Download An element

Chpt. 4: The Periodic Table
In this Chapter............
• What is an element???
• Chemists and order among the elements
• The Periodic Table
• Periodic Table & arrangement of electrons in
atoms
Robert Boyle from Co. Waterford
• first chemist to define the meaning of the word
element
An element is a chemical substance that cannot be split
up into simpler substances by chemical means.
History of the Elements
1. The Greeks (400 B.C.)
• The Greeks believed that
there were four
elements from which
everything else was
made:
- earth
- air
- water
- fire
2. Robert Boyle (1661)
• Boyle described elements as
“primitive and simple
substances”.
This is close to the modern
definition:
“An element is a substance
that cannot be split up into
simpler substances by
chemical means
• Predicted that compounds
were made from and could be
broken down into elements
3. Humphry Davy (early 1800’s):
• Used electricity to
separate compounds
into their elements
• He discovered
potassium, sodium,
calcium, barium,
strontium and
magnesium
4. Henry Moseley (1914):
• Using X-Rays he
discovered a method of
determining the
number of protons in
the nucleus of an atom.
• The atoms of each
element have different
numbers of protons in
them i.e. no two
elements have the
same number of
protons
No. of protons = Atomic No.
History of the Periodic Table
After 1800’s more and more elements were being
discovered and it was becoming increasingly difficult to
understand and memorise all the properties of each
element.
Four men were responsible for bringing order among
the elements in the form of the Periodic Table:
- Johann Dobereiner
- John Newlands
- Dmitri Mendeleev
- Henry Moseley
• The Periodic table was designed to classify
elements according to their properties and to show
trends in their physical and chemical properties
• The main groups are: I – Alkali metals; II – Alkaline
earth metals; VII – Halogens; VIII (also called group
0) – Noble/inert gases.
• The elements between groups II and III are called
the D-block or transition metals.
1. Dobereiner’s Triads (1829):
• Studied the properties of
various elements in a bid to
find order
• Discovered that when looking
at some groups of three
elements they had similar
chemical properties and the
atomic weight of middle
element was average of the
other two elements
• Group of three similar
elements called a triad
Dobereiner’s Triads
A triad is a group of three elements with similar
chemical properties in which the atomic weight of the
middle element is approximately equal to the average
of the other two.
2. Newlands Octaves (1864):
• John Newland
arranged known
elements in
order of increasing
atomic weights.
• Noticed properties
seemed to repeat every
eight elements.
Newlands Law of Octaves
An octave is a group of elements arranged in order of
increasing atomic weight, in which the first and the
eighth element of each group have similar properties
Problems with Newlands Octaves
• The properties repeat every 8 as noble gases hadn’t
been discovered yet!
• There were several problems such as iron being
grouped with oxygen and sulphur.
• Newlands tried to force all of the known elements
into the table instead of leaving gaps for elements not
yet discovered example: Li and Ag in same group.
3. Mendeleev’s Periodic Table (1869):
Dmitri Mendeleev listed all
known elements (63) in
order of increasing atomic
weight and again noticed
similar properties on every
eight element.
Mendeleev left gaps and
predicted elements that
had not been discovered.
Mendeleev’s Periodic Law
When elements are arranged in order of increasing
atomic weight (relative atomic mass), the properties
of the elements vary periodically.
Mendeleev:
What he did:
• Put elements with the same properties in the same
vertical group.
• Left gaps to make the elements fit into the proper
column (group).
• Predicted that elements (eg. Germanium and
Gallium) would be discovered to fill these gaps.
• Predicted their properties.
• Reversed the order of some elements so that their
properties matched their group e.g. Te and I
Mendeleev's Periodic Table
Mendeleev:
Gained acceptance when:
• new elements discovered fitted properties predicted
• reversals justified by discovery of atomic number
4. Moseley (1913):
• Moseley studied the frequencies of the
x-rays emitted by atoms of different
elements.
• Found frequencies varied depending on
the amount of positive change.
• In other words the difference between
the elements is the number of protons
in the nucleus
• Discovered method for determining
number of protons in nucleus of an
atom – known as atomic number
,
Atomic Number
The atomic number of a atom is the number of protons
in the nucleus of that atom
*Note: Once the atomic number was known it was seen
that Mendeleev’s table was in order of increasing
atomic number.
The Modern Periodic Table
The modern Periodic Table is an arrangement of
elements in order of increasing atomic number
The Periodic Law
When elements are arranged in order of increasing
atomic number, the properties vary periodically
Mendeleev’s
Periodic Table
Modern Periodic
Table
Elements listed in order of
increasing atomic weight
(mass)
Elements listed in order of
increasing atomic number
Noble gasses missing
(undiscovered)
Noble gasses now included as
Group 8 or 0
No atomic numbers listed
(undiscovered)
Atomic numbers now listed,
discovered by Moseley
d-Block (transition) elements
arranged as sub groups
d-Block (transition) elements
now listed in a separate block
Gaps left for undiscovered
elements
All gaps filled
Attention: New Additions to Periodic Table
WOMANIUM (WO)
Physical properties: Generally soft and round in form. Boils at
nothing and may freeze any time. Very bitter if not used well.
Chemical properties: Very active and highly unstable. Possesses
strong affinity with gold, silver, platinum, and precious stones.
Violent when left alone. Turns slightly green when placed next to a
better specimen.
Usage: An extremely good catalyst for dispersion of wealth.
Caution: Highly explosive in inexperienced hands!
MANIUM (XY)
Physical properties: Solid at room temperature but gets
bent out of shape easily. Difficult to find a pure sample.
Due to rust, aging samples are unable to conduct
electricity as easily as young samples.
Chemical properties: Attempts to bond with WO any
chance it can get. Also tends to form strong bonds with
itself. Becomes explosive when mixed with Childrium for
prolonged period of time.
Usage: Possibly good methane source.
Caution: In the absence of WO, this element rapidly
decomposes and begins to smell.
Uses of the Periodic Table
• Obtaining atomic numbers and mass numbers
• Obtaining relative atomic masses
• Writing electronic configurations.
Atomic Number & Atomic Mass Number
• Atomic Number (Z):
- number of protons an element has
- smaller of the two numbers
• Atomic Mass Number (A):
- of an element is the sum of the number of
protons and neutrons in the nucleus of an
atom of that element
- larger of the two numbers
- unit = atomic mass unit a.m.u.
1 a.m.u. = 1.66 x 10-24g
- always a whole number
Nuclear Formula of an element
Using the Nuclear Formula – Sample Questions
1. How many protons, neutrons and electrons are in
the following:
i) 3517 Cl
ii) 4521Sc3+
Mass Spectrometry
1919 an instrument called mass spectrometer was
built to measure the mass of atoms.
Mass Spectrometer
Isotopes
Isotopes are atoms of the same element (i.e. they
have the same atomic number) that have different
mass numbers due to different numbers of neutrons
in the nucleus.
Since the neutron has no charge, the quantity of
neutrons in an atom can change slightly without
having an effect on the atom.
Example: Neon – 2 isotopes, Chlorine – 2 isotopes,
Carbon – 3 isotopes and Hydrogen – 3 isotopes
3 Isotopes of Hydrogen
3 Isotopes of Carbon
Calculating the average mass of an atom:
A sample of chlorine is found to consist of 75% 3517Cl and
25% 3717Cl . Calculate the average mass of an atom of
chlorine.
Average Mass = (Mass isotope 1 x % abu) + (Mass isotope 2 x % abu)
100
Relative Atomic Mass (Atomic Weight) Ar
The relative atomic mass (Ar) of an element:
- is the average of the mass numbers of the
isotopes of the element,
- as they occur naturally,
- taking their abundances into account and,
- expressed on a scale in which the atoms of the
carbon-12 isotope have a mass of exactly 12
units.
(average mass of an atom, measured relative to the
mass of the carbon -12 isotope)
• The carbon-12 isotope has a perfect mass of 12amu.
• 1/12 of this mass is a perfect 1.000
• All average masses represented on the periodic table
are compared to this value of 1.000. We therefore say
that these atomic masses are relative to the mass of
1/12th of the C-12 isotope.
Ratio:
Ar=
mass of atom of element
1/12 mass of atom of carbon-12
Note: ratio therefore no units
Relative Atomic Mass Calculations
Formula:
Ar= (Mass isotope 1 x % abu) + (Mass isotope 2 x % abu)
100
Example 1:
An element, X, consists of 92.2% atoms with a mass
28, 4.7% of atoms with a mass 29 and 3.1% of atoms
with a mass 30. What is the relative atomic mass?
What is the element?
Example 2:
The two isotopes of chlorine have mass numbers of 35
and 37 respectively. Taking the relative atomic mass of
chlorine to be 35.46, calculate the % of each isotope
present in the element
Writing Electronic Configurations
• is the arrangement of electrons in an atom
• two methods:
- Bohr Model – in terms of energy levels (simple)
- Energy sublevels
Bohr Model Method:
• Find atomic number of element.
• As all elements are neutral - number of protons equals
number of electrons.
• Remember:
n = 1 energy level holds 2 electrons
n = 2 energy level holds 8 electrons
n = 3 energy level holds 8 electrons
n = 4 energy level holds 18 electrons
Bohr Model Example:
Write the electronic configuration for potassium showing
the number of electrons in each main energy level.
From periodic table potassium has atomic number 19
therefore since neutral atom:
no. of protons = no. of electrons = 19
1st energy level = 2 e2nd energy level = 8 e3rd energy level = 8 e4th energy level = 1 eSo, electronic configuration of potassium is (2,8,8,1)
Try the following:
Write the electronic configuration for the following
showing the number of electrons in each main energy
level:
i) Fluorine
ii) Calcium
Note:
- you must be able to write the electronic
configurations of the first 20 elements in terms of
the number of electrons in each main energy
level.
- number of electrons in outer shell is same as
group number e.g. Lithium group I has one outer
electron, Boron group III has three outer electrons
Element Atomi Electrons in
Electric
c No. each Shell
Config.
n=1
n=2 n=3 n=4
H
1
1
1
He
2
2
2
Li
3
2
1
2,1
Be
B
C
4
5
6
2
2
2
2
3
4
2,2
2,3
2,4
N
7
2
5
2,5
O
8
2
6
2,6
F
9
2
7
2,7
Element Atomi Electrons in
Electric
c No. each Shell
Config.
n=1
n=2 n=3 n=4
Ne
10
2
8
2,8
Na
11
2
8
1
2,8,1
Mg
12
2
8
2
2,8,1
Ca
20
2
8
8
2
2,8,8,2
Electron Configuration - in terms of sublevels
Remember:
Main energy
level
Sublevel
Electronic Configuration – in terms of sublevels
An energy sublevel is a group of orbitals, within an
atom, all having the same energy.
One orbital in an ‘s’ energy sublevel
Three orbitals in a ‘p’ energy sublevel - X, Y, Z
Five orbitals in a ‘d’ energy sublevel
In terms of energy s<p<d<f.
Electrons will fill sublevels in order of increasing energy.
S sublevel (orbital) can hold 2 electrons
P sublevel can hold 6 electrons, 2 in each px, py and pz
orbital
d sublevel can hold 10 electrons, 2 in each of the 5d
orbitals
Main Energy Level – 1, 2, 3, 4 etc.
Sublevel – s, p, d, f
Orbitals – 1s, 2s, 2px etc.
Electrons – 2 in each orbital
Aufbau Principle
Aufbau Principle states that when building up the
electronic configuration of an atom in its ground
state , the electrons occupy the lowest available
energy level.
• 1s orbital must be filled before the 2s, 2s must
be filled before 2p, etc….
• Electronic Configuration Order:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f
(4s is lower in energy than 3d)
Writing Electronic Configurations
No. of electrons
5
2p
Main energy level
Type of sublevel
Example 1:
Write the electronic configuration (s, p) for an atom of
oxygen
Solution:
• Determine number of electrons in oxygen
• Place the electrons into sublevels according to Aufbau
Principle and using periodic table
Remember – s sublevel 2 electrons
- p sublevel 6 electrons
- d sublevel 10 electrons
Try:
Write the electronic configuration for the elements from
1 to 24.
Remember:
The ‘4s’ energy sublevel is slightly lower in energy than
the ‘3d’ energy sublevel. Therefore, its orbital must be
filled with electrons before the orbitals in the 3d
sublevel.
Exceptions to the rule:
If p, d and f shells are exactly half filled or completely
filled they will have extra stability:
24) Cr = 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5
29) Cu = 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d10
You must be able to write the
electronic configurations of the first
36 elements in terms of their s, p, d
configurations!!!
Electronic Configuration of Ions
Ion – atom that has lost or gained electrons (has a
charge on it)
- loses electrons = positively charged ion
- gains electrons = negatively charged ion
Example:
What is electronic configuration of Al3+ ion? What
neutral atom has the same configuration?
Try:
Write the electronic configuration (s, p) of the S2- ion.
What neutral atom has the same configuration?
Example 2:
Identify the species represented by [ 1s2, 2s2, 2p6]2+
Try:
Identify the species represented by:
[1s2, 2s2, 2p6, 3s2, 3p6]2-
Arrangement of electrons in orbitals of Equal
Energy
The electronic configuration of nitrogen is:
1s2, 2s2, 2p3
The question is how are the three electrons in the p
sublevel distributed among the three p orbitals – x, y, z?
Answer:
Hund’s Rule of Maximum Multiplicity
(Bus Seat Rule)
Hund’s Rule of Maximum Multiplicity states that when
two or more orbitals of equal energy are available, the
electrons occupy them singly before filling them in pairs
Electronic Configuration of Nitrogen using
Hund’s Rule:
Before
1s2, 2s2, 2p3
After:
1s2, 2s2, 2px1, 2py1, 2pz1
2p sublevel – 1e- will enter px orbital
- 1e- will enter py orbital
- 1e- will enter pz orbital
Using ‘Arrows-in-Boxes’ Diagrams:
Distribution of electrons in nitrogen p sublevel:
px py pz
• Box = orbital
Arrow = direction of spin of electron
• Arrows represent spinning of electrons on own axis
as they revolve around nucleus
• Electrons can spin in either a clockwise direction or
an anti-clockwise direction
Pauli Exclusion Principle
The Pauli Exclusion Principle states that no more than
two electrons may occupy an orbital and they must
have opposite spin.
Note:
= electron spinning clockwise
= electron spinning anti-clockwise
Complete ‘arrows-in-boxes’ diagrams and
the extended form of electronic
configurations for the first ten elements
Note
The ‘arrows-in-boxes’ diagrams or extended
forms of electronic configurations need only
be given if you are specifically asked about
the distribution of electrons in the p
sublevel.