Dr. Saidane
Chem. 152
The Modern Atom
Skills you should have mastered
Conceptual
1.
Explain the correlation between light, energy, and photons.
2.
Explain how atoms emit light in a form of photons carrying energy.
3.
Compare and contrast the Bohr’s Model to the quantum model of an atom.
4.
List the four quantum numbers, and describe their significance.
5.
Distinguish and sketch the boundary surfaces of s-, p-, and d- orbitals.
6.
Determine the electron configuration of an element using the Aufbau
principle, the Pauli Exclusion Principle and Hund’s rule.
7.
Distinguish between and orbital diagram and an electron configuration.
8.
Explain the basis for the arrangement of the periodic table.
9.
Interpret the periodic table.
10.
Identify metals, Nonmetals, and metalloids on the periodic table.
11.
Relate the properties of various elements to their electron configurations.
12.
Define atomic radii and interpret the trend shown by atomic radii within the
periodic table.
13.
Define ionic radii and interpret the trend shown by ionic radii within the
periodic table.
14.
Define Ionization energy and interpret the trend.
Problem-solving
1.
For each energy level, determine the number of sublevels and orbitals
2. Determine the number of electron for each orbital, sublevel, and level.
3. Write the electron configuration and the orbital diagram for an element or ion.
4. Use the periodic table to write electron configurations and orbital diagrams for
various atoms.
5. Use the noble gas configuration to write the electron configuration.
6. Use the electron configuration to find the position of the element in the
periodic table.
7. Use the periodic table to predict and explain trends in atomic radius, ionic
radius, and ionization energy.
_________________________________________________________________________
Atoms emit light of many colors when their compounds are burned or exposed to
electric discharge. In order to understand their atomic behavior we need to understand
light and its properties.
Light
 Light is made of photons carrying energy.
 Photons, move as waves. Each wave has a wavelength, amplitude, a frequency, and a
speed.
o The wavelength,  ( is a Greek letter pronounced lambda) is the peak-to-peak
distance of the wave.
o Amplitude is the height of a peak. It depends on the intensity of the radiation.
The higher the peak of the wave, the higher its intensity.
o A wave oscillates in space and time both in direction and strength. The number
of oscillations is called frequency,  ( is a Greek letter pronounced nu), of the
radiation. The unit of frequency is hertz and is defined as 1 cycle per second (1
hz = 1/1s).
o A wave moves with a velocity = frequency x wavelength, v =  x 
Quantum Concept
 A quantum is the minimum amount of energy that can be lost or gained by an atom.
Atoms emit energy in a form of electromagnetic wave. The quantum concept states
that energy is present in small, discrete bundles. For example: A tennis ball that rolls
down a ramp loses potential energy continuously. A tennis ball that rolls down a
staircase loses potential energy in small bundles. The loss is quantized.

Later Louis De Broglie suggested that matter exhibit wave-like behavior and the
wavelength of matter is inversely proportional to the mass  = h/mv. Only
particles the size of an electron and moving at the speed of light or close to the
speed of light have large enough frequencies to be considered as waves. Electrons
in an atom are both particles and waves, which means that they are particles
carrying energy and moving like waves.
Bohr’s Model
 When an electric current is passed through a tube containing hydrogen gas at low
pressure, a series of lines were observed (called spectral lines or spectrum). In order to
explain hydrogen atomic spectra, Niels Bohr suggested that the electron of an atom of
hydrogen were confined to specific energy states. The electron moves around the
nucleus in definite discrete orbits. Electrons are found only in specific energy levels,
and nowhere else in between. The electron energy levels are quantized. When the
electron is in the orbit closest to the nucleus, the atom occupies the ground level (n=1),
which has the lowest energy level E1 (E for n=1).
 In the Bohr’s model of hydrogen; the energy levels that the electron can occupy are
similar to the lanes in a freeway. The lowest energy level, or ground state, is like the
slow lane. This is where the electrons are normally found. When the electron
receives more energy from an outside source it moves to a higher energy level, excited
state. Similarly a car in the slow lane that increases its speed (increase energy) moves
to a faster lane. Both electron and the car will return to their ground state energy, or
speed as their energy decreases. When the electron absorbs energy from an external
electric source, it moves from one orbit to another. The electron is then pulled back by
the nucleus and loses the gained energy by emitting a photon with an energy equal to
the energy gained E = hx. This relationship is called Bohr frequency condition. Each
spectral line arises from a specific transition. The greater the energy loss, the higher
the frequency (and the shorter the wavelength) of the radiation emitted. For each
frequency a line is observed with a different color. The spectrum of hydrogen consists
of discrete lines, which indicates that the electron in the atom can exist only in series of
discrete states, called energy levels, n.
Quantum Model
 In classical mechanics we speak of the path of electrons called orbits. However,
because of the wavelike properties of electrons we cannot say that an electron will
be found at a certain point in an atom. This is also known as the Heinsenberg
uncertainty principle. Ernest Shrödinger, an Austrian scientist, devised an
equation that describes electrons in terms of quantum mechanics. This equation
lets us calculate the probability that an electron is at a particular point in space
around the nucleus, called orbital.
 In order for these electrons to remain around the nucleus without colliding, they
must occupy different orbitals with different sizes, shapes, and energy. Each
orbital holds a maximum of two electrons. In order for the pair of electrons to
remain together in the same orbital without repelling, the electrons must spin in
opposite directions.
Atomic Orbitals

Each atomic orbital corresponds to an energy level of the electron. The higher the
energy, the larger the orbital. The various shapes of atomic orbitals can be
classified into four main types, which are labeled s, p, d, and f. There are many
orbitals of each type. They differ principally in the size of the cloud, which is
related to the energy level.

The simplest way of drawing an atomic orbital is as a boundary surface, a surface
within which there is a high probability (typically 90%) of finding an electron.
o An s-orbital has a spherical boundary surface, because the electron cloud is
spherical. s-orbitals with higher energies have spherical boundary surfaces
of bigger diameter.
o A p-orbital is a cloud with two lobes on opposite sides of the nucleus. The
nucleus lies on the plane that divides the two lobes, and an electron will, in
fact never be found at the nucleus itself if it is in a p-orbital. There are
three p-orbitals of a given energy, and they lie along three perpendicular
axes.
o The boundary surface of a d-orbital is more complicated than that of an sor p-orbital. There are five d-orbitals of a given energy; four of them have
four lobes, one is slightly different and is oriented along an axis. In each
case an electron that occupies a d-orbital will not be found at the nucleus.
o The f-orbitals have more complicated shapes. There are seven f-orbitals of
a given energy. The shapes of the f-orbitals are rarely needed to explain
chemical properties and therefore will not be studied here.
Quantum Numbers and Atomic Orbitals

Shrödinger found that each atomic orbital is identified by three numbers called
quantum numbers. One quantum number is called the principal quantum number,
n; the other two are the azimuthal quantum number, l, and the magnetic quantum
number, ml. These quantum numbers have another job: as well as labeling the
orbital, they tell us about the properties of the electron that occupies a given
orbital.
o The principal quantum number, n, is an integer that labels the energy
levels around the nucleus, from n =1 (ground state) to n = 7. Because the
clouds representing the orbitals get bigger as n increases, the average
distance of an electron from the nucleus also increases as n increases. On
average, an electron is closest to the nucleus in the ground state (n=1).
The orbitals form a series of thick shells, sometimes like fuzzy layers of an
imaginary onion. Shells of higher n surround the inner shells of lower n.
o The second quantum number, the azimuthal quantum number, l, governs
the shape of the orbital. Each value of l corresponds to one of the orbital
shapes. Thus, l = 0 corresponds to s; l = 1, corresponds to p; l = 2
corresponds to d; l = 3 corresponds to f. All the orbitals with a given value
of the azimuthal quantum number, l, are said to belong to the same
subshell of a given shell.
o The third quantum number, the magnetic quantum number, ml, labels the
different orbitals of a given subshell. The allowed values of ml are ml = l,
l-1, l-2, …, -l. Ex for the subshell with l =1, which consist of the porbitals, ml can have the values ml = +1, 0, and –1, so there are three porbitals in the subshell. These orbitals are denoted px, py, and pz. Each
label corresponds to a possible orientation of the lobes of the orbitals. In
general, a subshell with quantum number l consists of 2l +1 individual
orbitals.
o Electron Spin. An electron behaves in some respect like a spinning sphere.
It rotates on its axis. This property is called spin. An electron has two spin
states, represented by the arrows  (for a clockwise rotation) and  (for a
counterclockwise rotation). These two spin states are distinguished by a
fourth quantum number, the spin magnetic quantum number, ms. This
quantum number can have only two values: +1/2 indicates an  electron
and –1/2 indicates a  electron.
An orbital is specified by three quantum numbers; orbitals are organized
into shells and subshells. The relations between shells, subshells, and orbitals are
summarized in the table below:
n, (level,
shell)
or (sublevel or subshell)
ml,
# of Orbitals
Total
#
electrons
of
1
1s
1
2
2
2s
1
2
2p
3
6
3s
1
2
3p
3
6
3d
5
10
4s
1
2
4p
3
6
4d
5
10
4f
7
14
3
4
Orbital Energies
As well as being attracted by the nucleus, each electron is repelled by all the
other electrons in the atoms. As a result, it is less tightly bound to the nucleus than it
would be if those electrons were absents. We say that each electron is shielded from
the full attraction of the nucleus by the other electrons in the atom. The shielding
effectively reduces the pull of the nucleus on an electron. The less shielding effect the
closer the electron is of the nucleus and the lower its energy level. An s-electron can
be found close to the nucleus and is said to penetrate through the inner shells. A pelectron penetrates much less. For this reason, s-electrons lie at a lower energy than
p-electrons of the same shell. The order is s < p< d< f. Orbitals in different subshells
occupy different energy levels due to the shielding effect. The order is as follows: 1s,
2s, 2p, 3s, 3p, etc...
The Electron Configuration or Building-up Principle of Atoms
We report the electronic structure of an atom by writing its electron configuration.
Electron configurations tells us in which orbitals the electrons for an element are
located which is a list of all its occupied orbitals with the number of electrons that
each contain. In order to write the electron configuration of an atom we need to
follows three rules:
8. No two electrons can fill one orbital with the same spin. This is known as
the Pauli exclusion principle (discovered and stated by the Austrian scientist
Wolfgang Pauli):
9. Fill orbitals starting with lowest n and moving upwards, also known as the
Aufbau principle. It gives the order in which atomic orbitals are occupied as
follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, etc…
10. For degenerate orbitals, (more than one orbital in a subshell) add electrons
with parallel spin to different orbitals of that subshell rather than pairing
two electrons in one of them. This is also called Hund’s Rule (proposed by
the German scientist Fritz Hund).
Writing electron Configuration

The electron configuration of an atom is a shorthand method of writing the
location of electrons by sublevel. The sublevel is written followed by a
superscript with the number of electrons in the sublevel. Example: If the 2p
sublevel contains 2 electrons, it is written 2p2.

In order to write the electron configuration:
o First, determine how many electrons are in the atom. Example iron has 26
electrons.
o Arrange the energy sublevels according to increasing energy: 1s 2s 2p 3s 3p
4s 3d …
o Fill each sublevel with electrons until you have used all the electrons in the
atom: Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d 6
o The sum of the superscripts equals the atomic number of iron (26)
Electron Configuration and the Periodic Table
The periodic table is an arrangement of the elements in order of their atomic
numbers so that elements with similar properties fall in the same column, or group.
The periodic table can be used as a guide for electron configurations. Based on the
electron configuration of the elements, the periodic table can be divided into four
blocks, the s, p, d, and f blocks.

Groups 1A and 2A have the s-orbital filled. They form the s-block.

Groups 3A - 8A have the p-orbital filled. They form the p-block.

Groups 3B - 2B have the d-orbital filled. They form the d-block.

The lanthanides and actinides have the f-orbital filled. They form the f-block.
The s-block elements: Groups 1 and 2.
a) The elements of group 1 of the periodic table are known as the alkali metals.
Their electron configuration ends with ns1, n being the period of the element
and 1 is the number of valence electrons (outermost shell). Alkali metals are
soft shinny, and highly reactive.
b) The elements of group 2 are known as the alkaline-earth metals. Elements of
group 2 contain a pair of electrons in their outermost s subshell and therefore
their group electronic configuration ends with ns2. Alkaline earth metals are
less shinny, harder and less reactive than group 1 elements.
The p-block: Groups 13-18.
a) Elements in the p-block and in period 1 and 2, have an electron configuration
that ends with ns2np1-6.
b) For periods 3-7, the electron configuration ends with ns2(n-1)d1-10np1-6.
c) The properties of the elements in the p-block vary greatly, as the block
includes metals, non-metals, and metalloids.
d) The elements of group 17 are known as halogens. The elements of group 18
are known as noble gases.
e) The p-block elements together with the s-block elements are called the maingroup elements or representative elements.
The d-block elements: Groups 3-12.
They have a configuration that ends with (n-1)d1-10ns2. The d-block elements are
metals with typical metallic properties and are often referred to as transition
metals.
The f-block elements: Lanthanides and Actinides.
They have a configuration that ends with (n-2)f1-14(n-1)d1-10ns2.
a) There are 14 f-block elements between Lanthanum, La, and Hafnium, Hf, in
the sixth period called lanthanides (or rare earth elements). They are shiny
metals similar in reactivity to the group 2 alkaline-earth metals.
b) There are also 14 f-block elements, the actinides, between actinium, Ac, and
element 104, Unq, in the seventh period. They are all radioactive. The first
four are found naturally on earth. The remaining actinides are laboratorymade elements.
Condensed Electron Configurations
 The electron configuration for Na is: 1s2 2s2 2p6 3s1. We can abbreviate the
electron configuration by indicating the innermost electrons with the symbol of the
preceding noble gas.
 The preceding noble gas with an atomic number less than sodium is neon, Ne. We
rewrite the electron configuration: Na: [Ne] 3s1.
o [Ne] represents the electron configuration of neon. It also represents the
core electrons: electrons in [Noble Gas].
o And 3s1 means that there is only one electron in the 3s orbital, the outershell: electrons outside of [Noble Gas].
Orbital diagram for an atom

An orbital diagram is the representation of the electrons in each orbital.

In order to write the orbital diagram for an element, we need to first write the
electron configuration. We then represent each orbital with a box and each
electron in the orbital with an up or down arrow.

An orbital diagram helps us determine the number of paired and unpaired
electrons in each orbital. It also helps us determine the number of valence
electrons and write the electron dot notation of the atom.
Valence electrons
 Valence electrons are the outermost electrons and are involved in chemical reactions.
These electrons are of the highest energy and are furthest away from the nucleus. For
the representative elements, the group number represents the number of valence
electrons.
Electron Dot Notation
 An electron dot formula of an element shows the symbol of the element surrounded by
its valence electrons. We use one dot for each valence electron.

.
Example: H, (one valence electron) H , and Al (3 valence electrons)
.Al:
Ionic Charge
 Recall, that atoms lose or gain electrons to form ions.
 The charge of an ion is related to the number of valence electrons on the atom. Atoms
with fewer valence electrons (1-4) tend to lose electrons to achieve a noble gas


configuration. Whereas atoms with more valence electrons (5-7) tend to gain electrons
to achieve a noble gas configuration.
Metals lose their valence electrons to form cations.
Nonmetals gain electrons to form anions.
Predicting ionic charges
 Group IA/1 metals form 1+ ions, group IIA/2 metals form 2+ ions, group IIIA/13
metals form 3+ ions, and group IVA/14 metals from 4+ ions.
 Group VA/15 elements form -3 ions, group VIA/16 elements form -2 ions, and group
VIIA/17 elements form -1 ions.
THE PERIODIC TABLE
The periodic table is divided into boxes containing symbols of elements, atomic numbers,
and mass numbers. The periodic table is divided in 18 groups (columns) and 7 periods (rows).
Going from left to right, both atomic and mass numbers increase. The atomic number increases by
one unit, and the mass number increases by one or more units.
The s-block Elements
It is the first two columns of the periodic table (Groups 1A and 2A). Group 1A is
called the alkali metal family, and group 2A is called the alkaline-earth metals.
a) Physical properties
 They are good conductor of electricity.
 They are ductile, malleable, and lustrous.
 They are solids with high melting points and good conductors of heat.
b) Chemical Properties
 They form cations. A group 1 element is likely to form +1 ions, such as Li +,
Na+, and K+. Group 2 elements similarly form +2 ions, such as Mg2+, Ca2+,
and Ba2+.
 They react with the oxygen in air to form basic oxides, such as Na2O, which
react with acids.
 Because ionization energies are lowest at the bottom of each group and the
elements there lose their valence electrons most easily, the heavy elements
cesium and barium react most vigorously of all group1 and 2 elements,
respectively.
The p-block Elements
There are three types of elements in the p-block: nonmetals, metalloids and metals.
1. Non-metals are on the right side of the zigzag line in the p-Block. It includes group 17, 18
A and C, N, O, P, S, and Se. The nonmetals in group 18A are called noble gases. They
exist in nature as monatomic substances (single atom). They are gases and do not react
with other elements. The elements in group 17 or 7A are called halogens. Some non metal
elements exist in nature as monatomic such as the noble gases, other exist in nature as
diatomic such as O2, F2, Cl2. N2.
Chemical properties
 To the exception of noble gases (group 18), the p-block nonmetal elements tend to
gain electrons to complete their valence shell. They have high electron affinity.
 They react with metals to form ionic compounds with ionic bonds.
 They react with each other to form covalent bonds.
 They do not react with acids.
 They form acidic oxides such as Al2O3, which react with bases.
Physical properties
 Nonmetals in the p-block are poor conductors of electricity.
 They are not ductile, and not malleable.
 They can be solid, liquid, or gas.
 They have low melting points.
 They are poor conductors of heat.
2.
Metalloids or semi metals are the elements on both sides of the zigzag line with the
exception of aluminum and polonium. Metalloids have physical and chemical
properties between metals and nonmetals.
 They fall in the diagonal band across the periodic table, between the metals, and the
nonmetals.
 Their have a diagonal relationship. A diagonal relationship is a similarity in
chemical properties between diagonal neighbors. For example Si and As, As and Te,
Ge and Sb, Sb and Po are pairs with diagonal relationships.
 The elements on the metal side have physical and chemical properties similar to
metals. Those that are on the nonmetals have properties similar to nonmetals.
3. Metals. All the other elements in the p-block are metals. They have similar
chemical and physical properties as the metals in the s- and d-blocks.
The d-block Elements

The elements in the d-block are called transition metals. All d-block elements are metals
with properties between those of s-block and p-block metals. Many d-block elements
form cations with different oxidation numbers.
The f-Block
 The elements in the f-block are called the inner transition elements. They are also
referred to as the rare earth elements. They are considered rare because they are
radioactive and unstable because of their large size. They have similar properties as the
transition metals (d-block). The first row of the f-block is called the lanthanides; the
second row of the f-block is called the lanthanides.
PERIODIC TRENDS
Atomic Radius
It is defined as half the distance between the centers of neighboring atoms. There are two
factors that affect the atomic radius: the pull effect of the nucleus and the shielding effect
of the inner shell electrons.


Within a period, all atoms have the same number of shielding electron (inner shells
electrons), but have different number of valence electrons, and therefore have same
shielding effect but different pull effect. The more protons in the nucleus, the stronger
the pull effect, and therefore the smaller the atom. The number of protons increases
from left to right within a period, and therefore, the pull effect increases. The Radii
decreases from left to right within a period.
Down a group, atoms have increasing shielding effect and therefore less pull from the
nucleus, which increases the atomic size. The Radii increases down a group.
Ionic Radius

Metals lose electrons to have their valence shell filled, and become cations. In cations
there are more protons than electrons, which produces a strong pull effect from the
protons in the nucleus  Cations are smaller than their parent atoms. The higher the
charge of the cation, the stronger the pull, and the smaller the ion

A > A+ > A2+ > A3+

Non-metals gain electrons to fill their valence shell, and become anions. Anions have
more electrons than protons, which create less pull effect from the protons in the
nucleus  Anions are larger than their parent atoms. The higher the charge of the
anion, the weaker the pull, and the larger the ion.

A3- > A2- > A- > A

Ionic radii increase down a group (increasing shielding effect) and decrease across a
period (stronger pull).
Ionization Energy
The ionization energy is the energy needed to remove an electron from an atom in the gas
phase.





To remove an electron from a neutral atom, energy has to be supplied to drag the
electron away from the attraction of the nucleus. The stronger the pull in the neutral
atom, the harder it gets to remove the electron.  Larger atoms are easier to ionize
than smaller atoms.
For the first ionization energy, I1, we start with the neutral atom:
Cu (g)  Cu+ (g) + e- (g) Energy required = I1 (745.1 kJ. mol-1)
Because larger atoms are easier to ionize than smaller atoms, the first ionization
energies of the main group elements, is highest for elements close to helium and is
lowest for elements close to cesium.
The second ionization energy, I2, of an element is the energy needed to remove an
electron from a singly charged gas-phase cation:
Cu+ (g)  Cu2+ (g) + e- (g) Energy required = I2 (1955 kJ. mol-1).
Much more energy has to be supplied to drag a second electron away from a cation,
because now the nucleus has more protons, and exerts more pull. Second ionization
energies are higher than the first ionization energies (of the same element), and very
much higher if the electron is to be expelled from a filled shell.