Symbols and Formulae
Consult
Chapter
Chapter
Symbols of elements
There are over one hundred known elements in the periodic table. As research
is ongoing, new elements are still being discovered. When the elements are
first reported they are given a name and symbol that corresponds to the Latin
for their atomic number, for example element 118 was given the name
Ununoctium (Uuo). After it has been firmly established that a new element
has been discovered the element receives a formal name. The most up to date
periodic table can be found on the International Union of Pure and Applied
Chemists (IUPAC). http://www.iupac.org/
Formulae of compounds
The elements of the periodic table are bonded together in different ways to
form millions of different compounds with more being synthesised all the
time.
Covalent compounds achieve stability by sharing electrons.
Ionic compounds are composed of oppositely charged particles (ions) which
are held together by strong electrostatic forces.
In this unit we are interested not so much in how these particles bond
together, but the quantities and proportions in which they do so. We are going
to learn about moles, but not before we have learnt to write the formulae of
compounds.
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Formulae of ionic compounds
Until you learn the formulae and charges of the ions (see pages 2, 3 and 4)
you will make many avoidable mistakes. Get someone test you until you can
answer accurately and with complete confidence.
Metals form positive ions by electron loss (unless they are part of oxyions,
e.g. MnO4-). Non metals of Groups 5, 6 and 7 form negative ions by electron
gain. On the periodic table below, shade the area occupied by metals.
Comment on the balance of metals to non-metals: ............................................
............................................................................................................................
Complete the following table:
Symbol and charge
on ion
Positive ions (cations)
H+
Hydrogen
Group 1
Lithium
Sodium
Potassium
Group 2
Magnesium
Calcium
Strontium
Barium
2
Symbol and charge
on ion
Group 3
Aluminium
Transition metals
Iron
+2
+3
Copper
+1
+2
Zinc
+2
Silver
+1
Hydride
–1
nitride
–3
Oxide
–2
Sulphide
–2
Fluoride
–1
Chloride
–1
Bromide
–1
Iodide
–1
Negative ions
(anions)
Group 5
Group 6
Group 7
3
Other common ions including oxyions
Some chemical formulae contain Roman numerals. These refer to elements
that can exist in more than one valency/has more than one charge possible for
its ions. For example, iron has two common ions: Fe(II) (e.g. FeSO4) and
Fe(III) (e.g. Fe2O3).
High valencies are stabilised in oxyions. Mn7+ does not exist as a free ion, but
does exist in the stable ion MnO4-, manganate(VII).
Complete the following table:
Name of ion
Charge of ion
Formula
1. Ammonium
+1
NH4+
2. Hydroxide
-1
4. Carbonate
-2
5. Hydrogen carbonate
-1
7. Nitrate(V)
-1
8. Phosphate(V)
-3
10. Sulphate(VI)
-2
15. Dichromate(VI)
-2
Cr2O72-
17. Manganate(VII)
-1
MnO4-
4
PO43-
Writing chemical formulae
When a compound is formed,
The sum of all positive charges = The sum of all negative charges
Consider the formula of sodium chloride. What is the formula of:
the sodium ion? ................................ the chloride ion? ................................
Are the number of positive charges and negative charges equal? ....................
Circle the correct formula of sodium chloride:
Na2Cl
NaCl
NaCl2
Consider the formula of magnesium chloride. What is the formula of:
the magnesium ion? .......................... the chloride ion? ................................
Are the number of positive charges and negative charges equal? ....................
How many chloride ions balance the charge on the magnesium ion? ..............
Circle the correct formula of magnesium chloride: Mg2Cl
MgCl
MgCl2
Using brackets
Consider the formula of magnesium nitrate(V). What is the formula of:
the magnesium ion? ........................ the nitrate(V) ion? ................................
So that the number of positive charges equals the number of negative charges,
every magnesium ion needs to attract two nitrate(V) ions. Because the nitrate
ion is composed of more than one element, it is enclosed in brackets to
preserve its identity. The formula of magnesium nitrate(V) is written
Mg(NO3)2.
Consider the formula of ammonium carbonate. What is the formula of:
the ammonium ion? ........................ the carbonate ion? ................................
So the number of positive charges equals the number of negative charges. The
ammonium ion needs to be enclosed in brackets, not the carbonate ion. The
formula of ammonium carbonate is (NH4)2CO3.
5
Write formulae for the following compounds:
Compound
1. Potassium bromide
Formula
KBr
Compound
2. Sodium sulphate(VI)
3. Magnesium chloride
4. Silver nitrate(V)
5. Lead(II) sulphide
6. Calcium nitrate(V)
7. Ammonium hydroxide
8. Aluminium nitrate(V)
9. Aluminium chloride
10. Aluminium oxide
11. Magnesium sulphate(VI)
12. Lithium carbonate
13. Copper(I) oxide
14. Potassium sulphate(IV)
15. Copper(II) oxide
16. Sodium nitrate(III)
Name the following compounds
1. KI
2. BaSO4
3. PbSO4
4. NH4NO3
5. (NH4)2CO3
6. (NH4)3PO4
7. H2SO4
8. PbCO3
9. HNO3
10. Na2Cr2O7
11. HF
12. Zn(NO3)2
13. KHCO3
14. NaH
15. CuCl
16. FeCO3
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Formula
Relative atomic mass (Ar)
The relative atomic mass (Ar) is the average mass of the atoms of that
element compared to the mass of one twelfth of a carbon-12 atom.
Why is carbon-12 used as a standard? Would it not have been easier to use
one atom of hydrogen for comparative purposes? The problem relates to the
nucleus of the common 1H isotope, which, unlike all other atoms, does not
contain any neutrons.
There is a very small difference in mass between a proton and a neutron and
this cannot be ignored at the level of accuracy that scientists now work. The
most abundant carbon isotope 12C has equal numbers of protons and neutrons
– six of each.
•
Relative means the comparative statement must be made in the
definition.
•
Atomic masses have no units - they are relative or comparative values.
•
Relative atomic masses are not always whole numbers.
Chlorine, for example, has a relative atomic mass of 35.5, because chlorine is
made up of two isotopes:
75% chlorine-35 and 25% chlorine-37.
Ar = (75/100 × 35) + (25/100 × 37)
= 35.5
7
Relative formula mass (Mr)
To calculate the relative molecular mass of a compound, we need to know
that:
Relative Formula Mass = Sum of Relative Atomic masses
The following examples have been worked through for you.
Given: Ar H = 1; O = 16; Cl = 35.5; Ca = 40; Fe = 56
Mr Ca(OH)2
= mass of (1 atom Ca) + (2 atoms O) + (2 atoms H)
=
(1 × 40)
+
(2 × 16)
+
(2 × 1)
= 74
Mr Fe(ClO4)3.10H2O
= mass of
(1 × Fe)
+
(3 × Cl)
=
(1 × 56)
+ (3 × 35.5) + (12 × 16) + (10 × 18)
= 534.5
8
+ (12 × O) + (10 × H2O)
Calculate the relative formula masses of the following compounds:
Use a Periodic Table to find values for relative atomic mass (Ar) (do not
confuse Ar with the atomic number.)
Compound
Working
1. N2O4
2. C6H12O6
3. Al(NO3)3
4. CH3COOH
5. (NH4)3PO4
6. CuSO4.5H2O
7. Na2CO3.10H2O
8. (NH4)2SO4.Fe2(SO4)3.24H2O
9
The Avogadro constant
One atom of carbon is twelve times heavier than one atom of hydrogen.
Therefore 12g of carbon-12 contains the same number of atoms as 1g of
hydrogen-1.
It follows that the relative atomic mass in grams of all elements and the
relative molecular mass in grams of all compounds contain the same number
of particles.
This number is the Avogadro constant, which has a value of 6.02 × 1023.
The Avogadro constant is the number of atoms in twelve grams
of carbon-12 (12C).
Calculate the number of atoms (of elements) or molecules (of compounds) in:
Mass
9g C
Working
12g C contain 6.02 × 1023 atoms
∴ 9g C contains 9/12 × (6.02 × 1023) atoms
Ar = 12
=
7g Fe
Ar =
216g Ag
Ar =
3.6g H2O
18g H2O contain 6.02 × 1023 molecules
∴ 3.6g H2O contains:
Mr = 18
=
440g CO2
Mr =
15g CH3COOH
Mr =
135g C6H12O6
Mr =
10
The mole
A mole is the mass of a substance which contains the same number of
particles (atoms, molecules or ions) as there are atoms in 12 grams of
carbon-12 (12C).
How many particles are there in a mole? .........................................................
Moles of solids: elements
One mole of any element is equal to the atomic mass in grams.
mass
Number of moles =
Ar
Complete the following table
Mass → Moles
Mass / g No. moles
112g 56Fe
Moles → Mass
No. moles Mass / g
2 moles 4He
mass
No. Moles =
Ar
=
Ar He = 4
Mass
112
56
= Ar × moles
= 4 × 2
= 8g
= 2
64g 16O
0.1 moles 32Si
3g 12C
0.5 moles 35Cl
24g 32S
0.75 moles 24Mg
1.4g 14N
10 moles 23Na
11
Moles of solids: compounds
One mole of any compound is equal to the relative formula mass in grams.
mass
Number of moles =
Mr
Complete the following table
(Obtain relative atomic masses from a Periodic Table)
Mass → Moles
Mass / g No. moles
1.6g Fe2O3
No. Moles =
Mr = 160
=
Moles → Mass
No. moles Mass / g
mass
Mr
0.75 moles
Fe2O3
1.6
160
Mr =
2.0g CH3COOH
Mr =
8.0g NaOH
Mr =
21.5g Na2CO3
Mr =
= 160 × 0.75
= 120g
= 0.01
3.16g KMnO4
Mass = Mr × moles
0.75 moles
KMnO4
2.5 moles
CH3COOH
0.02 moles
NaOH
0.8 moles
Na2CO3
12