Course code : CHEM 11122
Course title : General Chemistry and Basic
Analytical Chemistry
Course content:
Structure and bonding (10 h)
Aqueous Solution Chemistry (20 h)
Observations in the macroscopic world
Matter is composed of atoms.
Understanding the structure of
atoms is critical to understand the
properties of matter
History of atomic theory
A revision of basic concepts
A Brief History of Atomic
Theory
This section will focus on
scientists who have had an
impact on the study of the
atom.
Democritus
Democritus proposed that
matter cannot be broken down
indefinitely.
470-380 B.C.
At some point you end up with a
piece that can’t be divided. That
smallest piece he called an atom,
from the Greek word atomos,
which means “indivisible”.
Democritus’ Model
“Atomos”
ATOMOS was the word Democritus used the point,
or stage
cannot
be broken down any
ATOMOS
waswhere
thematter
word
Democritus
used
further. ATOMOS literally means “indivisible”
the
point, or stage where matter cannot be broken
down any further. ATOMOS literally means
“indivisible”
The Birth of Modern Atomic Theory
John Dalton
• John Dalton was a British
chemist.
• He was the first modern
scientist to propose the
existence of atoms.
• He described an atom as
an invisible indestructible,
solid sphere, like a billiard
ball.
1766 - 1844
Dalton’s Model
The “Indivisible Sphere”
John Dalton
• In 1803, proposed an Atomic Theory which
states:
-Elements are made out of minute,
indivisible particles –Atoms
-Atoms of the same element are
exactly alike, and atoms of
different elements are different
- Atoms can not be created or destroyed
Image taken from:
chemistry.about.com/.../JohnDalton.htm
-Compounds are formed union of two or more
atoms of different elements in a simple whole
number ratios.
Subatomic Particles
J.J. Thomson
(1856 – 1940)
• J.J. Thomson was a British
physicist
• Proved that an atom can be
divided into smaller parts
Image taken from:
www.wired.com/.../news/2008/04/d
ayintech_0430
• While experimenting with
cathode-ray tubes,
discovered corpuscles,
which were later called
electrons
•Stated that the atom is neutral
•In 1897, proposed the Plum Pudding Model
which states that atoms mostly consist of
positively charged material with negatively
charged particles (electrons) located
throughout the positive material
Thomson “plum pudding” model
+
atoms mostly consist of positively charged material
with negatively charged particles (electrons) located
throughout the positive material
A = alpha
B = gamma
C = beta
J.J. Thomson, measured charge/mass of e(1906 Nobel Prize in Physics)
Rutherford’s experiment
(1871-1937)
Experimented with a radiation source that
sent out alpha particles through a thin piece
of gold foil to a detector screen that glowed
when it was hit.
•Suggested the following characteristics of the
atom:
oIt consists of a small core, or nucleus, that
contains most of the mass of the atom
oThis nucleus is made up of particles called
protons, which have a positive charge
oThe protons are surrounded by negatively
charged electrons, but most of the atom is
actually empty space
Ernest Rutherford experiments proved that atoms are
mostly empty space.
Discovered the nucleus, which contains positively
charged particles .
1932 A.D.
English scientist
James Chadwick
Using alpha particles
discovered a neutral atomic
particle with a mass close to
a proton. What he
discovered was the neutron.
-Discovered the third subatomic particle of the atom:
the neutron 450 BC 400 BC 350 BC
500-1600 AD
1650 AD Late 1700’s 1808AD 1831AD 1879AD1897AD1898AD
1911AD
1922AD
1932AD
1922 A.D.
Danish scientist
Niels Bohr
Came up with experimental
evidence proving that electrons
exist in energy levels (shells)
orbiting around a positively
charged nucleus.
450 BC 400 BC 350 BC
500-1600 AD
1650 AD Late 1700’s 1808AD 1831AD 1879AD1897AD1898AD
1911AD
1922AD
1932AD
The arrangement of electrons in an atom is called
its electronic structure.
Much of our present understanding of the electronic
structure of atoms has come from analysis of
the light emitted or absorbed by substances.
-To understand the basis for current model of
electronic structure we must first learn more about
light
Light and atomic spectra
Light
-The light we see with our eyes,
visible light, is one type of electromagnetic
radiation.
-There are many types of electromagnetic
radiation in addition to visible light.
-radio waves that carry music to our radios,
-infrared radiation (heat) from a glowing
fireplace,
-X-rays
-All types of electromagnetic radiation move
through a vacuum at , 3.00 x 108 m/s
the speed of light.
-All have wave-like characteristics similar to
those of waves that move through water.
-The distance between two adjacent peaks (or
between two adjacent troughs) is called the
wavelength
- The number of complete wavelengths, or cycles,
that pass a given point each second is the
frequency of the wave
There is an inverse relationship between the
frequency and wavelength of electromagnetic
radiation
Practice Exercise 1
Electromagnetic Spectrum
-shows
the various types of electromagnetic
radiation arranged in order of increasing
wavelength
the wavelengths of radio
waves can be longer than
a football field
The wavelengths of
gamma rays are
comparable to the
diameters of atomic
nuclei
Wave-like Behaviors of Light
How Do We Know that Light is a Wave?
Light reflects in the same manner that any wave
would reflect.
Light refracts in the same manner that any wave
would refract.
Light diffracts in the same manner that any
wave would diffract.
Since light behaves like a wave, one would
have good reason to believe that it might be
a wave
-Although the wave model of light explains many
aspects of its behavior, this model cannot
explain several phenomena
(1) the emission of electrons from metal
surfaces on which light shines
(the photoelectric effect),
(2) the emission of light from electronically
excited gas atoms (emission spectra)
In 1900 a German physicist named Max Planck
(1858–1947)
energy can be either released or absorbed by
atoms only in discrete amounts called “quanta”
quantum = (meaning “fixed amount”) to the
smallest quantity of energy that can be emitted
or absorbed as electromagnetic radiation.
He proposed that the energy, E, of a single
quantum equals a constant times the
frequency of the radiation:
-constant h is called Planck’s constant
and has a value of
= 6.626 x 10-34joule second (J s).
-Because the energy can be released
only in specific amounts, we say that the
allowed energies are quantized
-Planck’s revolutionary proposal that energy is
quantized was proved correct, and he was
awarded the 1918 Nobel Prize in Physics for
his work on the quantum theory
In 1905, Albert Einstein (1879–1955) used
Planck’s theory to explain the photoelectric
effect
The Photoelectric Effect
The photoelectric effect is the observation that
many metals emit electrons when light shines
upon them.
A minimum frequency of light, different for
different metals, is required for the
emission of electrons
-the radiant energy striking the metal surface
behaves like a stream of tiny energy packets.
Each packet, which is like a “particle” of
energy, is called a photon
Light has particle like properties
Under the right conditions, photons striking a
metal surface can transfer their energy
to electrons in the metal.
A certain amount of energy—called the work
function —is required for the electrons to
overcome the attractive forces holding them in
the metal.
If the photons striking the metal have less
energy than the work function, the electrons do
not acquire sufficient energy to escape from
the metal, even if the light beam is intense.
If the photons have energy greater than the
work function of the particular metal the
excess appears as the kinetic energy of the
emitted electrons
Einstein won the Nobel Prize in Physics in 1921
for his explanation of the photoelectric
effect.
Einstein’s theory of light
as a stream of photons rather than a wave
Light has both wave and particle like properties
Practice Exercise 2
Line Spectra
In 1913, the Danish physicist Niels Bohr offered
a theoretical explanation of line spectra, another
phenomenon that had puzzled scientists
during the nineteenth century.
spectrum
A spectrum is produced when radiation from
sources( light bulbs and stars) is separated
into its component wavelengths
A continuous visible spectrum is produced
when a narrow beam of white light is passed
through a prism.
The white light could be sunlight
Not all radiation sources produce a continuous
spectrum.
Atomic emission of hydrogen and neon
Different gases emit light of different
characteristic colors when an electric current is
passed through them.
When light coming from such tubes is passed
through a prism, only a few wavelengths are
present in the resultant spectra
Line Emission Spectrum of Hydrogen Atoms
four narrow bands of bright light are observed
against a black background
Line spectrum of neon
Each colored line in such spectra represents
light of one wavelength.
Every element has a unique emission spectrum
A spectrum containing radiation of only specific
wavelengths is called a line spectrum.
Atomic Spectrum of Hydrogen
• In 1885 Johann Balmer showed that the wave
length of any line in the visible spectrum of
atomic hydrogen, could be given by the simple
formula;
Later, additional lines were found in the
ultraviolet and infrared regions of hydrogen’s
line spectrum.
Soon Balmer’s equation was extended to a more
general one, called the Rydberg equation, which
allows us to calculate the wavelengths of all the
spectral lines of hydrogen
How could the remarkable simplicity of this
equation be explained?
Bohr theory
To explain the line spectrum of hydrogen, Bohr
assumed that
electrons in hydrogen atoms move in circular
orbits around the nucleus like planets around
the sun
Several Problems arise with this concept
1. Electrons are expected to slow down
gradually.
2. Why does electrons move in an orbit around
nucleus?
3. Since nucleus and electrons have opposite
charges, they should attract each other and
collide.
To explain these problems Bohr postulates
1. Electrons do not radiate energy if
stay in one orbit, ∴ do not slow down.
2. When electrons move from one orbit
to another they radiate or absorb
energy.
3. For an electron to remain in orbit the
electrostatic attraction between the
nucleus and electron must be equal to
the centrifugal force.
The Bohr Model of the Atom
Electronic transition
and origin of the
spectral lines of
hydrogen atom
Bohr assumed
that the discrete
lines seen in the
spectrum of the
hydrogen atom
were due to
transitions of an
electron from
one
allowed
orbit/energy
level to another.
What is most significant about Bohr’s model is that
it introduces two important ideas;
1. Electrons exist only in certain discrete
energy levels, which are described by
quantum numbers.
2. Energy is involved in the transition of an
electron from one level to another.
Limitations of Bohr's model of the atom
-The Bohr model could only successfully explain
the hydrogen spectrum.
-It could NOT accurately calculate the
spectral lines of larger atoms.
-The model only worked for hydrogenlike atoms
That is, if the atom had only one electron.
Depending on the experimental
circumstances,
radiation appears to have either a wave-like
or a particle-like (photon) character.
Wave-particle duality of matter
• In1924, Prince Louis Victor de Broglie(18921987)- French physicist- proposed that all
particles of matter (from single atoms to large
objects) moving at some velocity would have the
properties of a wave.
De Broglie used the term matter waves to
describe the wave characteristics of material
particles.
The Heisenberg Uncertainty Principle
• In 1927,Werner Heisenberg(1901-1976) German physicist- no experimental method
can be devised that will measure
simultaneously the precise position as well
as the momentum of an object.
The Heisenberg Uncertainty Principle
• It is impossible to know simultaneously both
the position and the momentum of an
object/electron exactly.
• ∴ The probability of finding an electron in a
particular position or in a particular volume of
space has to be considered.
Practice Exercise 4
Thus, we have essentially no idea where the
electron is located in the atom.
• ∴ The probability of finding an electron in
a particular position or in a particular volume
of space has to be considered.
QUANTUM
MECHANICS
AND ATOMIC ORBITALS
Quantum mechanics was used in 1926 by
Erwin Schrödinger(1887-1961) – Austrian
physicist – to describe electrons in atoms.
-proposed an equation, now known as Schrödinger’s
wave equation
Solving Schrödinger’s equation for the hydrogen
atom leads to a series of mathematical
functions called wave functions.
-These wave functions are usually
represented by the symbol
(lowercase
Greek letter psi).
-the square of the wave function,
at a given point in space represents the
probability that the electron will be found at
that location. For this reason, is called either
the probability density or the electron
density.
Where in the
figure is the region of
highest electron
density?
-Electron-density distributionThis rendering represents the probability of finding the
electron in a hydrogen atom in its ground state
Quantum numbers for electrons
The Bohr model was a one-dimensional model
that used one quantum, number to describe the
distribution of electrons in the atom/ orbit.
The quantum mechanical model uses three
quantum numbers, n, l, and , ml which result
naturally from the mathematics used, to describe
the distribution of electrons in the atom/orbitals
Quantum numbers for electrons
1. Principal Quantum Number (n)
n = integral values of 1, 2, 3, 4, …….
K L M N ……..
Determines the main energy level (shell) the
electron is in.
As n increases, the orbital becomes larger, and
the electron spends more time farther
from the nucleus.
An increase in n also means that the electron has
a higher energy and is therefore less tightly
bound to the nucleus.
2. Azimuthal (angular momentum) Quantum
Number (l)
l = 0, 1, 2, 3,…….(n-1)
-Describes the sub-shell that the electron
occupies. -This quantum number defines the
shape of the orbital.
3. Magnetic Quantum Number (ml)
can have integral values between -l and l,
including zero.
ml = +l…..0…..-l
This quantum number describes the
orientation of the particular orbital
that the electron occupies in space
•For s sublevel, l = 0
m=0
1 s type
•For p sublevel, l = 1
•
•
3 p types
m = +1, 0, -1
•For d sublevel, l = 2
m = +2,+1, 0, -1,-2
•
5 d types
•For f sublevel, l = 3
m = +3,+2,+1, 0, -1,-2,-3
7 f types
Practice problem 05
4. Spin Magnetic Quantum Number (ms)
ms = +1/2 or –1/2
-The electron behaves as if it were spinning about
an axis, thereby generating a magnetic field
whose direction depends on the direction of
spin.
-The two directions for the magnetic field
correspond to the two possible values for the spin
quantum number
Zeeman effect
•In 1896 Pieter Zeeman (1865-1943) - Dutch
physicist - discovered that the spectral lines of
a light source subjected to a strong magnetic
field were split into several components.
•(This phenomenon, known as the Zeeman effect.
A photo Zeeman took of the Zeeman effect
Comparing
probability
density
and radial
probability
function .
Radial probability distributions for the 1s,
orbital of hydrogen
Comparison of the 1s,
2s, and 3s orbitals
The Pauli exclusion principle
In 1925 the Austrian-born physicist Wolfgang
Pauli (1900–1958) discovered the principle that
governs the arrangements of electrons in manyelectron atoms.
The Pauli exclusion principle states that no
two electrons in an atom can have the same
set of four quantum numbers n, l, ml and ms.
-an orbital can hold a maximum of two
electrons and they must have opposite spins.
PRACTICE EXERCISE 06
ELECTRON CONFIGURATIONS
The way electrons are distributed among the
various orbitals of an atom is called the
electron configuration of the atom.
the orbitals are filled in order of increasing
energy, with no more than two electrons per
orbital.
A half arrow pointing up represents an electron
with a positive spin magnetic quantum number and
a half arrow pointing down represents an electron
with a negative spin magnetic quantum number
.
Electrons having opposite spins are said to be
paired when they are in the same orbital.
An unpaired electron is one not accompanied by
a partner of opposite spin.
In the lithium atom the two electrons in the
1s orbital are paired and the electron in the
2s orbital is unpaired.
Order of Filling Orbitals
Hund’s rule
•Within a sublevel,each orbital is occupied by
one electron before any orbital has
two(pairing).
•The first electrons to occupy orbitals with in a
sublevel must have parallel spins.
Hund’s rule, which states that for
degenerate orbitals, the lowest energy is
attained when the number of electrons having
the same spin is maximized.
Aufbau principle
General energy ordering of orbitals for a
many-electron atom.
•The filling begins with the sublevel lowest in energy and
continues upwards according to the “Aufbau principle”
PRACTICE EXERCISE 07
Condensed Electron Configurations
-In writing the condensed electron
configuration of an element, the electron
configuration of the nearest noble-gas element
of lower atomic number is represented by its
chemical symbol in brackets.
Anomalous Electron Configurations
The electron configurations of certain elements
appear to violate the rules we have just
discussed
This anomalous behavior is largely a
consequence of the closeness of the 3d and
4s orbital energies.
It frequently occurs when there are enough
electrons to form precisely half-filled sets of
degenerate orbitals (as in chromium) or a
completely filled d subshell (as in copper).
ELECTRON CONFIGURATIONS
AND THE PERIODIC TABLE
Dmitri Mendeleev
In 1869 he published a table of
the elements organized by
increasing atomic mass
1834 - 1907
The Modern Periodic Table
Element Organization
The Periodic Law:
Elements are arranged by atomic #,
Periodic table of the elements
(columns)
Groups
Periods
(rows)
Elements are presented in the periodic table
by increasing values of their atomic numbers,
the number of protons in their atomic nuclei
Symbols in Table
periodic table blocks
-named according to the subshell in which the
"last" electron resides.
Chemical compounds
A compound is a substance containing more than
one element
CHEMICAL BONDS
Two of the most common substances on our
dining table are salt and granulated sugar
NaCl
C12H22O1
1
The properties of substances are determined in
large part by the chemical bonds that hold their
atoms together
CHEMICAL BONDS
Atoms or ions are held together in
molecules or compounds by chemical bonds.
TYPES OF CHEMICAL BONDS
• Ionic bonds
• Covalent bonds
• Metallic bonds
• Coordinate bond
Ionic Bond
• Refers to electrostatic forces that exist
between ions of opposite charges.
• Between atoms of metals and nonmetals with
very different electronegativity
• Conductors and have high melting point.
• Examples; NaCl
-Ionic bond – electron from Na is transferred to Cl.
-The Na becomes (Na+) and the Cl becomes (Cl-),
charged particles or ions.
Atoms often gain, lose, or share electrons to
achieve the same number of electrons as the
noble gas closest to them in the periodic table.
The noble gases have very stable electron
arrangements
Octet rule:
Atoms tend to gain, lose, or share electrons
until they are surrounded by eight valence
electrons.
Metallic Bond
Metallic Bond
• Metallic bonding constitutes the electrostatic
attractive forces between the delocalized
electrons, called conduction electrons, gathered
in an electron cloud and the positively charged
metal ions
• Good conductors at all states,
• Examples; Na, Fe, Al, Au, Co
Covalent Bond
• Between nonmetallic elements of similar
electronegativity.
• Formed by sharing electron pairs
• Stable non-ionizing particles, they are not
conductors at any state
• Examples; O2, CO2, C2H6, H2O,
Lewis Structures
-we
usually show each shared electron pair as a
line and any unshared electrons as dots.
- the Lewis structures for H2 and Cl2 are
MOLECULAR GEOMETRY
AND
BONDING THEORIES
Valence Shell Electron Pair
Repulsion (VSEPR) Theory
1. Shapes of molecules are determined by the
repulsions between electron pairs in the valence
shell
2. A lone pair takes up more space around the
central atom than a bond pair. Hence repulsion
between
lone pair – lone pair > bond pair-lone pair > bond
pair- bond pair
The best arrangement of a given number of
electron domains is the one that minimizes the
repulsions among them.
Trigonal planar
Bent
0
1
2
Bent
Valence Bond Theory/VB theory
-Covalent bonding occurs when a valence
atomic orbital of one atom overlaps, with a valence
atomic orbital of another atom.
-The overlap of orbitals allows two electrons of
opposite spin to share the space between the nuclei,
forming a covalent bond.
Covalent bonds in H2, HCl, and Cl2 result from
overlap of atomic orbitals
Potential Energy Curve
-shows how the potential energy of a system
consisting of two H atoms changes as the atoms come
together to form an H2 molecule
-When the atoms are infinitely far apart, they
do not “feel” each other and so the energy
approaches zero.
-As the distance between the atoms decreases,
the overlap between their 1s orbitals increases
(the potential energy of the system decreases.
That is, the strength of the bond increases;
attractive forces dominate
-shows that as the atoms come closer together
than 0.74 Å, the energy increases sharply (is
due mainly to the electrostatic repulsion
between the nuclei)
-The internuclear distance, or bond
length, is the distance that corresponds to the
minimum of the potential-energy curve (This is
moe correctly known as the equilibrium bond
length).
-The potential energy at this minimum
corresponds to the bond strength.
Sigma bonds
•End to end or head on overlap of orbitals result in
bonds.
•The electron density is concentrated in between
the two atoms.
-the line joining the two nuclei passes
through the middle of the overlap region.
These bonds are called sigma bonds.
Pi (П) bonds
• Multiple bonds form by the sideway overlap of
orbitals, known as p orbitals.
• A bond is one in which the overlap regions lie
above and below the internuclear axis.
Hybridization and hybrid orbitals
The process of mixing atomic orbitals is a
mathematical operation called hybridization.
Formation of two equivalent Be-F
bonds in BeF2.
How many atomic orbitals contribute to
form the three sp2 hybrid orbitals?
Formation of sp2 hybrid
Formation of sp3 hybrid orbitals
Molecular orbital theory
Molecular orbital theory describes the electrons
in molecules by using specific wave functions
called molecular orbitals (MO).
The Hydrogen Molecule
The two molecular orbitals of H2, one bonding MO
and one an antibonding MO.
-In the bonding MO electron density is
concentrated in the region between the two
nuclei.
-By contrast, the antibonding MO has very
little electron density between the nuclei.
-The relative energies of two 1s atomic
orbitals and the molecular orbitals formed
from them are represented by an energy-level
diagram (also called a molecular orbital
diagram
Energy-level diagrams and electron
configurations for H2