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CHAPTER 3
ATOMIC AND ELECTRONIC
STRUCTURE
Name: Prachayanee Chueamsuwanna
date: October 6,2015
ATOMIC THEORY OF MATTER
DEFINING AN ATOM

An atom is the smallest part of an element that retains its identity
in a chemical reaction.
STRUCTURE OF THE ATOM: DISCOVERY
OF SUBATOMIC PARTICLES

ELECTRONS : particles that have a negatively charged, and was
discover by J.J. Thomson in 1890

PROTONS : particles that have a positively charged, and has a
mass of 1.67262 x 10-24 g. and was discover by Rutherford in 1919.

NEUTRON : has neutral charge which has a mass slightly bigger
than the proton, but without charge 1.67493 x 10-24 g.
MODEL OF ATOMIC STRUCTURE
1.DALTON’S MODEL

John Dalton proposed that all matter is composed of very small
things which he called atoms. This was not completely new
concept as the ancient Greeks had proposed that all matter is
composed of small, indivisible objects. When Dalton proposed his
model electrons and the nucleus were unknown.
2. THOMSON’S MODEL
THE PLUM PUDDING MODEL

Initially it was supposed that electrons and positively charged
particles were evenly distributed in a spherical atom. This
prevailing theory was that of the “plum pudding” model, put
forward by Thomson.
3. RUTHERFORD’S ATOMIC MODEL :
RADIOACTIVITY

The discovery of radioactivity lead to discovery protons as well as
the atomic model that the atoms consisted of the nucleus and
which housed protons and that electrons surrounded the
nucleus.

Radioactivity is the spontaneous emission of radiation by an
atom.
4. BOHR’S MODEL

A new model was proposed, by Niels Bohr(1885-1962) in 1913,
which says that an electron is found only in specific circular
paths, or orbits around the nucleus.

To move from one energy level to the other, an electron must
gain or lose just right amount of energy. A quantum of energy is
the amount of energy needed to move an electron from one
energy level to another energy level.
5. ERWIN SCHORDINGER : QUANTUM
MECHANICAL MODEL

The quantum mechanical determined the allowed energies and
electron can have and how likely it is to find the electron in
various locations around the nucleus of an atom.

The Bohr was limited. Erwin Schodinger in 1926 developed a
mathematical treatment into which both the wave and particle
nature of matter could be incorporated. This is known as
quantum mechanics.

Solving the wave equation gives a set of wave functions, or
orbitals and their corresponding energies. Each orbitals describes
a spatial distribution of electron density. An orbitals is described
by a set of three quantum numbers.
ATOMIC ORBITALS

Solutions to Schordinger’s equations give the energies, or energy
levels, an electron can have. For each energy level, the
Schordinger’s equation also leads to a mathematical expression
called an atomic orbital which describes the probability of
finding an electron at various locations around the nucleus of. An
atomic orbitals is represented pictorially as a region of space in
which there is a high probably of finding an electron.

Each sublevel corresponds to one or more orbitals of different
shapes. The orbitals describe where an electron is likely to be
found.
S ORBITALS

Are sphere shaped and are concentric and bigger as the main
shell number increases.

The value of l for S orbitals is 0.

They are spherical in shape.

The radius of the sphere increases with the value of n.
P ORBITALS

Two dumbbells on opposite sides of the nucleus, orientated along
respective three axis of 3 space. Found in the principal shell 2
and up, not in 1.

The value of l for P orbitals is 1.

They have two lobes with a node between them.
D ORBITALS AND OTHER HIGHER
ENERGY ORBITALS

Are 4 lobed(except one). These occur from the principal shell 3
and up. F orbitals are more complex( 8 lobed) and g…

The value of l for a d orbitals is 2.

Four of the five d orbitals have 4 lobes; the other resembles a p
orbital with a doughnut around the center.
ELECTRONIC CONFIGURATIONS

There are three rules that tells the position of an electron namely,
the Aufbau Principle, the Pauli Exclusion Principle and the Hand’s
Rule.
THE AUFBAU PRINCIPLE (BUILDING UP
PRINCIPLE)

According to principle, electrons occupy the orbitals of lowest
energy first. It dictates that for every further proton in the nucleus,
there is an electron in an orbital of that atom. This principle also
dictates the chemical and physical properties of an element,
and its position in the periodic table.
PAULI EXCLUSION PRINCIPLE

No two electrons in the same atom may have the same series of
quantum numbers. An atomic orbital describes at most two
electrons.

Electron Configurations
-
The electron configuration is a way of how electrons are
arranged in various orbitals around the nuclei of an atom each
component, e.g 4p5, consists of:
-
A number denoting the energy level, (4)
-
A letter denoting the type of orbital, (p)
-
A superscript denoting the number of electron in those orbitals (5)
HUND’S RULE

States that in subshells, as far as possible, electrons will half fill
orbitals with parallel spins.

The electron configuration for nitrogen therefore is 1s2 2s2 2p3 with
the p orbitals with one electron each, spinning in the same
direction.
ATOMIC NUMBER, MASS NUMBER, &
ISOTOPE

Atomic number : which equals its number of electrons if it is
uncharged. (number of protons or electrons)

Mass number : is the sum of the protons and neutrons in an
atom.(number of protons plus neutrons)

Isotope : same number of protons but different number of
neutrons or different mass number.
ATOMIC WEIGHT

Atomic and molecular can be measured with great accuracy
using a mass spectrometer. Because in the real world we use
large amounts of atoms and molecules, we use average masses
in calculations. Average mass is calculated from the isotopes of
an element weighted by their relative abundances.