Download File - Lenora Henderson`s Flipped Chemistry Classroom

Electrons In Atoms
5.1
 After
discovering the nucleus, Rutherford
used existing information about the atom to
create the nuclear model of the atom
(electrons rotate around the nucleus similar
to how planets orbit the sun)
 Rutherford’s
model could not explain the
chemical behavior of elements

Example: why metals or compounds of metals
give off characteristic colors when heated
 The
explanation of what leads to the
chemical properties of elements required a
model that would show the behavior of
electrons in atoms
 Neils
Bohr was a Danish physicist and a
student of Rutherford
 In 1913 he developed a new atomic model
that incorporated how the energy of an atom
changes when the atom absorbs or emits
light
Key Question
What did Bohr propose in his model of the
atom?
That an electron is found only in specific
circular paths, or orbits, around the nucleus.
 The
electron orbits in Bohr’s model has a fixed
energy
 The fixed energies an electron can have are called
energy levels
 The energy levels increase from bottom to top
 Electrons can move from one energy level to
another
 Electrons cannot exist between energy levels
 To move from one energy level to another the
electron must gain or lose the right amount of
energy
A quantum of energy is the right amount of
energy required to move an electron from one
energy level to another; therefore, the energy of
an electron is said to be quantized
 The amount of energy that an electron gains or
loses in an atom may vary, it is not always the
same
 These energy levels are not equally spaced

The higher energy levels are closer together
 The higher the energy level occupied by an electron,
the less energy it takes the electron to move from
that energy level to the next higher energy level


Disadvantage

Bohr only studied the simplest atom (Hydrogen);
therefore, his model did not explain the energies
absorbed and emitted by atoms with more than one
electron
 Rutherford
and Bohr both related the movement of
electrons to a large moving object
 However, later calculations and experimental
results proved otherwise
 Erwin Schrodinger, an Austrian physicist, used
calculations and results to devise and solve a
mathematical equation that describe the
movement of an electron in the hydrogen atom
 The quantum mechanical model came from the
mathematical solutions to the Schrodinger
Equation, which is the modern description of the
electrons in atoms
 Similarities

Both the Bohr and the Quantum Mechanical
Model restricts the energy of electrons to certain
values
 Differences

The QMM does not specify an exact path the
electron takes around the nucleus
Key Question
What does the QMM determine about the
electrons in an atom?
It determines the allowed energies an
electron can have and how likely it is to find
the electron in various locations around the
nucleus of an atom.






Concerning the QMM, probability describes how likely it is
to find an electron in a particular area surrounding the
nucleus
The probability of finding an electron within a certain
volume of space surrounding the nucleus can be
represented as a fuzzy cloudlike region
The cloud is more dense where there is a high probability
of finding an electron, and less dense where there is a low
probability of finding an electron
There is no outer boundary to the cloud because there is a
slight chance of finding the electron at a considerable
distance from the nucleus (opposites attract)
Therefore, attempts to show probabilities as a fuzzy cloud
are usually limited to the volume in which the electron is
found 90 percent of the time
It is impossible to know the exact location of an electron
at any given time
 Solutions
to the Schrodinger equation give
the energies, or energy levels, an electron
can have
 The Schrodinger equation leads to a
mathematical expression called an atomic
orbital
 An atomic orbital is a mathematical
representation that describes the probability
of finding an electron at various locations
around the nucleus; which is represented as
a region where there is a high probability of
finding an electron
 The
energy levels of electrons in the QMM are
labeled by principal quantum numbers (n), which
are assigned n = 1, 2, 3, 4, and so on
 The principal energy levels that are higher than 1
have several orbitals with different shapes and at
different energy levels
 These energy levels within a principal energy
level constitute energy sublevels
Key Question
How do sublevels of principal energy levels
differ?
Each energy sublevel corresponds to one or
more orbitals of different shapes. The
orbitals describe where an electron is likely
to be found.
 Different
atomic orbitals are distinguished by letters
 s orbitals are spherical in shape

p


The probability of finding an electron in this orbital does
not depend on direction because it is spherical in shape
orbitals are dumbbell-shaped
These orbital have different orientations in space
px, py,pz
d
orbital has five orbitals, four of the five are
clover-leaf shaped
 f orbitals are more complicated
See page 131
 The
number and types of atomic orbitals
depend on the principal energy level
 The number of principal energy levels equals
the number of sublevels in that principal
energy level
 The number of orbitals in a principal energy
level is equal to n2
 Only 2 electrons can occupy an orbital

The total number of electrons that can occupy a
principal energy level is given by the formula 2n2
Table 5.1 page 132
Summary of Principal Energy Levels and Sublevels
5.2
This diagram may not always be provided for
you; therefore, you must know the shorthand
way to create this diagram.
Key Question
What are the three rules for writing
electron configurations of elements?
Aufbau principle, Pauli exclusion principle,
and Hund’s rule are used to determine
electron configurations of atoms.
 Aufbau
principle states that electrons
occupy the orbitals of lowest energy first
 Orbitals for any sublevel of a principal
energy level are always of equal energy
 The s sublevel is always the lowest in energy
within a principal energy level

The range of energy levels within a principal
energy level can overlap the energy levels of
another principal energy level
Aufbau Diagram
Each box represents an orbital
The energy increases from the bottom to the
top
 Pauli
exclusion principle states that an
atomic orbital may describe at most two
electrons



The two electrons must have opposite spins
(paired spins)
Spin is a quantum mechanical property of
electrons and may be thought of as clockwise or
counterclockwise ( or )
Paired spins are represented as
 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
Example: Diagram on Board
(Will be shown in class!)
 The
electrons will then occupy each orbital
so that their spins are paired with the first
electron in the orbital
Extended Electron Configurations
The sublevels within the same principal energy
level are written together
Model
N
1s22s22p3
Sc
1s22s22p63s23p63d14s2