General Chemistry I
Dr. PHAN TẠI HUÂN
Faculty of Food Science and Technology
Nong Lam University
Course Introduction
Recommended texts:
– William L. Masterton and Cecile N. Hurley. 2009.
Chemistry: Principles and Reactions, 6th edition.
Brooks/Cole Cengage Lerning.
– John A. Olmsted and Gregory M. Williams, 2005.
Chemistry, 4th edition. John Wiley & Sons, Inc.
– Steven S. Zumdahl, 2005. Chemical Principles, 5th
edition. Houghton Mifflin Company.
– Kenneth W. Whitten, Raymond E. Davis, Larry Peck,
George G. Stanley, 2003. General Chemistry, 7th edition.
Brooks Cole Publisher.
– Steven L. Hoenig, 2002. Basic Training in Chemistry.
Kluwer Academic Publishers.
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1
Course Introduction
Attendance requirements
• Class attendance: 45 hrs.
Assessment
• Method Weighting
– Mid-term examination
– Final examination
30%
70%
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Course Introduction
Why should we care about chemistry?
– Chemistry is everywhere!
– Chemistry helps us to understand and be better
informed about the world in which we live!
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Course Introduction
– Module 1: Atomic structure and Periodic Table.
– Module 2: Chemical bonds and molecular
configurations.
– Module 3: The Three States of Matter.
– Module 4: Chemical Thermodynamics.
– Module 5: Chemical kinetics.
– Module 6: Solutions.
– Module 7: Oxidation-reduction reactions and
transformation of chemical energy.
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Module 1: Atomic structure and Periodic Table
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Classification of Matter
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Dalton's Atomic Theory (1803)
• Matter is composed of indivisible particles - atoms.
• An element is composed of only one kind of atom. These
atoms in a particular element have the same properties
such as mass, size, or even shape.
• A compound is composed of two or more elements
combined in fixed ratios or proportions.
• In a chemical reaction, the atoms in the reactants
recombine, resulting in products which represent the
combination of atoms present in the reactants. In the
process, atoms are neither created, nor destroyed. So a
chemical reaction is essentially a rearrangement of atoms.
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Ramifications of Dalton's Theory
• The Law of Conservation of Mass states that mass is
neither created nor destroyed in a chemical reaction or
physical change.
• The Law of Definite Proportions states that every chemical
compound is made up of elements in a definite ratio by
mass.
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Ramifications of Dalton's Theory
Exercise 1: Calculate the mass of sodium chloride formed by
the complete reaction of 10.0 g of sodium with 15.4 g of
chlorine. What law allows this calculation?
• Ans:
10
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Ramifications of Dalton's Theory
• Exercise 2: Calculate the mass of oxygen that will
combine with 2.00 g of magnesium if 0.660 g of oxygen
reacts with 1.00 g of magnesium. What law allows this
calculation?
• Ans:
11
The discovery of the electron
Schematic drawing of a gas discharge tube in operation.
When a high voltage is applied to the two perforated
plates, an electrical discharge occurs between them. The
positively charged and negatively charged particles that
form in the gas move to collectors at the ends of the tube. 12
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J. J. Thomson experiment (1897)
• Schematic drawing of a cathode-ray tube. A beam of
electrons is deflected by a pair of charged plates (bent line),
but magnetic force can be adjusted to exactly counterbalance
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the effect of the electrical force (straight line).
Millikan's oil drop experiment (1909)
Schematic view of Millikan's
oil drop experiment. An
atomizer generated a fine mist
of oil droplets. Bombarding
the droplets with X rays gave
some of them extra negative
charge. In the presence of
sufficient electrical force,
these negatively charged
droplets could be suspended
in space.
• The measurements gave several different values, but the
charge was always equal to n(-1.6 x 10-19C).
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Thomson´s plum pudding model
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Rutherford's scattering experiment
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The general structrure of the atom
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Exercise
• What is the net charge on an atom that contains 8
protons, 8 neutrons, and 10 electrons?
• Ans:
18
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Constituents of the atom
• The atomic number (Z) denotes the number of protons in an
atom's nucleus.
• The mass number (A) denotes the total number of protons
and neutrons. Protons and neutrons are often called nucleons.
• Some atoms have the same atomic number, but different
mass numbers. This means different number of neutrons.
Such atoms are called isotopes.
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Isotopes
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Exercise
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Periodic chart of the elements
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Atomic mass unit (amu)
• Atomic mass unit (amu) is defined as exactly 1/12 the mass
of a 12C atom. The mass of the 12C atom is taken to be
exactly 12 amu.
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Atomic mass
• The atomic mass of an element is the weighted average of
the masses of the individual isotopes of the element.
• Example: Naturally occurring copper consists of 69.17%
63Cu and 30.83% 65Cu. The mass of 63Cu is 62.939 598
amu, and the mass of 65Cu is 64.927 793 amu. What is the
atomic mass of copper?
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12
Exercise
• Naturally occurring carbon consists of two isotopes, 12C
and 13C. What are the percentage abundances of the two
isotopes in a sample of carbon whose atomic mass is
12.01112?
• Ans:
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The concept of mole
• The quantity of a given substance that contains as many
units or molecules as the number of atoms in 12 grams of
carbon-12 is called a mole.
• The numerical value of one mole is 6.023 x 1023 and is
referred to as Avogadro´s number.
• One mole of hydrogen atoms contains Avogadro number
of hydrogen atoms.
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Atomic structure
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Electromagnetic radiation
In 1864, James Clerk Maxwell
developed a mathematical
theory to describe radiation as
wave-like, or oscillating,
electric and magnetic fields in
space. The electric and
magnetic fields are
perpendicular to each other.
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Electromagnetic spectrum
• Visible light, infrared radiation,
microwaves, radio waves,
ultraviolet, x-rays and gamma
rays are all types of
electromagnetic radiation.
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Wave character of light
• The distance between two waves,
usually measured from the peak of
the waves, is the wavelength, given
the symbol lambda, λ.
• The frequency is a statement of the
number of waves passing a point in
space per second; it is given the
symbol nu, ν.
(The hertz is commonly used as the unit
for frequency; 1 Hz = 1 s−1)
• The product of the wavelength and
the frequency is equal to the velocity
of light, usually designated by c:
c = λν
(The value of c can be rounded to
c=2.998 × 108 m/s for most
calculations.)
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Particle character of light
• The energy of light is emitted, absorbed, or converted to
other forms of energy in individual units referred to as
quanta (singular: quantum).
• The unit of light energy is often referred to as the particle
of light called the photon.
• The energy of a photon is proportional to the frequency:
ε = hν = (6.626 × 10−34 J · s)ν
Planck’s constant, h, is the universal proportionality
constant.
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Exercise
• When a metal bar is heated, it emits electromagnetic radiation
observable as the red to white glow of the metal. What is the
energy of one photon of red light with a wavelength of 700
nm? What is the energy of a mole of photons with this
wavelength? Calculate the energy of one photon and a mole
of photons of blue light with a wavelength of 400 nm.
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Exercise
• What is the wavelength in meters of the radiation from (a)
a low-range TV station broadcasting at a frequency of 55
MHz, (b) anAM radio station at 610 kHz, and (c) a
microwave oven operating at 14.6 GHz?
• Ans:
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Interaction of light with matter
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17
Atomic spectra
• In the late 19th century, Johann Balmer (1825–1898) and
Johannes Rydberg (1854–1919) showed that the
wavelengths of the various lines in the hydrogen
spectrum can be related by a mathematical equation:
• Here R is 1.097 x 107 m-1 and is known as the Rydberg
constant. The n’s are positive integers, and n1 is smaller
than n2
• However this is an empirical equation.
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Bohr's Model of Hydrogen Atom (1913)
1) In each hydrogen atom, the electron revolves around the
nucleus in one of the several stable orbits.
2) Each orbit has a definite radius and thus has a definite
energy associated with it.
3) An electron in an orbit closest to the nucleus has the lowest
energy, and if the electron is in the lowest orbit the atom is
said to be in its ground state.
4) The electron in an atom may absorb discrete amounts of
energy and move to another orbit with higher energy, and
this state is called the excited state.
5) An electron in an excited atom can go back to a lower energy
level and this process will result in the release of excess
energy as light.
6) The amount of energy released or absorbed.
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The radius is given by
r= n2a0
a0 =5.292 x 10-11 m = 0.5292 Å (Bohr radius)
The potential energy is given by:
where h Planck’s constant, m the
mass of the electron
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• Each line in the emission spectrum represents the difference
in energies between two allowed energy levels for the
electron.
• Comparing this to
the Balmer-Rydberg equation
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Bohr's Model of Hydrogen Atom
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The wave nature of the electron
• De Broglie (1925) proposed that not only does light have
the dual properties of waves and particles, but also
particles of matter have properties of waves. The
wavelength of those particle waves is given by
λ = h/ mv
where m and v are the mass and velocity of the particle.
Planck’s constant, h, is so small that the wavelengths are in
an observable range only for particles of atomic or
subatomic mass.
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Quantum mechanical picture of the atom
• The Heisenberg Uncertainty Principle, stated in
1927 by Werner Heisenberg (1901–1976):
It is impossible to determine accurately both the
momentum and the position of an electron (or any
other very small particle) simultaneously.
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Basic ideas of quantum mechanics
• Atoms and molecules can exist only in certain energy
states. In each energy state, the atom or molecule has a
definite energy. When an atom or molecule changes its
energy state, it must emit or absorb just enough energy to
bring it to the new energy state (the quantum condition).
• When atoms or molecules emit or absorb radiation (light),
they change their energies.The energy change in the atom
or molecule is related to the frequency or wavelength of the
light emitted or absorbed by the equations:
• The allowed energy states of atoms and molecules can be
described by sets of numbers called quantum numbers.
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Basic ideas of quantum mechanics
• The mathematical
approach of quantum
mechanics involves
treating the electron in an
atom as a standing wave
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Orbitals
• A solution to the Schrödinger equation for an electron
must satisfy three quantum conditions corresponding to the
three dimensions of space.
•
Each quantum condition introduces an integer, called a
quantum number, into the solution.
• A separate solution, describing a probability distribution of
finding the electron at various locations, exists for each
allowed set of three quantum numbers.
• Such a solution is called an orbital.
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Quantum number
• Principal quantum number (n). The principal quantum
number denotes the energy level of electrons. The larger the
principal quantum number is, the larger the energy.
• The orbital size depends on n. This means that the larger
the n value, the larger the orbital. Orbitals with the same n
belong to the same shell.
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Quantum number
• Angular momentum quantum number (l). Angular
momentum quantum number denotes the shape of the
orbital. The values range from 0 to n – 1.
• The angular momentum quantum numbers correspond to
different subshells.
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Quantum number
• Magnetic quantum number (ml). Magnetic quantum
numbers define the different spatial orientations of the
orbitals. The values range from –l to +l.
• There are three p orbitals corresponding to ml = 1, 0, and
−1. However, it is usually more convenient in chemistry to
use a new set of three orbitals oriented along the x, y, and z
axes to display the shapes and directions of these orbitals.
• Further, there are 5 d orbitals and 7 f orbitals having
different shapes and orientations in space.
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Orbitals
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Orbitals
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Quantum number
• Spin quantum number (mS) Spin quantum number has
to do with the spin orientations of an electron. The two
possible spins are denoted by the spin quantum numbers
+ ½ and -1/2.
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Quantum number
• The values of n, l, and ml describe a particular atomic
orbital. Each atomic orbital can accommodate no more
than two electrons, one with ms=+1/2 and another with
ms=-1/2
.
• A set of quantum numbers: (n, l, ml, ms).
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Permissible Values of the Quantum Numbers
Through n=4
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“Building” of Electron configurations
by the Aufbau Principle
• No two electrons in an atom may have identical sets of four
quantum numbers (Pauli exclusion principle).
• Orbitals are filled in the order of increasing energy
(Klechkowski´s rule).
• Electrons occupy all the orbitals of a given subshell singly
before pairing begins. These unpaired electrons have
parallel spins (Hund´s rule).
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Electron configurations
Two general rules help us to
predict electron configurations.
• Electrons are assigned to
orbitals in order of increasing
value of (n + l ).
• For subshells with the same
value of (n+l ), electrons are
assigned first to the subshell
with lower n.
1s < 2s < 2p <3s < 3p<4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f <
5d...
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Exercise
• Arrange the electrons in the following list in order
of increasing energy, lowest first:
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Exercise
• How many electrons are permitted in each of the
following subshells? (a) 2s (b) 6p, and (c) 4d.
• Ans:
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Exercise
• How many electrons are permitted in each of the
following subshells?
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Exercise
• Write the electronic configuration of sulfur and
also show the filling of electrons with orbital
notation.
• Ans:
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Exercise
• Write detailed electronic configurations for (a) N (Z=7),
(b) P (Z=15), (c) As (Z=33), and (d) Sb (Z=51).
• What makes their chemical properties similar?
• Ans:
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The periodic table
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The periodic table
• The properties of elements are periodic functions of their
atomic numbers
• The vertical columns of elements represented in the
periodic table are called groups, and the horizontal rows
are called periods.
• There are seven periods in the periodic table. The groups
are usually designated by roman numerals followed by the
letter A or B as shown in the periodic table.
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The periodic table
• The groups IA through VIIA are called the representative
elements. These elements have either s or p orbital valence
electrons. The last group in the periodic table is the noble
gas group otherwise known as the zero group.
• The groups ranging from IB through VIIIB are called
transition metals, and finally the metals from lanthanum
through hafnium and metals from actinium onward are
called the inner transition metals.
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Period 1
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Period 2
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Period 3
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Period 4
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Periodicity of Atomic Properties
As principal quantum number n increases, atomic orbitals
become larger and less stable. As atomic number Z increases,
any given atomic orbital becomes smaller and more stable.
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Atomic size
Exercise: Arrange the following elements in terms of
increasing atomic radius: Mg, Cl, K and Cs.
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Atomic size
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Ionization energy
• Ionization energy (IE) is the minimum amount
of energy required to remove an electron from an
atom.
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Ionization energy
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Ionization energy
• First IE < Second IE < Third IE < Fourth IE < ...
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Electron affinity
• Electron affinity is the energy change associated
with the addition of an electron to a gaseous atom.
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Electron affinity
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Electron affinity
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Electronegativity
• The relative tendency of an atom to attract the bonding
electrons to itself is called electronegativity. The popularly
used electronegativity scale is based on a system called
Pauling's scale, according to which fluorine (the most
electronegative element) has an electronegativity value of
4.0. Nonmetals are the most electronegative elements.
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Electronegativity Values of the Elements
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Summary
After you have studied this module, you should be able to
• Describe the evidence for the existence and properties of
electrons, protons, and neutrons.
• Predict the arrangements of the particles in atoms.
• Describe isotopes and their composition.
• Calculate atomic weights from isotopic abundance.
• Descriptions of waves play an important role in our
theories of light and of atomic structure.
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Summary
• Describe the four quantum numbers, and give possible
combinations of their values for specific atomic
orbitals.
• Describe the shapes of orbitals and recall the usual
order of their relative energies.
• Write the electron configurations of atoms.
• Relate the electron configuration of an atom to its
position in the periodic table.
• Periodicity of atomic properties.
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