Fundamentals of Inorganic Chemistry
Lectures 1 & 2
SCH 100
Dr. E. Changamu
Office: Biochemistry Rm 11
Evaluation
• 2 CAT 20 marks• Practical 10 marks (Rem. Lab coats)• End of term Exam 70
Syllabus• The early theories of atomic structure: the fundamental
particles of the atom; • Rutherford’s planetary model• The Bohr theory of the atom; failure of the Bohr theory. • Planck’s quantization of energy • The photoelectric effect; • Qualitative treatment of the atomic orbitals (s, p d and f).
• The Aufbau Principle and the periodic table. • Common oxidation states of the elements.
Syllabus• Naturally occurring and artificially made
isotopes, and their applications. • Atomic properties of the elements
– electronegativity – electron affinity– ionization energy
• Chemical bonding – ionic, – covalent, – metallic, – coordinate
Syllabus• Hybridization of atomic orbitals and shapes of simple
molecules and ions.
• The nature of ionic and covalent compounds as influenced by the above factors.
• The mole concept and its application. • General concepts of acids and bases - strong and weak acids
and bases; pH calculations. • Balancing of redox reactions.
Inorganic Chemistry
• Is the custodian of all the elements known.• Is concerned with the
– Occurrence of the elements in nature– Extraction from natural sources– The reactions of the elements with other
elements and compounds– Safety and application of the elements and/ or
their compounds
Inorganic Chemistry• Is related to the other divisions of chemistry
– Organometallic chemistry bridges inorganic and organic chemistry. It deals with compounds containing direct metal-carbon bonds.
The study of the composition, properties, and transformations of matter
Chemistry
Inorganic Chemistry
– Bioinorganic chemistry bridges biochemistry and inorganic chemistry
– Environmental chemistry includes the study of both inorganic and organic compounds.
Lecture Objectives At the end of this lecture you should be able to
– describe the historical development of atomic structure
– descried the nature of electrons, protons and neutrons
– explain the Thomson model of the atom– explain Rutherford model of the atom– describe the nature of electromagnetic radiation– discuss the Bohr model and the atomic hydrogen
spectrum
Atomic Structure
460 BC Democritus develops the idea of atoms
He pounded up materials in his pestle and
mortar until he had reduced them to smaller
and smaller particles which he called Atoma
ATOMA
(greek for indivisible)
HISTORICAL DEVELOPMENT OF THE ATOMIC THEORY
1808 John Dalton1
Suggested that all matter was made up of tiny
spheres that were able to bounce around with
perfect elasticity and called them Atoms.
John Dalton, A New System of Chemical Philosophy, 1808; reprinted with an introduction by Alexander Joseph, Peter Owen Limited, London, 1965.
HISTORICAL DEVELOPMENT OF THE ATOMIC THEORY
Dalton’s Atomic Theory • Elements are made of extremely small particles called atoms.• Atoms of a given element are identical in size, mass, and other
properties; atoms of different elements differ in size, mass, and other properties.
• Atoms cannot be subdivided, created, or destroyed.• Atoms of different elements combine in simple whole-number
ratios to form chemical compounds.• In chemical reactions, atoms are combined, separated, or
rearranged.
HISTORICAL DEVELOPMENT OF THE ATOMIC THEORY
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1898 Joseph John Thompson
Found from his cathode ray tube experiments
that atoms could sometimes eject a far smaller
negative particle which he called an electron.
The electrons came from the atoms making up the cathode.
HISTORICAL DEVELOPMENT OF THE ATOMIC THEORY
Joseph John Thompson
• Thomson found out that the electrons were the same regardless of
the metal he used for the cathode. Therefore he concluded they
were part of the structure of all atoms.
• He calculated the ratio of their charge to their mass
Charge/mass = 1.76 x 1011 C/kg
• Robert Millikan determined the charge of the electron
Charge of electron, e = 1.602 x 10-19 C
• Thomson calculated the mass of the electron to be
Mass of electron, me = 9.109 x 10-31 kg
HISTORICAL DEVELOPMENT OF THE ATOMIC THEORY
Thompson develops the idea that an atom was made up of electrons scattered unevenly within an elastic sphere surrounded by a soup of positive charge to balance the electron's charge like plums surrounded by pudding.
PLUM PUDDING
MODEL
HISTORICAL DEVELOPMENT OF THE ATOMIC THEORY
Ernest Rutherford, 1910
• Oversaw Geiger and Marsden carrying out his
famous experiment.
• They fired α-particles (a type of naturally-occurring
radiation consisting of positively charged helium
atoms) at a piece of gold foil which was only a few
atoms thick.
• They found that although most of them passed
through. About 1 in 10,000 hit something and
bounced right back.
HISTORICAL DEVELOPMENT OF THE ATOMIC THEORY
gold foil
helium nuclei
They found that while most of the α-particles passed through the foil, a small number were deflected and, to their surprise, some helium nuclei bounced straight back.
helium nuclei
HISTORICAL DEVELOPMENT OF THE ATOMIC THEORY
The Proton
• In 1918, while Rutherford was performing his various experiments in the field of radioactivity, bombarding nitrogen gas with alpha particles, he noticed that one of the experimental results was a surge of hydrogen.
• He correctly deduced that the hydrogen atoms must have come from within the nitrogen atoms themselves, which would mean that there was something within all of these atoms which was divisible, the amount of which would determine what element the atom represented.
The proton• The particle Rutherford isolated was the proton,
which by itself constitutes the nucleus of a single hydrogen atom, though in this case it was “ionized” (missing its electron, thereby giving it a net positive charge), which Rutherford determined by exposing the resulting hydrogen to magnetic fields.
• Rutherford’s new evidence
allowed him to propose a more
detailed model with a central
nucleus.
• He suggested that the positive
charge was all in a central
nucleus. With this holding the
electrons in place by electrical
attraction.
HISTORICAL DEVELOPMENT OF THE ATOMIC THEORY
The Neutron
• By James Chadwick in 1932.• Last to be discovered due to its neutrality.• Now the structure was complete
(somewhat)• See details of discovery in
http://isaacmmcphee.suite101.com/the-discovery-of-the-neutron-a46060
Particle Charge (e) Mass (amu)Proton (p) +1 1Neutron (n) Neutral 1Electron (e-) -1 1/1836
Fundamental particles
1 atomic mass unit (amu) = 1.672621777(74)×10−27 kg
1 electric charge (e) = 1.602176565(35)×10−19 C
Nature of Electromagnetic radiation
Niels Bohr, 1913
• Studied under Rutherford at the Victoria
University in Manchester.
• Bohr refined Rutherford's idea by adding
that the electrons were in orbits. Rather like
planets orbiting the sun. With each orbit only
able to contain a set number of electrons.
HISTORICAL DEVELOPMENT OF THE ATOMIC THEORY
Bohr’s Model
Nucleus
Electron
Orbit
Electron orbits26
HELIUM ATOM
+N
N
+-
-
proton
electron
neutron
Shell
What do these particles consist of?
ATOMIC STRUCTURE
the number of protons in an atom
the number of protons and neutrons in an atom
He2
4 Atomic mass = protons + neutrons
Atomic number = No of protons
number of electrons = number of protons
ATOMIC STRUCTURE
Electrons are arranged in Energy Levels
or Shells around the nucleus of an atom.
• first shell a maximum of 2 electrons
• second shell a maximum of 8
electrons
• third shell a maximum of 8
electrons
ATOMIC STRUCTURE
There are two ways to represent the atomic
structure of an element or compound;
1. Electronic Configuration
2. Dot & Cross Diagrams
ELECTRONIC CONFIGURATION
With electronic configuration elements are
represented numerically by the number of electrons
in their shells and number of shells. For example;
N
Nitrogen
7
14
2 in 1st shell
5 in 2nd shell
configuration = 2 , 5
2 + 5 = 7
ELECTRONIC CONFIGURATION
Write the electronic configuration for the following elements;
Ca O
Cl Si
Na20
40
11
23
8
17
16
35
14
28B
11
5
a) b) c)
d) e) f)
2,8,8,2 2,8,1
2,8,7 2,8,4 2,3
2,6
DOT & CROSS DIAGRAMS
With Dot & Cross diagrams elements and
compounds are represented by Dots or Crosses to
show electrons, and circles to show the shells. For
example;
Nitrogen N XX X
X
XX
X
N7
14
DOT & CROSS DIAGRAMS
Draw the Dot & Cross diagrams for the following elements;
O Cl8 17
16 35a) b)
O
X
XX
X
X
X
X
X
Cl
X
X
X
X X
X
XX
X
X
X
X
X
XX
X
X
X
SUMMARY
1. The Atomic Number of an atom = number of
protons in the nucleus.
2. The Atomic Mass of an atom = number of
Protons + Neutrons in the nucleus.
3. The number of Protons = Number of Electrons.
4. Electrons orbit the nucleus in shells.
5. Each shell can only carry a set number of electrons.
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