Download Atomic Orbitals - Harding Charter Preparatory High School

Chapter 10
Modern Atomic Theory
Models of the Atom
• Rutherford’s model of an atom stated that an atom is
mostly empty space with single dense positive core with
negatively charged electrons moving around the nucleus
in some way
– Rutherford's model of an atom could not explain why metals or
compounds of metals give off characteristic colors when heated
in a flame, or why objects, when heated to higher and higher
temperatures, first glow dull red, then yellow, then white
– Rutherford’s atomic model could not explain the chemical
properties of elements
Electromagnetic Radiation
• Our current model of atoms and how electrons are
arranged around the atom came from the study of light
Electromagnetic Radiation
• Light is a form of electromagnetic radiation
• Electromagnetic radiation
– Radiant energy that exhibits wavelike behavior and travels
through space at the speed of light in a vacuum
Wave Length (λ)
Crest
Amplitude
Trough
𝐹𝑟𝑒𝑞𝑢𝑒𝑛𝑐𝑦 ν = 𝐶𝑦𝑐𝑙𝑒𝑠 𝑆𝑒𝑐𝑜𝑛𝑑
1 Second
4 Cycles
Electromagnetic Radiation
– Electromagnetic radiation consists of waves
• Waves cycle
– Each complete wave cycle starts at zero, increases to highest
value (the crest), passes through zero and down to its trough
and back up to zero
– The amplitude of the wave is its height at the crest
– The wavelength, represented by λ (lambda), is the distance
between the crests
– The frequency, represented by ν (nu), is the number of waves
that to pass a given point per unit of time
» The units of frequency are usually cycles per second called
a hertz (Hz) or S-1
– The product of frequency and wavelength always equals a
constant (C), the speed of light 3.00x108 m/s
c = λν
– The wave length and frequency of light are inversely
proportional to each other
Electromagnetic Radiation
• Light is a form of electromagnetic radiation that includes
radio waves, microwaves, infrared waves, visible light,
ultraviolet light, x-rays, and gamma rays
– All electromagnetic rays travel in a vacuum at a speed of 3.00 x
108 m/s
• Sunlight consists of light with a continuous range of
wavelengths and frequencies
• When the different frequencies of light pass through a
prism they are separated into a spectrum
Electromagnetic Radiation
http://upload.wikimedia.org/wikipedia/commons/c/cf/EM_Spectrum_Properties_edit.svg
Atomic Spectrum of Hydrogen
• When H2 molecules absorb energy, some H-H
bonds are broken and the resulting hydrogen
atoms are excited
– When excited, the hydrogen atoms contain excess
energy that is released in the form of light at specific
wavelength producing an special emission spectrum
called a line spectrum
– The spectrum indicated that only certain energies are
allowed for the electron in the hydrogen atom
• In other words, the energies were quantized
Atomic Spectra
• When atoms absorb energy, electrons move into higher
energy levels, and these electrons lose energy by
emitting light when they return to lower energy levels
• Light emitted by an element separates into discrete lines
to give an atomic emission spectrum of the element
• Each discrete line in an emission spectra corresponds to
one exact frequency of light emitted by the atom
• The emission spectrum of each element is like a
person’s finger print and can be used to identify each
element
Atomic Spectra and the Bohr Model
• Niels Bohr proposed that an electron is found
only in specific circular paths, or orbitals, around
the nucleus
– Each possible electron orbit in Bohr’s model has a
fixed energy.
• The fixed energies an electron can have are called energy
levels
• To move from one energy level to another, an electron must
gain or lose just the right amount of energy called a quantum
– Thus, the energy of electron is said to be quantized
– The energy levels are not equally spread, the higher energy
levels are closer together
The Bohr Model
• The electron in a hydrogen
atom moves around the
nucleus only in certain allowed
circular orbits
n=5
n=4
n=3
n=2
n=1
 The orbits are known as
principle energy levels aka
shells
Line
Spectrum
Wavelength
 When not excited, the electron in
the hydrogen atom resides in the n
= 1 energy level.
 A certain quantum of energy is
required for an electron to move to
a higher shell (energy level)
An explanation of Atomic Spectra
• The lowest possible energy of the electron is its ground state
• When an electron absorbs energy and is excited, it jumps in
quanta’s to higher energy levels
• A quantum of energy in the form of light is emitted when the
electron drops back to a lower energy level
• The emission occurs in a single abrupt step
• The quanta of energy released is related to the frequency of the
emitted light by the equation E = hν where h = 6.626x10-34 J-s
• The light emitted by an electron moving from a higher to a
lower energy level has frequency directly proportional to the
energy change of the electron
• Light has both wave and particle properties
Light Frequency, Wavelength, and Energy Relationship
𝐶 = λν
𝐸 = ℎν
• Increasing wave length result in decreasing frequency which
result in light with less energy
• Increasing frequency results in decreasing wavelength which
result in light with more energy
The Quantum Mechanical Model
Schrodinger
• Like the Bohr model, electrons are restricted to certain
energy levels
• Unlike the Bohr model, the quantum mechanical model
does not involve an exact path the electron takes around
the nucleus
– Electron paths are not circular
• 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
Quantum Mechanics
• Photons have wave an particle properties
• Everything travels in wave like motion but the mass of the particle
must be small enough to observe the wave like nature
• Heisenberg uncertainty principle states that is impossible to know
exactly both the velocity and the position of a particle at the same
time
Quantum Mechanical Model
• Heisenberg uncertainty principle
– There is a fundamental limitation to how
precisely both the position and momentum of
a particle can be known at a given time
Atomic Orbitals
• Atomic orbitals are often thought of as a region of space
in which there is a high probability of finding an electron
– Each orbital is characterized by a series of numbers called
quantum numbers, which describe various properties of the
orbital:
• Energy levels of electrons in the quantum mechanical model are
called principle quantum numbers (n) and are assigned the
numbers 1,2,3,4, and so forth
Atomic Orbitals
•
An atomic orbital is often thought of as a region of space in which there is a
high probability of finding an electron
•
For each principle energy level, there may be several orbitals with different
shapes and different energy levels that constitute energy sublevels
(subshells)
•
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: s, p, d, and f
–
–
–
–
•
The number and kinds of atomic orbitals depend on the energy sublevel
–
–
–
–
•
There is only one type of s orbital
There are three types of p orbitals
There are five types of d orbitals
There are seven types of f orbitals
The lowest energy level n = 1 can only hold one orbital s
The second energy level n = 2 can hold four orbitals s, p1, p2, p3
The third energy level n = 3 can hold 9 orbitals s, p1, p2, p3, d1, d2, d3, d4, d5
The fourth energy level n = 4 can hold 16 orbitals ...
Two electrons with opposite spin can potentially occupy each orbital
– The maximum number of electrons that can occupy a principle energy level is
given by the formula 2n2, where n is the principle quantum number
Atomic Orbitals
Electron Arrangements in Atoms
• In atoms, electrons and the nucleus
interact to make the most stable
arrangement possible
• Three rules you need to know about
electron arrangements are: the Afbau
principle, the Pauli exclusion principle, and
Hund’s rule that tell how to determine the
electron arrangement of atoms
Electron Arrangements in Atoms
• Aufbau principle – electrons occupy the orbitals of lowest energy
first
– The range of some energy levels within a principle energy level can
overlap the energy levels of another principle level
• Pauli exclusion principle – an atomic orbital may describe at most
two electrons
– To occupy the same orbital, the two electrons must have opposite spin
represented with an up or down arrow ↑↓
• Hund’s rule – 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
– One electron enters each orbital until all the orbitals contain one
electron with the same spin direction
• Some actual electron configurations differ from those assigned using
the aufbau principle because half-filled sublevels are not as stable
as filled sublevels, but they are more stable than other
configurations
• Find the electron
configuration of Zinc
1s2
1. Find Zinc’s atomic number
• Zinc has an atomic number
of 30.
2s2
2p6
3s2
3p6
3d10
4s2
4p6
4d10
4f14
5s2
5p6
5d10
5f14
6s2
6p6
6d10
6f14
2. Follow the arrows adding the
superscripts until you reach 30
Answer 1s2 2s2 2p6 3s2 3p6 4s2 3d10
Notice 2 + 2 + 6 + 2 + 6 + 2 + 10 = 30
3. Rearrange electron
configuration in order of
increasing energy level
Answer 1s2 2s2 2p6 3s2 3p6 3d10 4s2
• Find the electron configuration
of Ruthenium (Ru)
1s2
1. Find Ruthenium’s atomic number
•
2s2
2p6
3s2
3p6
Ruthenium has an atomic number of 44.
2. Follow the arrows adding the super scripts
until you reach 44
3d10
Answer 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d6
4s2
4p6
4d10
4f14
5s2
5p6
5d10
5f14
6s2
6p6
6d10
6f14
Notice 2 + 2 + 6 + 2 + 6 + 2 + 10 + 6 + 2 + 6 = 44
Notice Only 6 electrons were needed for the 4d
sublevel
3. Rearrange electron configuration in order of
increasing energy level
Answer 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d6 5s2
• Writing Orbital Diagrams
1s2
2s2
2p6
3s2
3p6
3d10
4s2
4p6
4d10
4f14
5s2
5p6
5d10
5f14
6s2
6p6
6d10
6f14
1s2
2s2
2p6
1.
Write orbital boxes for each sublevel that
contains electrons in the same order as you
would an initial electron configuration
2.
Add one electron (represented by an up
arrow) to each box (orbital) in a sublevel
until each orbital in that sublevel contains
one electron
3.
Once all orbitals in the sublevel contain one
electron start paring them up with a down
arrow until all boxes in the sublevel are filled
4.
Once one sublevel is filled, move to the next
sublevel repeating steps 3 and 4 until you
have as many arrows total as you do
electrons in the atom
5.
The orbital diagram for
3s2
3p6
4s2
26𝐹𝑒
is shown below
3d10
• Writing Orbital Diagrams
1s2
2s2
2p6
3s2
3p6
3d10
4s2
4p6
4d10
4f14
5s2
5p6
5d10
5f14
6s2
6p6
6d10
6f14
1s2
2s2
2p6
1.
Write orbital boxes for each sublevel that
contains electrons in the same order as you
would an initial electron configuration
2.
Add one electron (represented by an up
arrow) to each box (orbital) in a sublevel
until each orbital in that sublevel contains
one electron
3.
Once all orbitals in the sublevel contain one
electron start paring them up with a down
arrow until all boxes in the sublevel are filled
4.
Once one sublevel is filled, move to the next
sublevel repeating steps 3 and 4 until you
have as many arrows total as you do
electrons in the atom
5.
The orbital diagram for
3s2
3p6
4s2
26𝐹𝑒
is shown below
3d10
Aufbau Principle and the Periodic Table
Valence Electrons
• Valence electrons are the electrons in the
outermost principal quantum level of an atom
– The inner electrons are known as core electrons
• Valence electrons are the most important in
determining the chemical properties of an
element
• The elements in the same group have the same
valence electron configuration
Periodic Trends
• Atomic Size
• Ionization Energy
• Electronegativity
Periodic Trends: Atomic Size
Periodic Trends
– The atomic radius is one half of the distance between
the nuclei of two atoms of same element when atoms
are joined
– In general, atomic size increases from top to bottom
within a group and decreases from left to right across
a period
– Why
• Periodic Trends
– The increases of positive charge in the nucleus of the atom
draws in it’s electrons in its highest energy level as you move
left to right across a period and thus, the atomic size decreases
• Group Trends
– As you move down a group, new principle energy levels are
filled with electrons. These filled inner energy levels shield the
electrons in the outer energy levels from the positive charge of
the nucleus and thus atoms are larger down a group in the
periodic table
Periodic Trends
• Ions
– Cations are smaller than the atoms they were
formed from
– Anions are larger than the atoms they were
formed from
Periodic Trends
• Trends in Ionization Energy
– Electrons can move to higher energy levels when atoms absorb
energy
– If there is enough energy, electrons overcome the attraction of
protons in the nucleus and escape
– The energy required to remove an electron from an atom is
called ionization energy
• Measured with when an element is in its gaseous state
• The energy required to remove the first electron from an atom is
called the first ionization energy and produces a cation with a 1+
charge
• The first ionization energy tends to decrease from top to bottom
within a group and increases from left to right across a period
– Atoms with low ionization energy levels tend to lose electrons
easily in chemical reaction becoming cations
– Atoms with high ionization energy tend to gain electrons in
chemical reaction becoming anions
Periodic Trends: Ionization Energy
Periodic Trends
• Trends in Ionization Energy
– Periodic Trends in ionization energy
• Increases from left to right due to increasing
nuclear charge, and constant shielding of the outer
most electrons
– Results in increased attraction of the outer electrons to
the nucleus and greater energy required to remove
outermost electron
– Group Trends in ionization energy
• Decreases from top to bottom in a group due to
increased shielding of outer electron by completely
filled inner energy levels and the outer electrons
further distance from the nucleus
– Results in decreased attraction of the outer electrons to
the nucleus and less energy required to remove
outermost electron
Periodic Trends: Electronegativity
Trends in Electronegativity
• Electronegativity
– When atoms react to form a compound, two kinds of bonds can
result between the atoms called ionic bonds and covalent bonds
– Whether an ionic or covalent bond forms can be predicted by the
electronegativity of the atoms involved in making the compound
• Electronegativity is the ability of an atom of an element to attract
electrons when the atom is in a compound
• Scientist use factors such as ionization energy to calculate values of
electronegativity
– In general, electronegativity values decrease from top to bottom
within a group. For representative elements, the values tend to
increase from left to right across a period
• Metals at the far left of the periodic table have low electronegativity
values
• Nonmetals at the far right of periodic table (excluding noble gases)
have high electronegativity values
• Electronegativity among transition metals are not as regular
Bond Type and Polarity
• The difference in electronegativity
between the two atoms involved in the
bond will determine the most probable
type of bond that will form
– If the difference is:
•
•
•
•
0.0—0.4 the bond is nonpolar covalent
0.4—1.0 the bond is moderately polar covalent
1.0—2.0 the bond is very polar covalent
>2.0 the bond is ionic