V.C.E Chemistry
Year 11 Revision Notes
Atomic Theory
List the three fundamental particles and their
Properties.
Draw a diagram to show electronic arrangement around the
nucleus of an atom.
Given the nuclear structure of an atom approximate its
mass in Atomic Mass Units..
Define the Relative Atomic mass of an element.
State properties of an element given its position on the
periodic table.
Boiling Points and Melting Points
(B.P. and M.P.)
Describe the outer electronic structure of the atoms in a
molecule.
which has a negative charge of 1.602 X 10–19 coulomb. The mass of an electron is 9.102 X 10–31 kg – which is negligible compared to the mass of a proton or neutron. The number of electrons revolving around the nucleus is equal to the number of protons in the nucleus.
which has a positive charge of 1.602 X 10–19 coulomb – the same magnitude as that of the electron. A proton is obtained by removing the electron from an atom of Hydrogen, leaving H+. A proton has a mass of 1.6725 X 10–27 kg, which is much greater than that of an electron but slightly less than that of a neutron
which is electrically neutral. The proton and neutron make up the nucleus of atoms. The neutron has a mass of 1.6748 X 10–27 kg
The atom is mostly empty space. The nucleus has a radius of 10–15 m whereas the atom is 10–10 m. The nucleus is about 99.97% of the weight of the atom
An
element is composed of atoms all of which have the same atomic number.
The
Atomic Number of an atom is the number of protons in the nucleus of the atom.
Usually the number of protons on the nucleus is equal to the number of
electrons.
The
Mass Number of an atom is the total number of protons and neutrons in the
nucleus.
Both
the atomic number and mass number are shown with the relevant chemical symbol.
Example
|
Mass Number |
40 |
|
|
|
|
Ca |
|
Atomic Number |
20 |
|
ie
20 Protons
40 Protons and Neutrons
20 Protons therefore 20 Neutrons
(40 – 20)
20
Electrons (equals number of Protons)
Investigation
of the spectra of many different elements shows that the electrons in their
elements are rest5ricted to certain permitted energy levels. The results of a
great deal of experimental work can be summarized as flows.
i)
A
number of electrons can have very close levels of energy. Such a set of
electrons is called an electron shell. From the nucleus outwards the shells are
either numbered or called K, L, M, N, O, O etc shells.
ii)
The
maximum number of electrons in each shell is limited
ie
2 electrons in the first of K shell
8 electrons in the second of L shell
18 electrons in the third of M shell
32 electrons in the fourth of N shell
2n2 electrons in the Nth shell
iii) The outer shell of electrons in an
atom never holds more than 8 electrons regardless of the maximum number
electrons for the shell indicated in (ii)
above. For example gthe third shell can hold up to 18 electrons but will hold no
more than 8 electrons unless there are electrons in the fourth shell, as it is
then not an outer shell
If
we make an approximation and say the
mass of a proton is nearly equal to the mass of a neutron and then say that the
mass of a proton or neutron is equal to 1, we have what is called an
Atomic Mass Unit.
Thus
the mass number of an atom is a measure of the mass of an atom relative to the
mass of a proton or neutron having a mass of 1 Atomic Mass Unit.
So
|
12 |
|
|
|
C |
|
6 |
|
Would
have a mass of 12 Atomic Mass Units
|
16 |
|
|
|
O |
|
8 |
|
Would
have a mass of 16 Atomic Mass Units
Isotopes
are atoms of the same element with the same atomic number but with
different mass numbers. A group of atoms with the same atomic number,
even oif some have
a different mass number are of the same element as they have the same
number of protons.
These
elements with a different mass number must have a different number of neutrons.
Thewy are all ISOTOPES of that element.
An
isotope is identified by writing the mass number after the name
|
|
12 |
|
|
Carbon 12 |
|
C |
|
|
6 |
|
|
|
235 |
|
|
Uranium 235 |
|
U |
|
|
92 |
|
There are three isotopes of hydrogen all of which have
special names
Hydrogen 1 Protium
(99.985% of all Hydrogen atoms)
Hydrogen 2 Deuterium (0.015% of all Hydrogen atoms)
Hydrogen 3 Tritium (Trace only)
The Hydrogen 1 atom has a Relative Atomic Mass of 1.0078
Natural Hydrogen has a Relative Atomic Mass of 1.008
Carbon occurs as a mixture of three isotopes
Carbon 12 (99,892%
of all Carbon atoms)
Carbon 13 (1.108% of all Carbon atoms)
Carbon 14 (<0.001% of all Carbon atoms)g objects, it
slowly dexays, hence the life of the object can be determined
Note Carbon
14 is used in the process called “carbon dating”. It is manufactured daily
in the upper atmosphere, but when absorbed with carbon in living objects it
slowly decay, and a result the age of the object can be calculated
Natural Carbon has a Relative Atomic Mass of 12.0111
Carbon 12 has, by definition a Relative Atomic Mass of
12.000
Carbon 13 has a Relative Atomic Mass of 13.0034
Calcium occurs as a mixture of six isotopes
Calcium 40 (96.97% of all Calcium atoms)
Calcium 42 ( 0.64%
of all Calcium atoms)
Calcium 43 ( 0.145%
of all Calcium atoms)
Calcium 44 ( 2.06%
of all Calcium atoms)
Calcium 46 ( 0.0033%
of all Calcium atoms)
Calcium 48 ( 0.2%
of all Calcium atoms)
As nearly all elements have at least two
isotopes, and as the isotope(s) with extra neutrons must weigh more, this will
affect the average mass of an element.
The Relative Atomic Mass of an element is
the weighted mean of all the naturally occurring isotopes of that element
Carbon 12
Isotopic Mass = 12.00
% Abundance = 98.893%
Carbon 13
Isotopic Mass = 13.0034 %
Abundance = 1.107%
Relative Atomic Mass is given by
Carbon 12 is taken as the standard of
relative atomic mass. All other elements’ relative atomic mass is related to
Carbon 12 as 12 precisely (12.00 and as many zeroes required). An atomic mass
spectrometer is used to measure the isotopic mass. The spectrometer is
calibrated with the Carbon 12 isotope, and all other isotopes are relative to
this
The
reason there is a difference between the atomic mass of an element and the mass
number of its most common isotope can be readily seen from the exercise above.
As we know, each different isotope of an element has a different atomic weight.
Thus, to achieve a relative atomic mass (that includes all is0topes) we must
take the weighted mean of all isotopes. As this relative atomic mass is the
weighted mean of all the isotopes, it follows that the relative atomic mass will
differ slightly from the mass number of that element’s most common isotope
e.g.
C = 12.011 Carbon 12 is
the most common isotope
H = 1.008 Hydrogen 1 is the most common isotope
In
a Mass Spectrometer, ions are accelerated by an electric fields caused by two
charged plates, and deflected by a magnetic field along part of a circular path
in a vacuum.
The
path of the ions is registered on a photographic plate, or on an electron
detector, and form the record or spectrogram. From these, the relative masses
and their properties can be estimated (lighter particles have a smaller radius
or curve)
NOTE
A positive ion is produced when an electron from a heated filament colides with
a neutral particle and breaks off an electron leaving a positively charged
particle.
The
Mass Spectrograph can be used to obtain
1)
The
atomic masses of the element
2)
The
relative abundance of different isotopes
3)
The
mass of a molecule
4)
Clues
about the structure of complex molecules
To
determine the Relative Atomic Mass of an element when you have been given the
mass spectrograph
i)
measure
the peak height of each isotopic peak
ii)
calculate
Peak Height X Isotopic Mass of each isotope
iii)
calculate
sum of all (ii) / 100
d)
When
measuring quantities of matter such as mass and volume, the units gram and cubic
metre are used. However, chemists have defined a different and distinct physical
quantity called Amount of Substance. The unit chosen for this quantity is the
mole
A
mole is the amount of substance that contains as many elementary entities as
there are atoms in 0.012 kg (12 g) of Carbon 12.
Ie
12 g of Carbon 12 = 1 mole of Carbon 12
The
number of elementary entities contained in an amount of substance equal to 1
mole is the same for all substances. This quantity is found to be 6.0225 X 1023
mol-1, and is known as the Avogadro Constant.
12
g of Carbon 12 contains 6.0225 X 1023
atoms
32
g of Oxygen gas (O2) contains 6.0225 X 1023 molecules
58.5
g of Sodium Chloride contains 6.0225 X 1023 Na+ ions and
6.0225 X 1023 Cl- ions
Number
of Mole =
In
the periodic table the elements are classified in order of increasing atomic
order, but are also arranged in vertical groupings according to the number of
electrons in the outer shell. Thus the horizontal lines have elements with atoms
having the same number of shells while the vertical columns contain elements
with the same number of outer shell
electrons. As nearly all physical and chemical properties depend on the the
electronic configuration, properties vary periodically as you progress through
the table. The physical and chemical properties show a trend down a group.
Although each group has elements with the same number of outer shell electrons,
their properties are affected because when going down a group there is one extra
shell of electrons causing the outer shell to be further from the nucleus.
First
Four Periods
All
of Groups I, II and III are metals (except Boron)
All
of Group VIII are Inert (or Noble) Gases (their outer shells are full)
All
of Groups IV, V, VI and VII and Boron are non-metals (Boron has a few metallic
characteristics)
Group
I metals are more reactive than Group II metals which are more reactive than
Group III, and in each group, activity increases down the group.
Group
VII has the most reactive non-metals.
Elements
in groups 1,2, 3 and 4 are bonded in infinite arrays and have high M.P.’s and
B.P.’s. The M.P.’s and B.P.’s of groups 1, 2 ,3 and 4 increase from left to right and from top to bottom down
the group.
For
elements in Groups 5.6 and 7, M.P.’s and B.P.’s decrease from left to right
and going down the group.
For
each group, density increases going down the group.
Electrical
conductivity decreases as the number of electrons in the outer shell increases.
Less electrons present in the outer shell means that those electrons are more
easily lost, resulting in the electrons being able to move more freely and
consequently have higher electrical conductivity.
The
metals, (Groups 1, 2 and 3) are good conductors of electricity as the outer
shell electrons are free to move. The electrical conductivity increases down to
group because the outer shell is further from the nucleus.
The
non-metals (Groups 5, 6 and 7) have
very poor electrical electrical conductivity because the outer shell electrons
are not as free to move.
The
Group IV elements are semi-conductors and need the presence of a Group 3 or 5
element as an impurity before they conduct electricity.
The
metals (groups 1,2 and 3) are all electron donors. The order of ability to
donate electrons is
Group
1 > Group 2 > Group 3
The
groups 5. 6 and 7 are all electron acceptors. The order of ability to accept
electrons is
Group
7 > Group 6 > Group 5
The
group 4 elements can act as both electron acceptors and donators
The
atoms in a molecule are joined by either ionic bonds or covalent bonds.
A
covalent bond is formed by the sharing of outer shell electrons between the
atoms.
An
ionic bond is formed by the complete transfer of electrons from one atom to
another
Ionic
solids consist of positive and negative ions held together by the attraction
between their opposite electrical charges
Network
solids, like diamond, have all the atoms held together by covalent bonds.
Pure
metals are a different case altogether. Metals consist of positive ions held
together by a sea of electrons.
is
a measure of the electron attracting power of the element. Electronegativity
increases from left to right across the periodic table due to the increasing
number of electrons in the outer shell. Electronegativity decreases when going
down a group because the outer shell is getting further away from the nucleus.
Group
I Very low electronegativity –
form positive ions eg K+,
Li+, Na+
Group
II Fairly low electronegativity – form positive ions eg Ba2+, Ca2+,
Mg2+
GroupIII
Low electronegativity – form positive ions eg B3+ , Al3+
, Ga3+
Group
IV Medium electronegativity – less likely to form either positive or negative
ions
Group
V High electronegativity – forms negative ions eg N3-, P3-
Group
VI Fairly high electronegativity - forms negative ions eg O2-,
Group
VII Very high electronegativity – form negative ions eg Cl-, F-, I-
The
relative positions of elements in the periodic table determines the degree of
purity of the type bond (covalent or ionic) formed between the two elements. If
the elements are close together in the periodic table then a covalent bond will
result.
Eg
H2 BrCl, Na2
Conversly
if the elements are well separated then the bond type will be ionic, eg NaBr,
BeF2, NaCl. The reasons for this as follows
Covalent
bonding involves the sharing of outer shell electrons, usually in pairs, to make
up the complete outer shell of electrons in the atom. Each shared pair of
electrons constitutes on covalent bond. In the case where each atom is two
electrons short in the outer shell, two electrons from each atom are shared.
This is called a double covalent bond. Covalent bonds are generally formed
between non metals, but some metals form covalent bonds in the vapour state, eg
Na2
If
the covalent bond is between two identical atoms, the bond is one in which the
electrons are equally shared, This is called a pure covalent bond.
Other
covalent bonds have a slightly positive charge on one atom and a slightly
negative charge on the other. This is because of differences in the
electronegativities, eg H – Cl.
Chlorine has a higher electronegativity than Hydrogen, so the Cl atom has a
slightly negative charge and the H atom has a slightly positive charge. As a
result, H-Cl bond has some ionic character. The same is true for carbon monoxide
(CO). Since the oxygen attracts electrons more than the carbon does, the O is
slightly negative while the carbon is slightly positive.
The
same is true when one atom is larger than the other, with the larger atom
attracting the shared electrons more than the smaller atom. Thus one end of the
bond is slightly positive and the other end is slightly negative. Such covalent
bonds are said to be polar.
Ionic
bonds are formed by the transfer of outer shell electrons to the outer shell of
a non metal. The result is that each atom (or atoms) now has an outer shell of 8
electrons. The metal is left with a positive charge (called a positive ion or
cation) and the non metal is left with negative charge (called a negative ion or
anion).
The
charges on the ions are called electro valencies eg Na+, Mg2+,
O2–, Cl–. An electrostatic force of attraction exists
between the oppositely charged metal and non metal. The conventional way of
writing bonding is
Ionic
bonding may be represented in two other ways
a)
symbols
and electronic configuration
Na +
Cl –
Na
+ +
Cl –
2,8,1
2,8,7
2,8 2,8,8
b)
symbols
and valency electrons
Some ionic bonds are not purely ionic.
With metal to non-metal bonding, although the metal transfers its outer shell
electron(s) to the non metal (eg NaCl) the fact that the metal is now positively
charged means that it pulls the electron back a little so that the electron
spends part of its time in the Na’s orbit. Thus salts are slightly covalently
bonded in nature.