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Unit V: Atomic Theory
Vocabulary: atoms, ions, compounds, molecules, quantum theory, block structure of the Periodic Table,
sub-atomic, particles, electrons, protons, neutrons, isotopes, radioisotopes, radioactive decay, average
atomic mass vs. mass number, standard notation, plum pudding model, spectroscopy, electromagnetic
spectrum, absorption and emission spectra, line spectra vs. continuous spectra, cations and anions,
mass spectrometer, mass spectra, electron arrangement vs. electron configuration, ground state,
excited state, photon, quantization, Planck’s equation and Planck’s constant, ionization energy (first and
successive), Bohr model, valence electrons, sublevels of electrons (s,p,d,f), wave-particle duality,
Heisenberg’s Uncertainty Principle, atomic orbitals, Pauli Exclusion Principle, Aufbau Principle, Hund’s
third rule
1. Introduction
The idea that all matter is composed of atoms is one of the most important discoveries in human
history. Although Greek and Indian philosophers had the idea 2500 years ago, we have only had
experimental evidence since the 19th century. This unit is mostly theoretical, based on the following key
The everyday notion of Newtonian, classical physics does not apply at the sub-atomic level.
Particles do not follow fixed paths.
Light behaves as a wave and as a particle. All matter behaves as a particle and as a wave. The
wave-like properties are more significant the smaller the particle.
Electrons are not found in defined orbits but in regions of probability called orbitals.
2. Historical background
a) Dalton
John Dalton, in the 19th century, noticed that hydrogen and oxygen always combine in the same
proportions to make a new substance called a compound (in this case, water). Compounds have
very different properties than their component elements. Dalton proposed:
All matter is composed of tiny indivisible particles called atoms
Atoms cannot be created or destroyed
Atoms of the same element are alike in every way
Atoms of different elements are different
Atoms can combine together in small numbers to form molecules
*Note: the term “molecule” is now usually used for covalent compounds i.e. where e- are
shared between non-metals.
b) Thompson
J.J. Thompson, at the end of the 19th century, discovered that different metals produce a stream of
negatively charged particles when a high voltage is applied across two electrodes. This device is called a
cathode ray tube (CRT) and is the technology behind tube TVs and other devices. He called these
negative particles “corpuscles”, now called electrons.
The electrons were the same regardless of the metal used, so Thompson concluded they were found in
all atoms, and created the Plum Pudding model of the atom:
c) Rutherford
ACT – “Atomic Target Practice”
Rutherford’s Scattering Experiment a.k.a. “Gold Foil Experiment” was one of the most important
experiments in history. The key findings:
The mass of an atom is concentrated in a very small area (nucleus) and atoms are mostly empty
The nucleus has a positive charge
“It was quite the most incredible thing that has happened to me. It was if you had fired a (artillery) shell
at a piece of tissue paper and it came back to hit you.” – Ernest Rutherford (1871-1937)
3. Subatomic particles:
Relative mass
Relative charge
Atomic number: the # of protons in the nucleus.
Mass number: the # of protons plus the # of neutrons in an atom
Standard notation:
4. Isotopes
Isotopes: atoms of the same element with different mass numbers.
On the periodic table, the relative atomic masses (a.k.a. molar mass) are often given as decimal (e.g.
chlorine is 35.45) because that number is an average relative atomic mass for all of the naturally
occurring isotopes.
Scientists can measure the masses of atoms, and therefore find the percent abundance of isotopes,
using a mass spectrometer.
ACT – Bean Bag Isotopes
Radioisotope: an isotope that is radioactive. Having too few or too many neutrons can result in unstable
nuclei. These nuclei decay (called nuclear decay) into a more stable nucleus. Recall from Grade 10 the
types of decay: alpha, beta and gamma.
e.g. iodine-131, a product of nuclear fission in nuclear power plants, was found in Vancouver seaweed
shortly after the Japanese power plant disaster of 2011.
Radioisotopes can also be of great benefit when used safely.
e.g. cobalt-60 is used to kill cancer cells via gamma decay, commonly called “radiation therapy”
Other uses of radioisotopes:
iodine-131 and iodine-125 used to treat cancer
sterilizing surgical instruments
preserving food
fighting crime
detecting cracks in structural materials
5. Ions
When an atom loses or gains electrons it becomes an ion. Ions have charges because the number of
protons (positive charges) is no longer equal to the number of electrons (negative charges).
Cation: a positively charged ion.
Anion: a negatively charged ion
Ionic compound: a compound that results from the attraction of cations and anions to each other e.g.
6. The Mass Spectrometer
The mass spectrometer allows us to measure the masses of atoms. Therefore it also allows us to find the
relative abundance of isotopes for an element.
5 Basic Operations:
a) Vaporization – allows individual atoms of the element to be analyzed
b) Ionization – the atoms are bombarded with high-energy electrons which knock the electrons out
of the atoms, producing positively charged ions
c) Acceleration – the positive ions are attracted to negatively charged plates. The electric field
accelerates them and they pass through a hole in the plates
d) Deflection – the positive ions then get deflected by a magnetic field placed at right angles to
their path. The amount of deflection is proportional to the charge/mass ratio. Ions with a
smaller mass get deflected more than larger ions. Also, ions with a higher charge will be
deflected more because they are more affected by the magnetic field.
e) Detection – the ions are detected and a signal is sent to a recorder. The strength of the signal is
a measure of the number of ions with that particular charge mass ratio.
The mass spectrometer can give us the mass of individual atoms, which are all in the range of 10-24g to
10-22g. As these numbers are so small, it makes sense for us to use relative values. Scientists chose
carbon-12 as the standard, because it is common, easy to transport, and a solid. In other words, at atom
of carbon-12 was given a mass of 12, and every other atom is compared to that.
Results of a mass spectrometer are presented as a mass spectrum:
e.g. mass spectrum for copper. Calculate the average mass
e.g. Boron, used as a control for nuclear reactors, exists as 10B and 11B. Use your periodic table to find
the abundances of the two isotopes.
7. Electron arrangement
ACT – flame test
a) The electromagnetic spectrum
Electromagnetic radiation comes in different forms, from high-energy gamma rays to low-energy radio
All waves travel at the speed of
light (c) = 3x108 m s-1
Wavelength (λ)= the distance
between successive crests
Frequency (ƒ) = the number of
waves that pass a point in 1 s.
The higher the frequency, the
higher the energy.
Related by equation c=λƒ
ACT: Spectroscopy lab
continuous spectra – when white light passes through a prism and all the wavelengths appear ROYGBIV
emission line spectrum – when high voltage is applied to a gas, lines that correspond with distinct
wavelengths of light are produced. As each element has a different emission line spectrum, they are like
barcodes that identify the element.
b) Bohr models
In Grade 10 you learned to the rules for drawing Bohr models for the first 20 elements. Here you will
learn why we have those rules, and the limitations of the Bohr model, but first some background:
The photons (packets of light) of light that are emitted from an element to give a line spectrum come
from the energy released when an electron falls from its excited state to its ground state.
Bohr used the line spectrum of hydrogen to propose that because the light is found in discrete
wavelengths, the energy of the electrons is also quantized – meaning it exists at discrete energy levels
and not in between. (staircase analogy)
Ephoton = ΔEelectron
Max Planck related the energy to the frequency of the radiation using ΔE=hf
(‘h’ = Planck’s constant = 6.626068 × 10-34 m2 kg / s)
Hydrogen spectrum:
Transition to n=1 produces UV light
Transition to n=2 produces visible light
Transition to n=3 or higher produces infrared light
Ionization energy: the energy re quired to remove an electron from the ground state of each atom in
one mole of gaseous atoms, ions or molecules
First ionization energy: the minimum energy needed to remove one mole of electrons from one mole of
gaseous atoms in their ground state
Valence electrons: the outer electrons (assume ground-state)
Image from:
Key points:
c) Quantum Mechanics
Those “shells” of the Bohr model are referred to as the principal energy levels, or principal quantum
numbers. Experiments show there are sub-levels, referred to as s, p, d and f. These patterns only make
sense if we treat electrons as waves instead of particles.
Wave-particle duality:
Light can be described by its frequency (a wave property) or by the E of the individual “particles” (called
photons or quanta of light). Both properties are related by Planck’s equation E = hf
The same is true of electrons, or any matter, although the significance of the wave-like properties is
more significant for subatomic particles. Electrons fired from an “electron gun” through a slit make a
diffraction pattern, just like waves.
Heisenberg’s uncertainty principle: we cannot know where an electron is at any given moment. The best
we can hope for is a “probability picture” of where the electron might be.
Atomic orbitals: a region around a nucleus where there is a 90% chance of finding an electron. An
electron is characterized by its orbital and its spin
ACT – quantum leap lab
Principal Quantum Number, n
is roughly equivalent to the n of the energy levels in the Bohr atom.
has whole number values n = 1, 2, 3, 4,… and designates what are called the main energy levels in an
all orbitals with the same n value are said to be in the same shell.
as n increases, electrons are generally further from the nucleus and have higher total potential
Subshell or Secondary Quantum Number, l
divides the shells into smaller groups of orbitals called subshells
only certain values of l are allowed
n: l = 0, 1, 2, …, (n-1).
The total possible subshells in each level is equal to n
For n =
l = 0, 1, 2, 3, …, (n-1)
0, 1
(s, p)
0, 1, 2
(s, p, d)
0, 1, 2, 3
(s, p, d, f)
(subshell letters)
whereas the principal quantum number describes the energy and size of an orbital, the secondary
quantum number determines the shape of the orbital
energies of the subshells within the same main energy level increase from s < p < d < f
Shapes: you should know the shapes of the s and p orbitals. See p. 332
Magnetic Quantum Number, ml
(also known as Orbital Quantum Number)
splits the subshells into individual orbitals
describes how an orbital is oriented in space relative to other orbitals
restrictions on the value of ml can range from -l to +l
For l =
ml = -l,…-1, 0, 1…, +1)
0 (s)
(one s orbital)
1 (p)
+1, 0 -1
(three p orbitals)
2 (d)
+2, +1, 0, -1, -2
(five d orbitals)
3 (f)
+3, +2, +1, 0, -1, -2, -3
(seven f orbitals)
Spin Quantum Number, ms
electrons behave as tiny magnets and like a toy top can spin in one of two directions
electron spin has two possible orientations
ms = +1/2, -1/2
Pauli exclusion principle – states that no two electrons in the same atom may have identical values for
all four quantum numbers
Aufbau principle – electrons are placed into the lowest energy orbitals first.
e.g. Draw and write electron configurations for the first five elements
Hund’s Rule – when electrons are placed in a set of orbitals of equal energy, they are spread out as
much as possible to give as few paired electrons as possible (separate orbitals and spins in the same
8. Electron Configuration for Atoms
Note: Electron arrangement refers to the 2,8,8 etc of the Bohr models – the number of electrons in the
principal energy levels only. Electron configuration refers to the specific orbitals.
e.g. write the full electron configuration for:
Energy level diagram:
As we get further from the nucleus, the
energy levels get closer together. Notice that
the 4s orbital will fill before the 3d.
Block structure of the periodic table
The shape of our modern periodic table is explained by the electron configuration. If you are writing the
electron configuration for an element and finish with electrons in the p-orbitals, we call them “p-block
Use coloured pencils to identify the s, p, d, and f-block elements.
Understanding this will save you a lot of time when writing electron configurations!
Core notation: a shorthand notation for electron configuration. The “core” refers to the electron
configuration of the noble gas that precedes the element you are dealing with. Write the core in square
brackets, followed by the additional electrons.
e.g. Core notation for silicon: [Ne] 3s23p2
e.g. Core notation for iron:
e.g. Core notation for krypton:
The 4s and 3d orbitals are so close together, that it results in a more stable atom if an e- jumps from the
4s orbital to either:
a) half-fill the 3d subshell. 3d4  3d5
b) completely fill the 3d subshell. 3d9  3d10
e.g. chromium
e.g. copper
9. Electron configuration for Ions
a) Anions – when adding electrons, simply follow the order of orbital filling given above
e.g. oxide
e.g. chloride
b) Cations – when removing electrons, remove from the highest principal energy level (n) first. For
example, for transition metals, remove the 4s electrons before the 3d electrons.
e.g. Fe3+
e.g. S2-
10. Ionization energy
We have already learned about ionization energy as it relates to a single atom. Here you will compare
first ionization energy for different elements.
ACT – Graphing Ionization Energy for the first 20 elements.
a) Explain the first ionization energy trend as you move across a period?
b) What are the exceptions?
c) What is the first IE trend as you move down a family?
d) How does the trend and the exceptions provide evidence for the existence of sub-shells?