UNIT 2:ELECTRONIC CONFIGURATION OF ATOMS AND IONS
UNIT 2: ELECTRONIC CONFIGURATION OF ATOMS
AND IONS
Key unit Competence
To relate Bohr’s atomic model with atomic spectrum of Hydrogen, write electronic
configuration of atoms and ions using s, p, d and f atomic orbitals and interpret
graphical information related to ionization energy of elements.
Learning objectives
By the end of this unit, students should be able to:
• Explain the stability of atoms using the concept of quantization of energy.
• Explain the achievements and limitations of Bohr’s atomic model.
• Explain the existence of energy levels using the data from emission spectra.
• Describe hydrogen spectral lines and spectral line series
• Explain the types of spectra in relation with the nature of light
• Explain the quantum theory of the atom using the quantum numbers.
• Determine the number and shapes of orbitals in a given energy level or
principal quantum number
• Explain the rules governing the electronic configuration: Aufbau principle
and Hund’s rule
• Explain the relationship between the electronic configuration and the
stability of the atoms
• Interpret the graphs of first ionisation energy against the atomic number.• Describe the factors which influence the first ionisation energy.
1. What can you see on the image above?
2. What type of motion is performed by the people on the image?
3. How does their potential energy change?
2.1. Bohr’s atomic model and concept of energy levels
The potential energy of a person walking up ramp increases in a uniform and
continuous manner whereas the potential energy of a person walking up steps
increases in a stepwise and quantized manner. This can be explained by the values of
energy which are continuous for the person walking up ramp while they are discrete
(discontinued) for the person walking up steps (Figure 2.1(a) and Figure 2.1(b).
Niels Bohr (1885-1962) a young Danish physicist working in Rutherford’s laboratory,
suggested a model for the hydrogen atom and predicted the existence of line
spectra. In his model, based on Planck’s and Einstein’s ideas about quantized energy,
Bohr proposed three postulates:
• An electron can rotate around the nucleus in certain fixed orbits of definite
energy without emission of any radiant energy. Such orbits are calledstationary orbits.
• An atom can make a transition from its stationary state of higher energy
E2
to a state of lower energy E1 and emit a single photon of frequency ν.
Conversely, an atom can absorb an energy at the lower level E1
and transit
to the higher energy level E2
. That is, the change in energy for a system,
,which can be represented by the equation:
where n is an integer (1,2,3,…) and h is Planck’s constant
determined from experiment and has a value of 6.626x10-34J.s; ⱱ is the frequency of
the electromagnetic radiation absorbed or emitted. Each of these small “packets”
of energy is called a photon or a quantum of energy. Energy can be gained or lost
only in whole-number multiples of the quantity hν ,
That is, the change in energy for a system can be represented by the equation:
, where n is an integer (1,2,3,…).
An electron does not release energy when it is in its stationary orbit. That is, the
electron does not change energy while it is moving around on a given orbit.
When an electron on a given orbit absorbs an appropriate quantum of energy, it
jumps, it is promoted to a higher energy level; this process is called “excitation” of
electron. On the contrary, if the electron loses an appropriate quantum of energy,
it falls on the lower energy level by emission of a light corresponding to the lost
quantum of energy and the process is called “de-excitation” of electron. As there
are many energy levels electrons can be excited to and de-excited from, an atom
will have many lines of absorption, each corresponding to a quantum of energy
absorbed: this appears as a series of lines called absorption spectrum. In the same
way the series of emission lines will produce an emission spectrum( see figure 2.3.a)
2.1.1. Achievements of Bohr’s Atomic Model
• Explanation of the stability of an atom
Based on Rutherford’s atomic model, the electrons move around the nucleus in
circular paths called orbits. According to the classical theory of electromagnetism,
a charged particle revolving around a charged nucleus would release energy and
end up by spiraling into the nucleus; thus the atoms would be unstable. The Bohr’s
atomic model makes an assumption of discrete orbit (allowed orbits) to explain
why an atom is stable; by doing so, Bohr introduces the concept of quantization of
energy.
• Explanation of the production of the absorption and emission spectra
The Bohr’s atomic model explains the origin of atomic absorption and emission
spectra.
2.1.2. Limitations of Bohr Model
1.Bohr’s theory fails to explain the origin of the spectral lines of multi-electron
atoms.
It only explains the origin of the spectrum of hydrogen-like species having only one
electron such as H, He+, Li2+, Be2+, ........
The model fails to explain the spectral lines of atoms or species with more than one
electron.
2. According to Bohr, the circular orbits in which electrons revolve are planar.
However, modern research has shown that an electron moves around the nucleus in
the three dimensional space.
3. Bohr’s theory fails to account for Zeeman Effect and Stark Effect. Zeeman Effect
is the splitting of the spectral lines into thinner and closely- spaced lines when an
excited atom is placed in a magnetic field. Stark Effect consists of the splitting of the
spectral lines into thinner and closely-spaced lines in presence of electric field.
4. Bohr’s theory is in contradiction with Heisenberg’s uncertainty principle. Bohr
assumes that the electron revolves around the nucleus in circular orbits at fixed
distance from the nucleus and with a fixed velocity. However, according to W.
Heisenberg, it is not possible to know simultaneously the accurate position and the
velocity of a very small moving particle such as an electron.
Checking up 2.1
1. Find out two more examples that you can use to illustrating the concept of
quantization.
2. Discuss the main weakness of Rutherford’s nuclear atom.2.2. Hydrogen spectrum and spectral lines
Activity 2.2
Bohr’s atomic model allows to explain the emission spectra of atoms. This happens
when excited electrons lose energy in form of electromagnetic radiation and fall tolower energy levels.
The wave-particle nature of the light
Light as a wave
The light is a wave-like phenomenon as shown in Figure 2.2.
It is characterized by its wavelengths, generally symbolized by the Greek letter
lambda, λ, and its frequency, represented by the Greek letter nu1
, ν.
As shown in the Figure below, the wavelength represents the distance between two
(Source: http://psychelic-information-theory.com/em_spectrum)successive summits/peaks (or two successive troughs).The frequency represents the
number of complete wavelengths made by the light per second, also called cycles
per second.
Visible light is composed by different visible lights with different λ and ν.
But all those lights have the same speed, the speed of light, which, in a vacuum, is
1 The letter gamma, γ, may also be used
When an electron is excited or de-excited, the energy absorbed or emitted
corresponds to the difference of energy, ΔE, between the final energy level of the
electron, E2, and the starting energy level of the electron, E1: E2 – E1 = ΔE = hν. ΔE is
positive when E2>E1, this is the case of absorption and excitation of electron; on the
other hand ΔE may be negative when E2<E1, in case of emission and de-excitation
of electron.
Figure 2.4 below shows the different series of emission spectra of hydrogen. As
you can see, the difference between those series is the final energy level where the
electron fall after de-excitation.
The series have been named according to the scientists who discovered them.
Ionization of an atom or loss of an electron corresponds to excitation of an electronto the level n=∞.
Figure 2.3.a: continious, absorption and emission spectra
Checking2.2
1. What is the meaning of infinity level in the hydrogen spectral lines?
2. Given a transition of an electron from n=5 to n=2. Calculate
i) energy
ii) Frequencyiii) Wavelength
2.3. Atomic spectra
Activity 2.3
Observe the picture above, discuss in groups and answer the following
questions.
a. What do you see on the above photo?
b. State the physical phenomenon which is related to the above photo.
c. Think of any other means of producing the same pattern. List two of them.
d. What property can you attribute to light with reference to the above process?
The atomic spectrum is the range of characteristic frequencies or electromagnetic
radiations that are readily absorbed or emitted by an atom. It is also known as a line
spectrum.
When white light is passed through a prism, we see a myriad of colors – specifically
what we call a rainbow. This dispersion of white light demonstrates that white light
contains all the wavelengths of radiations and is thus considered to be continuous.
Each color blends into the next with no discontinuity and we get continuous or light
spectrum.
When elements are vaporized and then thermally excited, they emit light. However,
this light is not in the form of a continuous spectrum as observed with white light.
Instead, a discrete line spectrum is seen. A series of fine lines of different colors
separated by large black spaces is observed. The wavelengths of those lines are
characteristic of the element producing them – thus, elements can be identified
based on the spectral line data that they produce.
Two types of atomic spectra are known: Emission and absorption spectra.
1. Emission is the ability of a substance to give off light when it interacts with
heat.
2. Absorption is the opposite of emission, where energy, light or radiation is
absorbed by the electrons of a matter.
NB: A combination of the emission and absorption spectra of a given atom gives acontinuous spectra.
Checking up 2.3
1. Different metals, when exposed to a flame, emit different flame colors.
Explain the origin of that difference.
2. Would you expect to see the emission of lines and the absorption lines
of a given element to appear at the same place on a photographic plate
or not. Explain your answer.3. How do you explain the many spectral lines for the same element?
2.4 Orbitals and Quantum Numbers
Activity 2.4
1. a)Write the electronic configuration of aluminium atom(Z=13)
b) Indicate the number of electrons in each energy level/quantum shell
c) The shells are numbered from inside-outward starting from 1, 2, 3, 4 …
which other name is given to these shells?
d) How did you obtain the exact number of electrons in each energy level/quantum shell in (c) above?
To answer the questions that could not be answered by Bohr’s atomic model, other
atomic models were proposed. One of those models is the Quantum model that
has been developed by the Austrarian physicist Erwin Schrödinger (1887-1961).
The model is based on a mathematical equation called Schrödinger equation. This
model is based on the following assumptions or hypotheses:
• An electron moves around the nucleus continuously. However, it is not
possible to determine its precise position and velocity at the same time. We
can only determine the region, around the nucleus, where there is a high
probability of finding that electron at a given time.
• The region where the probability of finding electron is high, at more than
95%, is called “orbital”; in other words, the orbital is the volume or the
space (tridimensional) around the nucleus where there is a high probability
of finding the electron.
Without going into the mathematical development of the Schrödinger equation, we
can say that the energy of the electron depends on the orbital where it is located.
And an atomic orbital is described by a certain number of “quantum numbers”
according to the solution of Schrodinger equation, i.e. 3 whole numbers:
1) The principal quantum numberis a positive integer which varies from 1
to ∞. The principal quantum number indicates the energy level in an atom where
electrons can be located: the higher the n value, the higher the energy level. An
electron in energy level n=1 has lowest energy in an atom. The principal quantum
number, n, has been traditionally given names with the letters: K(n=1), L(n=2),M(n=3), N(n=4), O(n=5), P(n=6)
2)The angular momentum quantum number (l)
The second quantum number is the angular momentum quantum number
represented by the letter, l: it is an integer which can take any value from zero or
higher but less than n-1, i.e. equal to: 0,1, 2, 3,….up to n-1. For example if n= 1, l
is equal to 0, if n= 2, l can be 0, 1. It is also called secondary or azimuthal quantum
number. It indicates the shape of the orbital and is sometimes called the orbital
shape quantum number. By tradition, those different shapes of orbitals have been
given names or letter symbols: l = 0 = s, l =1 = p, l = 2 = d, l=3 = f
3) Magnetic quantum number (ml)
The magnetic quantum number describes the orientation of the orbital. It is an
integer that varies from -l to +l. For example if: l = 0, ml
can only be 0; if l = 1, ml =
-1, 0, +1; if l=2, ml
= -2, -1, 0, 1, 2. As you can see for each value of l there are (2l+ 1)
values of ml corresponding to (2l + 1) orientations under the influence of magnetic
field. The s orbital where l is zero and ml has no orientation; it has the shape of asphere as shown in figure 2.4
4) The spin quantum number (S)
The fourth quantum number is the spin quantum number, represented by the
symbol S (or ms
in some books). The electron behaves as a spinning magnet.The
spin quantum number is the property of the electron, not the orbital.
This number describes the spinning direction of the electron in a magnetic field.
The direction could be either clockwise or counterclockwise. The electron behaves
as if it were spinning about its axis, thereby generating a magnetic field whose
direction depends on the direction of the spin. The two directions for the magnetic
field correspond to the two possible values for the spin quantum number, S (ms).Only two values are possible: s = +1/2 and -1/2 as shown in the Figure 2.7 below.
Orbital box representation
An orbital box representation consists of a box for each orbital in a given energy
level, grouped by sublevel, with an arrow indicating an electron and its spin.
Note that two electrons in the same orbital have necessarily opposite spins as
indicated in the examples below.
The table 2.4 shows the electronic configuration of some elements using orbital box
representation and applying Hund’s rule.
N.B: An orbital box representation doesn’t show the real form of the orbital; the forms of thedifferent orbitals are shown in Figures 2.4, 2.5 and 2.6.
Noble Gas Notation
All noble gases have completely filled subshells and can be used as a shorthand way
of writing electron configurations for subsequent atoms.
When using this method, the following steps are respected.
a. Identify the noble gas whose electronic configuration is included in that of
the concerned element.
b.Write the chemical symbol of the identified noble gas within square brackets.
We call this the noble gas core.
c. Add electrons beyond the noble gas core. Note that electrons that are added to the electronic level of the highest principal quantum number (theoutermost level or valence shell) are called valence electrons.
2.6. Relationship between ionization energy, energy levels
and factors influencing ionization energy
2.6.1. Concept of Ionization energy
The ionization energy is a measure of the energy needed for an atom, in gaseous
state, to lose an electron and become positive ion.
The first ionisation energy is the energy required to remove one mole of electrons
from one mole of atoms in their gaseous state. The example below shows how torepresent the successive ionization energies of an atom M.
b) To remove the second electron needs greater energy because this electron is closer
to the nucleus in a 2p orbital. There is a steady increase in energy required as electrons are removed from 2p and then 2s-orbitals.
c) The removal of the tenth and eleventh electrons requires much greater amounts
of energy, because these electrons are closer to the nucleus in 1s orbital.
2.6.3. Factors influencing the extent of ionization energy
The ionization energy is a physical property of elements that can be influenced by
some factors:
1) Size of atom
The atomic size is the distance between the nucleus and valence shell. As the
number of energy levels (shells) increases, the force of attraction between nucleus
and valence electron decreases. Therefore, the valence electrons are loosely held
by the nucleus and lower energy is required to remove them, i.e. ionization energy
decreases with an increase in atomic size and vice versa.This is what happens when
you go down a Group.
2) Nuclear charge
The nuclear charge is the total charge of all the protons in the nucleus. As the nuclear
charge increases, the force of attraction between nucleus and valence electrons
increases and hence makes it difficult to remove an electron from the valence
shell. The higher the nuclear charge, the higher the ionization energy. This is what
happens when you cross a period from left to right.
3) Screening effect or Shielding effect
The Screening effect or Shielding effect is due to the presence of inner electrons
which have a screening or shielding effect against the attraction of the nucleus
towards the outermost electrons. The electrons present in inner shells between
the nucleus and the valence shell reduce the attraction between nucleus and the
outermost electrons. This shielding effect increases with the increasing number
of inner electrons. A strong Shielding effect makes it easier to remove an external
electron and hence lowers the ionisation energy.
2.6.4. Importance of ionization energy in the determination of the
chemistry of an element
Ionization energy provides a basis to understand the chemistry of an element. The
following information is provided.
Determination of metallic or non- metallic character.
The I.E informs us how the atom will behave chemically: a low I.E indicates that the
element behaves as metal whereas a high I.E indicates that the element behaves as
non-metal.
The first ionization energies of metals are all nearly below 800kJ mol-1 while those of
non- metals are all generally above 800 kJ mol-1.
Down the group ionization energies decrease so that the elements became more
metallic. In groups 14 and 15 there is change from non metallic to metallic character.
Across a period from left to right 1st I.E. increases. The elements become less metallic
to non- metallic
2.7. End unit assessment
1. Which of the following is the correct representation of the ground-state
electron configuration of molybdenum? Explain what is wrong with eachof the others
12. Four possible electron configurations for a nitrogen atom are shown below, but only one represents the correct configuration for a nitrogen atom
in its ground state. Which one is the correct electron configuration? Which
configurations violate the Pauli Exclusion Principle? Which configurationsviolate Hund’s rule?
a) What factors determine the magnitude of the first ionization energy?
b. To which group does element W belong? Explain
c. Would you expect W to be a metal or a non-metal? Explain your answer.d. Write an equation representing the second ionization of W.