THE BOHR MODEL OF THE ATOM
Experimental Evidence
Theories are based on observations. What did Bohr observe that made it necessary for him to
develop a new model for the structure of the atom?
Bohr noticed that when white light was passed through a prism, the visible spectrum of light was
obtained. In contrast to this, Bohr saw that when light of a particular colour (like from the hydrogen
discharge tube) was passed through a prism, a line spectrum was obtained. A line spectrum contains
distinct lines of a particular colour.
To Bohr, the line spectra phenomenon showed that atoms could not emit energy continuously, but
only in very precise quantities (he described the energy emitted as quantized). Because the emitted
light was due to the movement of electrons, Bohr suggested that electrons could not move
continuously in the atom (as Rutherford had suggested) but only in precise steps. Bohr
hypothesized that electrons occupy specific energy
levels. When an atom is excited, such as during heating,
electrons can jump to higher levels. When the electrons
fall back to lower energy levels, precise quanta of energy
are released as a specific wavelengths (lines) of light.
The Bohr Model of the Atom
The atom consists of two main parts:
• The nucleus
• The electron cloud
The nucleus is the centre of the atom and contains the
protons and neutrons. The nucleus comprises most of
the mass of the atom but little of its volume. The electron
cloud surrounds the nucleus. The electron cloud contains
electrons. The electrons are arranged in levels around the
nucleus. There is a maximum capacity for electrons at
each of these levels.
Electron Capacity
The electrons are arranged in levels around the nucleus. There is a maximum capacity for
electrons at each of these levels.
Level (n) Electron Capacity (2n2)
1
2(1)2 = 2
2
2(2)2 = 8
3
2(3)2 = 18
4
2(4)2 = 32
5
2(5)2 = 50
6
2(6)2 = 72
7
2(7)2 = 98
The levels get closer together as you move further
and further from the nucleus. "n" is the outermost
level where electrons can be and still be
considered a part of the atom.
An electron is at infinity if the electron is removed
from the atom.
Why Do Spectral Lines Occur?
Ground State vs. Excited State
Electrons are normally found in the lowest available energy orbitals. This is known as the ground
state.
If for some reason, an electron has absorbed energy and moved to an orbital of higher energy, then
the electron is in an excited state.
The electrons in an excited state will eventually lose their extra energy and return to their ground
state. When they release their excess energy, they emit electromagnetic radiation. When this excited
electron returns to the second energy level in the atom, the electromagnetic radiation is in the form
of visible light of a particular colour.
Much of your chemistry will deal with the elements with atomic numbers from 1 to 20. Let's
examine where the electrons are in atoms of these elements.
We will have to remember the electron capacity at each level.
Lewis Symbols or Diagrams
Elemental properties and reactions are determined only by electrons in the outer energy levels.
Electrons in completely filled energy levels are ignored when considering properties. Simplified
Bohr diagrams which only consider electrons in outer energy levels are called Lewis Symbols.
A Lewis Symbol consists of the element symbol
surrounded by "dots" to represent the number of
electrons in the outer energy level as represented
by a Bohr Diagram. The number of electrons in
the outer energy level is correlated by simply
reading the Group number. Lewis symbols for
oxygen, fluorine, and sodium are given in the
diagram.
Lewis Symbols for the elements of the second
period. Correlate the number of dots with the
group number.
Valence Electrons
Electron(s) located in the outermost energy level of the atom are called the valence electrons.
Why do you think that the valence electrons are of particular interest to the chemist?
Look at your periodic table in which you wrote the location of all the electrons of each of the first
twenty elements. Can you remember a quick way to determine?

the number of valence electrons that an atom of an element has?

the level at which the valence electrons are found (the valence level)?
Lewis Structures of Atoms
Chemists draw Lewis structures to show the valence electrons in an atom.
To draw a Lewis structure you:
a. write the symbol for the element
b. locate the element on the periodic table and determine the number of valence electrons it has.
For elements in the principal groups IA to VIIIA the number of valence electrons equals the
group number.
c. Imagine that the symbol for the element has four sides, like sides of a square.
d. Place dots on each side of the square, one per side before doubling up.
e. Each side can have at most 2 dots.
f. When two dots are on the same side, place them as a pair.
When looking at the spectral lines of elements other than hydrogen, scientists wondered why more
than one spectral line of a particular colour occurred. For example: In the spectrum for lithium there
were 4 red lines, 1 orange line and 2 blue lines. If Bohr's model was accurate, then more than one
line of the same colour would not be possible.
(negative of spectral lines for printing purporses)
Record in your notebook the line spectrum of three elements having more than one line of a
particular colour. Many elements show this characteristic. Choose any 3 elements that do this.
Be sure to label your diagrams of the line spectra that you choose. Ensure that you clearly
record which element belongs to each of your recorded line spectra. (Use coloured pencils to
assist with spectral lines)
Element 1) _______________________
Element 2) _______________________
Problems-Explanations
Why would the line spectrum for an element contain more than one red line?
Bohr's model of the atom could not explain why there is more than one line of a particular colour in
the line spectrum of various elements. His model did explain why one line of a particular colour
occurred. Do you remember why?
Refresh your memory.
Bohr's model of the atom needed to be modified in order to explain the occurrences of more than
one line of a particular colour in the line spectrum of elements.
What was this explanation?
The Modified Model of the Atom
•
The atom still consisted of a nucleus and an electron cloud.
•
The electron cloud still contained levels on which electrons would be located.
•
The energy levels for electrons were subdivided into sublevels: s, p, d and f. The sublevels
were called orbitals.
•
The first energy level has only one sublevel or orbital, the 1s orbital
•
The second energy level has two types of sublevels or orbitals, the 2s and three 2p orbitals
(making 4 sublevels in total).
•
The third energy level has three types of sublevels or orbitals, the 3s, three 3p, and five 3d
orbitals.
•
The fourth energy level has four types of orbitals: the 4s, three 4p, five 4d and seven 4f
orbitals.
•
Each of the fifth, sixth, and seventh energy levels has similar sublevels or orbitals as the
fourth energy level.
•
Each orbital can hold 2 electrons at most.
•
The electrons within an orbital or sublevel have opposite spins.
•
For any orbitals having the same energy, one electron goes into each orbital before you
double up.
•
The order of electrons filling into sublevels is in order of lowest energy to highest energy.
•
The sublevels overlap as you move further from the nucleus.
•
The order of filling electrons is not as straightforward as it may seem. This is because the
energy levels start to overlap as you move further from the nucleus.
Orbital Representation
Orbital Representation is a
way for chemists to show
where the electrons in an
atom are located. To show the
orbital representation,
chemists use the atomic
orbital chart sometimes called
an Aufbau diagram. This
atomic orbital chart can be
designed using lines, squares
or circles to indicate each of
the orbitals (sublevels).
Explanation
More Dead Chemists!
In 1926, Wolfgang Pauli (1900-1958) determined the assignments of electrons in various orbitals.
Known as the Pauli exclusion principle, it formed the basis of the electronic structure in atoms.
What this principle states is that each atomic orbital holds at most two electrons. The ground-state
electronic structures are then obtained following Pauli's principle: by arranging the atomic orbitals
in order of increasing energy and filling them one electron at a time starting from the lowest-energy
level. This process of starting at the bottom and working up is commonly known as the Aufbau
principle, Aufbau being derived from German, meaning "building-up". A group of orbitals with the
same energy (such as the three different 2p orbital) make up a sub-shell, and a group of sub-shells
with similar energy compromise a shell.
Friedrich Hund (1896-1997) suggested that if more than one orbital in an orbit is available,
electrons will fill empty orbitals before pairing in one of them. This is known today as Hund's rule.
To summarize, the ground-state electron configuration is obtained simply by "filling-up" orbitals of
the lowest energy in accord with the Pauli exclusion principle and Hund's rule.
MOLECULE LEWIS DIAGRAMS
A Lewis diagram depicts a mmolecule using an element symbol to represent the nucleus and core
electrons of each atom. Valence electrons are represented by lines for electron pair bonds and dots
for unbonded electrons.
The following procedure can be followed to derive Lewis diagrams for most molecules.
1. Find the total number of electrons:
Tabulate the total number of outer energy level electrons for all atoms in the molecule. For each
atom, read the group number.
2. Draw a first tentative structure:
The element with the least number of atoms is usually the central element. Draw a tentative
molecular and electron arrangement attaching other atoms with single bonds as the first guess.
Single bonds represented with a line
represent 2 electrons
3. Add electrons as dots to get octets
around atoms:
When counting electrons for the octet
around an atom, count both electrons in a
bond for each atom and any lone pair
electrons. Hydrogen, of course, gets only 2
electrons.
4. Count the total number of electrons in
the final structure to see if the total agrees
with the number tabulated in step #1. If not,
then move a lone pair of electrons into a
double bond. Or add more lone pairs of
electrons.
5. Cycle through steps 3 and 4 several times until
you get it right by trial and error.
Sketch of the Periodic Table
1H
3Li
11Na
19K
2He
4Be
12Mg
20Ca
5B
13Al
6C
14Si
7N
8O
9F
10Ne
15P
16S
17Cl
18Ar
35Br
36Kr