Strcuture 1.3 Electron configurations by Anoosha Qaisar

IB chemistry notes by Anoosha Qaisar

1. Lecture notes SL/HL
By
Ms. Anoosha Qaisar

2. Electron
configurations
Guiding questions: How can we model the energy states of electrons in atoms?

3. Emission Spectra
Learning Objectives (S1.3.1)
Understand
Emission spectra are produced when photons are emitted from
atoms as electrons in excited states return to lower energy
levels.
Apply your knowledge to:
Describe qualitatively the relationship between colour,
wavelength, frequency and energy across the electromagnetic
spectrum.
Distinguish between a continuous spectrum and a line spectrum.

4. Demonstration “Flame TEST”

5. CONTINUOUS SPECTRUM
In the 1600s, Sir Isaac Newton showed
that sunlight can be broken down
into different coloured components
using a prism. This generates a
continuous spectrum. This type of
spectrum contains light of all
wavelengths, and appears as a
continuous series of colours, in which
each colour merges into the next, and
no gaps are visible. The classic
example of a continuous spectrum is
the rainbow. The wavelength of
visible light ranges from 400nm to
700nm

6. A continuous spectrum is produced when white light passed
through a prism.
It shows all colors in an unbroken sequence of frequencies,
such as the spectrum of visible light.

7. ●A line spectrum is an emission spectrum that has
sharp lines produced by specific frequencies of
light.
●It is produced by excited atoms and ions as they
fall back to a lower energy level.
●Different elements have different line spectra so
they can be used to identify unknown elements.
LINE SPECTRUM

8. A pure gaseous element subjected to a high voltage under reduced
pressure will glow — in other words, it will emit light. When this light
passes through a prism, it produces a series of lines against a dark
background. This is known as an emission spectrum
In contrast, when a cold gas is placed between the prism and a
source of visible light of all wavelengths, a series of dark lines
within a continuous spectrum will appear. This is known as an
absorption spectrum.
Emission and Absorption Spectrum

10. Why the Helium element is called Helium?
Discovery of Helium (He):
●Discovered in 1868 by French astronomer
Pierre Jules César Janssen
●Observed during a solar eclipse through a
prism
●Detected by a bright yellow line from the
Sun's chromosphere
●Named after the Greek word 'helios' meaning
Sun
Spectroscopy:
The study of the interaction between matter and
light
Elements can be identified by their unique emission
spectra.Each element’s spectrum acts like a
fingerprint

11. The electromagnetic spectrum
●Electromagnetic radiation comes in different forms.
●All forms travel at the same speed of light but have
different wavelengths.
●The higher energy forms have shorter wavelengths and
higher frequencies.

12. The frequency of a wave is inversely proportional to its
wavelength.
That means that waves with a high frequency have a short
wavelength, while waves with a low frequency have a longer
wavelength.
λ = V/f
(where ' V' is the speed of the wave and 'f' is the frequency
of the wave)It has units of distance (m)

13. The wavelength of a wave describes how long the wave
is. The distance from the "crest" (top) of one wave to the
crest of the next wave is the wavelength. Alternately, we
can measure from the "trough" (bottom) of one wave to
the trough of the next wave and get the same value for
the wavelength.

14. A photon is a quantum of energy, which is proportional to the frequency of
the radiation as follows:
E = h × f Where
E = the specific energy possessed by the photon, expressed in joules, J
h = Planck’s constant, 6.63×10–34 Js
f = frequency of the radiation.The frequency, f, is the number of waves that
pass a point in one second. It has the units hertz (Hz) or s–1.

All regions of the EM spectrum travel at the same speed – the speed of
light, which is 3 × 108 m s–1

15. Order of the EM spectrum

16. Speed of light

17. Relationship between energy, frequency and wavelength

18. Figure shows the visible region in more detail, which is known as the
continuous spectrum. Our eyes see the continuous spectrum as white light,
but as you can see it is actually made up of different colours.In the visible
region, red light has the lowest energy, lowest frequency and longest
wavelength. As we go across the spectrum, the energy and frequency both
increase and the wavelength decreases. So violet light has the highest
energy, highest frequency and shortest wavelength.

20. The line emission spectrum of hydrogen
Learning Objectives (S1.3.2 and 1.3.3)
Understand:
The line emission spectrum of hydrogen provides evidence for the
existence of electrons in discrete energy levels, which converge at higher
energies.
Apply your knowledge to:
●Describe the emission spectrum of the hydrogen atom, including the
relationships between the lines and energy transitions to the first,
second and third energy levels.

●Deduce the maximum number of electrons that can occupy each
energy level.

21. Build your own spectroscope (kids crafting activity) | Cambridge Festival 2021
Emission line Spectra

22. Differences between the continuous spectrum and the
emission line spectrum


●A continuous spectrum shows all the wavelengths or frequencies of
visible light
●An emission line spectrum only shows specific wavelengths or
frequencies of light.

24. Emission line spacing: Lines get closer together at the
high-energy end of the spectrum.

Convergence: Lines converge toward the violet/blue end (high
frequency, short wavelength).

Energy and wavelength relationship: Shorter wavelengths
correspond to higher energy (violet/blue), while longer wavelengths
correspond to lower energy (red).

Observation in spectrum: Distance between the red and blue
lines is larger than between the blue and violet lines.

25. A selection of emission
line spectra for certain
elements – can you
identify the unknown gas?

26. The main energy
levels occupied by
electrons, assigned
the letter n.


n = 1 has the lowest
energy and n = 6 the
highest
Ground state
Excited states
Formation of
emission line
spectrum
As the value of n
increases, so
does the distance
from the nucleus
and its energy
also increases
Energy
absorbing
Energy
emitting

27. Bohr’s Model of the Hydrogen Atom
Hydrogen's Abundance:
Most abundant element (90% of atoms in the universe)
Located in Group 1 of the periodic table (alkali metals) despite not being a
metal .Hydrogen has only one electron
Bohr’s Planetary Model (1913):
Electrons orbit the nucleus like planets around the Sun
Only specific orbits (shells) are allowed; electrons can’t exist between them
Electrons move between shells by absorbing or emitting energy
Significance:
Bohr’s model explains hydrogen’s emission line spectrum
Model only works for single-electron atoms (e.g., hydrogen), not
multi-electron atoms

28. The hydrogen
emission
spectrum
Electron
Transitions:Electrons
transition from higher to lower
energy levels, emitting light
Specific transitions produce
distinct lines in the spectrum
Example: n = 3 to n = 2
transition produces the red
line (visible light)

29. The full hydrogen emission spectrum
Visible light spectrum
shows transitions from
higher energy level to n =
2 (visible region)
Transitions from higher
energy level to n = 1
correspond to UV
radiation (high energy)
Transitions to n = 3
correspond to infrared
radiation (low energy)
Energy levels converge
at high energies
n = ∞ represents
complete ionization
(electron fully removed)

30. How is possible that we see 4 lines in the
emission spectra of the Hydrogen even If it
has just 1 electron?

31. Answer:

Even though hydrogen has only one electron, we see four distinct lines in its
emission spectrum because the electron can transition between different energy
levels. When hydrogen gas is energized, the single electron can be excited to
various higher energy levels (n = 3, 4, 5, etc.). As the electron falls back to lower
energy levels (particularly to n = 2 in the visible region), it emits energy in the form
of light, creating different lines. Each transition from a higher level to n = 2
corresponds to a different wavelength of visible light, resulting in the four lines
observed in the spectrum.
So, the four lines represent different possible transitions, not four electrons.

32. Electrons and Energy Level
The main energy level or shell is given an integer number, n, and can hold
a maximum of electrons, 2n
2
where n is the principal energy level number.


A single atomic orbital can hold a maximum of two electrons

33. Quantum mechanical model
The scientists Heisenberg, de Broglie and Schrodinger developed the
current model of the atom called the Quantum Mechanical Model.

The electrons do not travel in precise orbits, but in wave functions called
Orbitals.

HEISENBERG UNCERTAINTY PRINCIPLE:
We are limited in just how precisely we can know both the position and
momentum of a particle at a given time.
The wave function or orbital has a 90% probability of finding the electron
within it.

36. Main energy levels and sublevels

Atomic orbitals
A region of
space where
there is a high
probability of
finding an
electron.

39. Electron configurations
Electronic configuration shows the arrangement of electrons in their
different levels around the nucleus of an atom.
s < p < d < f

40. Notation of the electron configuration
Electron configuration simulation

41. 1.The Aufbau principle
When adding electrons to an atom, the lower energy orbitals must be filled first.

Degenerate orbitals:
Atomic orbitals
that have equal
energy levels. For
example, the three
3p orbitals are
degenerate orbitals.



There is an overlap
in energy between
the 3d and 4s
sublevels. The 4s
sublevel is of lower
energy and fills
before the 3d
sublevel.

42. Notes about sublevels
Degenerate orbitals: Atomic
orbitals that have equal energy
levels. For example, the three 3p
orbitals are degenerate orbitals.

There is an overlap in energy
between the 3d and 4s sublevels.
The 4s sublevel is of lower energy
and fills before the 3d sublevel.

43. 2. The Pauli exclusion principle
An atomic orbital can only hold two electrons and they must have opposite spins.

44. 3. Hund’s rule
When we have degenerate orbitals (orbitals of the same energy) then each orbital is
filled with a single electron before being doubly occupied.

46. Trends in ionization energy
Energy level n= ∞. This
represents the point
when the electron has
been completely
removed from the
attraction of the nucleus
and the atom has been
ionised.

47. 1)Write the full electron configurations for the first ten elements of the
periodic table, from hydrogen to neon.


NOTE!!! add the correct number of electrons (according to the charge) for
negative ions or to remove the correct number of electrons for positive ions.


2) Write the full electron configurations for the following ions.
(a) Na+
(b) Mg2+
(c) F–
(d) O2–

48. Condensed electron configurations
This notation uses the symbol of a noble gas to represent the core electrons.

49. Exceptions to the Aufbau principle

These are for the elements copper (symbol Cu) and chromium (symbol
Cr). We expect the electron configurations to be as follows:
Cr = [Ar] 4s2 3d4
Cu = [Ar] 4s2 3d9


The actual electron configurations for copper and chromium are:
Cr = [Ar] 4s1 3d5
Cu = [Ar] 4s1 3d10

51. 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p

52. To deduce electron configurations,
follow the order of orbital filling: 1s, 2s,
2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f,
5d, 6p, 7s, 5f, 6d, 7p.

53. (a) Carbon, C
(b) Phosphorus, P
(c) Potassium, K
(d) Cobalt, Co
(e) Bromine, Br
Write the condensed electron configurations for the following atoms: