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Chapter 5 Electrons in Atoms
5.1
Models of the atom
Rutherford's atomic model could not explain the chemical properties of elements due to the lack of
explanation regarding the "electron cloud".
Electrons – determine the chemical properties of an atom ( i.e. is it highly or slightly reactive? or
unreactive?).
These properties cannot be understood without knowing where the electrons are, and how they are arranged
around the nucleus
The Bohr Model
In 1913, Neils Bohr, a student of Rutherford, refined Rutherford's atomic model in order to explain why
atoms of all elements give off light when heated.
He proposed a planetary model of the atom with the electrons orbiting around the nucleus in a specific
circular paths. Each electron has an energy level.
Each energy level of the electron can be thought of as rungs on a ladder. The energy levels closest to the
nucleus are like rungs of a ladder closest to the ground. The rungs at the top of the ladder are analogous to
the energy levels furthest from the nucleus.
People can go up a ladder by going from rung to rung the way electrons in an atom can go from energy level
to energy level.
A person can never stand between the rungs of a ladder and electrons must always be at a specific energy
level as well (and not between them). In order for a person to go from rung to rung, they must move a
specific distance. In order for electrons to move from energy level to energy level, they must gain or lose a
certain amount of energy. The amount of energy is called a quantum of energy.
quantum - amount of energy required to move an electron from one energy level to another energy level
The amount of energy an electron gains is not always the same. The energy levels in an atom are not equally
spaced like a normal ladder.
The Bohr model was successful in understanding hydrogen's electron, but failed to account for the energy
emitted by electrons in other atoms
The Quantum Mechanical Model - The quantum mechanical model determines the allowed energies an
electron can have and how likely it is to find the electron in various locations around the nucleus.
Electrons only have a probability of being in a certain location, the same way the exact location of a fast
moving propeller blade at any time cannot not be determined.
In the quantum mechanical model, the probability of finding an electron within a certain volume of space
surrounding the nucleus can be represented as a fuzzy cloud. The cloud is more dense where the probability
of finding the electron is high.
Atomic Orbitals- are regions of space in which there is a high probability of finding an electron.
(a) Principal Energy levels (sometimes called shells) -The energy levels of electrons are labeled
by principal quantum numbers (n), which are integers starting at n = 1, 2, 3, 4, and so on.
(b) Energy sublevels - For each principal energy level there may be several orbitals with different
shapes that have different energy levels. These various energy levels within the principal energy
level are called sublevels.
**Each energy sublevel corresponds to an orbital of a different shape, which describes
where the electron is likely to be found.
**Different atomic orbitals are denoted by letters. The s orbitals are spherical, and p
orbitals are dumbbell-shaped, and the d orbitals are mostly clover shaped.
5.2 Electron arrangement in Atoms
Electron Configurations (Shown in the bottom of each elements box on the periodic table)- the way that
electrons are arranged in various orbitals around the nuclei of atoms (electrons and protons always interact to
make the most stable arrangement possible, which is the configuration with the lowest amount of energy).
Ground state - when an atom's electrons are in the lowest possible energy state
Three main rules for finding/writing electron configurations
1. Aufbau Principle - electrons occupy the orbitals of lowest energy first. Each box below represents an
atomic orbital
The orbitals for any sublevel of a principal energy level are always of equal energy
The s sublevel always has the lowest energy in the principal energy level
Notice that the 4s orbital has a lower energy than the 3d orbital
2. Pauli Exclusion principle - only two electrons can occupy an orbital and the two electrons must have
opposite spins (this is called paired electrons)
3. Hund's Rule - states that electrons occupy orbitals of the same energy in a way that makes the number of
electrons with the same spin direction as large as possible. ( i.e. when filling orbitals, only one electron can
be placed in orbitals of equal energy until they are all filled, only then can they be paired up)
Ex. Consider the electron configuration of oxygen
It has 8 electrons. The orbital of lowest energy is 1s (only one atomic orbital). It can hold two electrons of
opposite spin. The first electron is placed in arrow up, while the next electron is placed in arrow down. This
leaves 6 remaining electrons that must fill higher energy electron orbitals. The next highest sublevel is 2s
which can hold two electrons with opposite spin as well. This leaves 4 remaining electrons which must fill
the next highest energy sublevel: 2p. The 2p sublevel contains 3 atomic orbitals that can hold a maximum of
six electrons, but we only have 4 electrons remaining to place in. Each orbital must receive 1 electron
pointing in the up direction before any electrons are paired. This leaves one electron remaining which must
be placed into the first orbital that already contains one electron in the opposite direction or spin.
To the right of the above orbital diagrams, the short hand method for writing electron configurations is
shown, which shows the number of electrons in the sublevel as a superscript. For instance, Na is shown as
1s22s22p63s1
There are exceptions to the orbital filling diagrams for specific elements - this occurs after vanadium (Z=23)
because the d orbital fills after 4s and will take an electron from the 4s orbital to exactly half fill or
completely fill the d orbital - this configuration is more stable
Incorrect
Correct
Ex. Cr 1s22s2 2p63s23p63d44s2 --------------> Cr 1s22s22p63s23p63d54s1
Ex. Cu 1s22s22p63s23p63d94s2 ---------------> Cu 1s22s2sp63s23p63d104s1
More exceptions occur in the higher energy levels due to the energy decrease between principal energy
levels difference at higher principal quantum numbers (n).
Section Review Questions
1. Write the electron configuration and the expanded electron configuration of chlorine.
2. Write the electron configuration and the expanded electron configuration of titanium.
3. Write the electron configuration and the expanded electron configuration of bromine.
5.3 Physics and Quantum Mechanical Model
Light is electromagnetic energy and is part particle and part wave. The color of light corresponds to a
specific wavelength or frequency of electromagnetic energy.
Where: red light has the longest wavelength (above 700 nm) and a relatively low energy (frequency)
violet light has a shorter wavelength (around 400 nm) and a relatively high energy or frequency
Atomic Spectra - When atoms absorb energy in the form of excessive heat or electrical charge, electrons
move into higher energy levels. This is called the excited state as opposed to ground state
Ex. Na 2-8-1 ----------------------------------------> 2-7-1-1
ground
heat/electricity
excited state
state
These electrons then lose energy by emitting light when they return to lower energy levels. When this light
is passed through a diffraction grating or a spectroscope the wavelength/frequencies of light are separated
into bands of light called an atomic emission spectrum.
Every element has its own unique atomic emission spectrum. The same way no two people have the same
fingerprints, no two elements have the same emission spectrum. Ex. Helium
Atomic emission spectrums are used to identify elements
Ex. The spectrums of hydrogen, mercury, and neon are compared below. Each element always look the
same!
Ex. Use the four bright-line spectra data of lithium, hydrogen, helium and sodium to determine
the elements in the unknown sample.
Bohr used physics and mathematics to explain the bright-line spectrum of hydrogen.
Based on his data, he proved that hydrogen’s 1 electron could produce all 4 spectral lines in the bright line
spectrum by showing that the electron normally existed in a “ground state” energy level. But when heated,
the electron would become excited.
E absorbed
----------------->
Ground state
E released
---------------->
Excited state
ground state + light (spectral line)
or photon
Explanation of a all four spectral lines
Therefore – The Bohr model proposed that all atoms had electrons orbiting the nucleus in electron shells
(also called principle energy levels). These shells exist even when electrons are not present to fill them.
The electrons always remained in the ground state or lowest possible E level unless they were excited by E.
Bohr also proposed that only 2 e- could fit in the first energy level, 8 in the 2nd, and 8 in the 3rd.
Valence electrons – electrons in the highest occupied energy level of an atom
Valence electrons determine an atoms chemical properties, and are typically the only electrons involved in
chemical bonding. They are the "glue" that holds atoms together
The electron configurations of all atoms are shown in the bottom of the elements box in the periodic table
The number farthest to the right in these configurations is the number of valence electrons for the atoms for
each element
Quantum Mechanics - A new branch of physics that takes into account the wave and particle nature of
matter. This discovery of "matter waves" has lead to the modern wave-mechanical model of the atom
Main ideas are:
1. all moving objects have wavelike behavior
2. wavelengths of large objects is unobservable but the wavelengths of small particles like subatomic
particles exhibit wavelengths that are easily detectable.
3. Heisenburg uncertainty which applies to very small objects like electrons states: It is impossible to know
both the position and velocity of a particle at the same time.
Section Review Questions
1. What is the ground state electron configuration of calcium?
Write a possible excited electron configuration for calcium?
2. Which of the following electron configurations is an excited state for magnesium
2-8-2
2-7-4
2-8-3
2-7-3
3. Describe how an atomic emission spectrum is produced.
4. What is an excited electron?
5. What is a ground state electron?
6. Do any different elements have the same emission spectrum? explain.
7. What element could have a possible electron configuration of 2-8-17-3? Is this a ground state electron
configuration