Nihon Akbar
SCH3U- R.H King Academy
Wiki ISU
Day 11: Electron Configuration for Ions and Flame Tests
Electron configuration can be used to show the configuration of an ion. You can show the change in the
configuration of the ion by either adding electrons(for a negative ion) in the orbitals, or subtracting
them(for a positive ion).
It is recommended that you write the original electron configuration for the atom and then determine
what will happen to its ion.
Cations and Anions have different rules for figuring out the configuration of the atom's ion.
Cations: Metals
Cations are ions who loses an electron, or in other words, are positively charged.
The ion will have to lose electrons from its original configuration. So the exponents representing the
number of electrons on the orbitals will change. In some cases, the entire orbital may disappear, as the
only electron on it is taken away. This will help the atom become stable by completing its final orbital
and make it stable.
Easy way to remember is that the last number in the electron configuration for the ion will be less than
what the original amount is.
For example, if we take Na+(Sodium ion) and try to write its configuration, we will get:
Original: 1s22s22p63s1
Sodium ion: 1s22s22p6
Because sodium loses an electron when it reacts to form its ion, its configuration shows 1 less electron,
which also results in 1 less orbital. Therefore, 3s1 which is the last orbital on sodium, is gone as the
electron is given to which ever element sodium reacts with. This results in the ion configuration to have
2p6 as the last orbital.
Other examples of electron configuration for cations:
Calcium ionOriginal: 1s22s22p63s23p64s2
Calcium ion: 1s22s22p63s23p6
Anions: Non-Metals
Anions are ions who gains an electron, or in other words, are negatively charged.
The ion will have to gain an electron from its original configuration. Therefore, an additional electron
will be added. This additional electron(s) will then complete the final orbital of the element, and give it a
Noble Gas structure.
Easy way to remember this is that the last number in the electron configuration will be more than what
the original amount is.
For example, if we take an Cl-(Chlorine ion) and write its configuration, we will get this:
Original: 1s22s22p63s23p5
Chlorine ion: 1s22s22p63s23p6 or [Ar]
Nihon Akbar
SCH3U- R.H King Academy
Wiki ISU
Because chlorine gains an electron after a reaction to form its ion, its ion configuration has 1 extra
electron added to the last orbital, 3p. This is why the number of electrons change from 3p5 to 3p6. This
gives chlorine a full final orbital, making it stable and allowing it to have a noble gas structure of Argon
gas([Ne]3s23p6 ).
Other examples of electron configuration for anions:
Oxygen ionOriginal: 1s22s22p4
Oxygen ion: 1s22s22p6 or [Ne]
Isoelectronicity
Isoelectronic means "equal electronic". This means that there are elements or ions which have the same
number of electrons in their configuration. They also tend to have similar chemical properties.
Isoelectronic can also mean a ccharged atom and neutral atom having the same electron configuration.
For example,
Li+ & He: Both has an electron configuration of 1s2
P-3 & S-2 : Both has an electron configuration of 1s22s22p63s23p6
S-2 & Ar: Both has an electron configuration of 1s22s22p63s23p6
Electromagnetic Spectrum and Flame Test
Electromagnetic Spectrum
An electromagnetic spectrum can be used to figure out how 'excited' electrons are in an atom.
An electromagnetic spectrum shows a continuous spectrum of coloured light. These colours can be
illustrated, in order, by R O Y G B V (red, orange, yellow, green, blue & violet). Each colour has its own
frequency and wavelength.
High Frequency and Short Wavelength = Higher Energy
Low Frequency and Long wavelength = Lower Energy
Red has the lowest frequency and longest wavelength, and violet has the highest frequency and shortest
wavelength. Other colours increase in energy as the order goes from red to violet.
In terms of atoms and electrons, an electromagnetic spectrum can help us understand the energy
required by the atoms to jump up and down levels when excited.
When gaseous atoms are heated, energy is absorbed by electrons. This absorption of energy allows the
electrons to 'escape' from their ground state, and jump to a higher energy level. When this happens, the
atom is said to be 'excited'. These excited atoms are unstable and return back to their ground state,
causing them to release the absorbed energy in energy emission patterns. Using a spectroscope, we can
see the emissions as colours in the electromagnetic spectrum.
Nihon Akbar
SCH3U- R.H King Academy
Wiki ISU
Flame Tests
Flame tests prove that each element has its own, unique characteristic colour which shows when the
element is heated. In class, we did the flame tests with the following elements, and the colour they
produced is recorded. Pictures of the result is also added .
Compound
Calcium chloride
Cation
Ca+
Colour
Dark Orange
Potassium chloride
K+
Purple/Violet
Copper (II) chloride
Cu2+
Aqua Green
Sodium chloride
Na+
Yellow-ish
Strontium chloride
Sr2+
Orange/ Red or Crimson
Calcium Chloride
Sodium Chloride
Potassium Chloride
Copper(II) Chloride
Strontium Chloride
Nihon Akbar
SCH3U- R.H King Academy
Wiki ISU
These characteristic colours that show when an element is heated is taken advantage of in fireworks.
Fireworks use these elements and heat to form colours. For example, an aqua green coloured firework
which is launched uses Copper(II) Chloride to form that colour. A crimson firework would use Strontium
Chloride, as shown in the picture below.
Strontium is used to
make crimson coloured
fireworks.
More on flame tests and their uses in the textbook on Pg.439-441
Homework:
Pg.33 # 1, 2, 3, 4, 6
Pg.441 # 5, 6, 9
Websites that may be helpful:
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http://chemwiki.ucdavis.edu/Physical_Chemistry/Quantum_Mechanics/Quantum_Theory/Trapped_
Particles/Atoms/Electron_Configuration#Electronic_Configurations_of_Cations_and_Anions
http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Electronic_Configurations
http://chemtutor.com/struct.htm#con
http://en.wikipedia.org/wiki/Electron_configurations#Ionization_of_the_transition_metals
http://www.chemguide.co.uk/atoms/properties/ionstruct.html
http://www.iun.edu/~cpanhd/C101webnotes/chemical-bond/isoelectronic.html
http://chemwiki.ucdavis.edu/Physical_Chemistry/Spectroscopy/Fundamentals/Electromagnetic_Radi
ation