl. state symbols).
Hypothesis
Predict the % Purity of the sample of Limestone you have been given.
Method
Plan an experiment to calculate the % purity of the sample of the
Limestone you have been given.
Percentage purity = mass of calcium carbonate
mass of limestone sample
x 100%
Safety Precautions
Write down the precautions you will take to minimise the risk of harm
during this experiment.
Results
3. Record your results.
Mass of impure Limestone at the start = __________________
Volume of Carbon Dioxide collected = ____________________
Number of moles of CO2 collected =
Number of moles of Calcium carbonate in the original sample =
Mass of Calcium Carbonate in the original sample =
Percentage purity calculation:
% Purity = mass of pure substance X 100%
total mass of sample
!
Paper 3 Moles Questions