The Electronic
Structure of Atoms
CHAPTER 1
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1. The Atomic Theory
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The Atomic Theory
1.
In the fifth century b.c., the Greek philosopher Democritus expressed the belief that all matter consists of very
small, indivisible particles, which he named atomos (meaning uncuttable or indivisible)
2.
Dalton’s work marked the beginning of the modern era of chemistry and was summarized in : All atoms of a
given element are identical, having the same size, mass, and chemical properties and a chemical reaction
involves only the separation, combination, or rearrangement of atoms; it does not result in their creation or
destruction.
3.
An English physicist, J. J. Thomson, used a cathode ray tube and his knowledge of electromagnetic theory to
determine the ratio of electric charge to the mass of an individual electron.
4.
Rutherford proposed, that the atom’s positive charges are all concentrated in the nucleus, a dense central core
within the atom
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2. The Structure of the Atom
2.1.The Electron
According to electromagnetic theory, a moving charged body behaves like a magnet and can interact with electric
and magnetic fields through which it passes.
• The experiment consist in a cathode ray produced in a discharge tube traveling from the cathode (left) to the anode (right).
The ray itself is invisible, but the fluorescence of a zinc sulfide coating on the glass causes it to appear green.
• Two charged deflection plate are added to either side of the cathode ray tube and a current was
applied.
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 it must consist of negatively charged particles. We know these negatively charged particles as electrons.
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2.2.The Proton and the Nucleus
Experiment.2: Rutherford devised a model of atomic structure, suggesting that most of the atom must be empty space.
This structure would allow most of the alfa particles (Alpha radioactive rays that consist of positively charged particles) to
pass through the gold foil with little or no deflection. The atom’s positive charges, Rutherford proposed, are all
concentrated in the nucleus
 Rutherford’s experimental design for measuring the scattering of a particles by a piece of gold foil shows that :
• Most of the alfa particles passed through the gold foil with little or no deflection.
• A few were deflected at wide angles. Occasionally an alfa particle was turned back.
The positively charged particles in the nucleus are called protons.
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2.3.The Neutron
A typical atomic radius is about 100
pm, whereas the radius of an atomic
nucleus is only about 510-3 pm.
in 1932 , James Chadwick, bombarded a thin sheet of beryllium with alfa particles, a very high energy radiation similar to
gamma rays was emitted by the metal. (The third type of radioactive radiation consists of high-energy rays called gamma
rays . Like X rays, ɣ rays have no charge and are not affected by an external field).
Later experiments showed that the rays actually consisted of electrically neutral particles having a mass slightly greater
than that of protons. Chadwick named these neutrons.
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Atomic Number, Mass Number and Isotopes
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3. The Quantum Mechanical Description of the Hydrogen Atom
Trying to locate a subatomic particle that behaves like a wave, the German physicist Heisenberg formulated what is
now known as the Heisenberg uncertainty principle:
It is impossible to know simultaneously both the momentum (mass times velocity) and the position of a particle with
certainty.
Although quantum mechanics tells us that we cannot pinpoint an electron in an atom, it does define the region where
the electron might be at a given time.
The concept of electron density gives the probability that an electron will be found in a particular region of an atom.
An atomic orbital can be thought of as the wave function of an electron in an atom means the distribution of the
electron density in 3D space around the nucleus.
• A representation of the
electron density distribution
surrounding the nucleus in
the hydrogen atom. It shows
a high probability of finding
the electron closer to the
nucleus.
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3.1. Quantum Numbers
To describe the distribution of electrons in hydrogen and other atoms and label them we need to use
quantum numbers.
The principal quantum number,
The angular momentum quantum number,
The magnetic quantum number
The Electron Spin Quantum Number (ms)
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3.1.1. The Principal Quantum Number (n)
• The principal quantum number (n) can have integral values 1, 2, 3, and so forth.
• In a hydrogen atom, the value of n determines the energy of an orbital, but this is not the case for a many-electron atom.
• The principal quantum number also relates to the average distance of the electron from the nucleus in a particular orbital.
 The larger n is, the greater the average distance of an electron in the orbital from the nucleus and therefore the larger the
orbital.
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3.1.2. The Angular Momentum Quantum Number
The angular momentum quantum number (l) tells us the “shape” of the orbitals.
The values of l, depend on the value of the principal quantum number, n. For a given value of n, l, has possible integral
values from 0 to (n -1).
The value of l, is generally designated by the letters s, p, d, . . as follows:
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3.1.3. The Magnetic Quantum Number (m,l)
The magnetic quantum number (m,l) describes the orientation of the orbital in space.
Within a subshell, the value of m,l depends on the value of the angular momentum quantum number, l ,.
For a certain value of ,l, there are (2l+1) integral values of m,l as follows:
• Values are integers ranging from –l to l. example, for l=1 then m,l=3, values are -1,0,+1
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3.1.4. The Electron Spin Quantum Number (m,s)
If electrons are thought of as spinning on their own
axes, as Earth does, their magnetic properties can be
accounted for.
According to electromagnetic theory, a spinning charge
generates a magnetic field, and it is this motion that
causes an electron to behave like a magnet.
To take the electron spin into account, it is necessary to
introduce a fourth quantum number, called the
electron spin quantum number (ms), which has a value
of +½ or -1/2
The magnetic fields generated by
these two spinning motions are
analogous to those from the two
magnets. The upward and downward
arrows are used to denote the
direction of spin.
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3.2. Atomic Orbitals
One of the important questions we ask when studying the properties of atomic orbitals is, What are the
shapes of the orbitals?
3.2.1. S Orbital
All s orbitals are spherical in shape but differ in size, which increases as the
principal quantum number increases.
hydrogen 1s, 2s, and 3s orbitals. Each
sphere contains about 90 percent of the
total electron density. All s orbitals are
spherical. Roughly speaking, the size of
an orbital is proportional to n squared,
where n is the principal quantum number
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3.2.2. p Orbitals
The p orbitals start with the principal quantum number n =2.
If n = 1, then the angular momentum quantum number, can assume only the value of L= zero; therefore, there is only a 1s orbital.
l= 1, the magnetic quantum number m,l can have values of 3. Starting with n = 2 and , l =1, we therefore have three 2p orbitals: 2px, 2py,
and 2pz
 The boundary surface diagrams of p orbitals show that each p orbital can be thought of as two lobes on opposite
sides of the nucleus.
• These orbitals are identical in shape and energy, but their orientations are different .
• Like s orbitals, p orbitals increase in size from 2p to 3p to 4p orbital and so on.
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3.2.3. d Orbitals and Other Higher-Energy Orbitals
When , l=2, there are five values of m,l, which correspond to five d orbitals. The lowest value of n for a d orbital is 3.
• All the 3d orbitals in an atom are identical in energy.
• The d orbitals for which n is greater than 3 (4d, 5d, . . .)have similar shapes.
• Orbitals having higher energy than d orbitals are labeled f, g, . . . and so on.
• The f orbitals are important in accounting for the behavior of elements with atomic numbers
greater than 57, but their shapes are difficult to represent
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Relation Between Quantum Numbers and Atomic Orbitals
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