Transcript Ch #9 Electrons and the Periodic Chart

Chapter #9
Atomic Theory
Quantum Model of the Atom
Physics Review
• J.J. Thompson determined all mater contains electrons
with cathode ray tube experiments.
• Rutherford demonstrated that atoms contain mostly empty
space by bombarding metals with alpha particles (helium
nuclei)
• Previously we talked about the plumb pudding model of
an atom and Rutherford’s planetary model.
• Then along came Bohr who applied Planck’s the ideas of
energy quantization to Rutherford’s planetary model of an
atom.
Evidence for Quantization of Energy
The circular orbits of the Bohr model are characterized by
the principle quantum number, n, which has positive integer
values. e.g.1, 2, 3….
About Energy
Energy is anything that has the capacity to do work.
Energy Examples:
Food; how much energy for a heart beat?
Gasoline; how many $ goes to heat and how much to work?
Dynamite; what can it move?
Electricity; what can it move?
Apple on a tree; what can it move?
The Bohr Model of an Atom
Rutherford proposed the electrons were located in
orbits around the nucleus similar to planets around a
star. This is sometimes called the planetary model,
which is not the modern day model of an atom.
Li atom
3p, 3n, 3e-
The Bohr Model of an Atom
Neils Bohr developed a mathematical model based
on Rutherford's proposal using Planck’s quantized
energy levels.
In Bohr’s model electrons could only have fixed
Potential energy levels.
Classification of Electromagnetic Radiation
White light is the collection of all of the colors in the
rainbow. When white light is passed through a glass
prism the waves are bent by the glass. The shortest
(most energetic) are bent the most.
Potential and Kinetic Energy
Potential energy; energy due to position above the
ground.
Kinetic energy; energy due to matter in motion.
Electromagnetic energy; energy possessing both
particle and wave properties traveling at the speed
of light.
Units of energy, joule, or calorie.
Properties of Waves
Wavelength (m)
Amplitude (m)
Speed 3.0X108 m/s
Energy (j)
Frequency (1/s, Hz)
Electromagnetic Radiation
10-12
Gamma rays
10-10
X- rays
10-8
Uv-rays
10-7
Visible
rays
10-4
10-2
Infrared
rays
Micro
Waves
(Radar)
Increasing wave length in meters
Increasing energy
100
102
Radio and Television waves
Visible Radiation Passing Through a Prism
Continuous Spectrum
Line Spectrum
Evidence for Quantization of Energy
The red-orange light from hydrogen gas passes
through a prism to form a line spectra. Each different
colored light has its own unique energy. This is called
an emission spectra
Additional Emission Line Spectra
Note: Each element has its own unique
emission spectra (Finger print)
Emission vs. Absorption Spectra
Early on physicists theorized that light emitted by
the sun should be a continuous spectrum. They were
troubled by black lines when observing the suns
spectrum through a glass prism. Bunsen (Bunsen
burner) and Kirchoff studied emission spectra from
emission tubes containing various gaseous elements.
Interesting the wavelengths of the black lines match
know wave lengths of different elements. This lead
physics to conclude that gases surround the sun that
absorb the emitted wave lengths of the sun.
Absorption Spectra
What is Quantization
Quantization means no in between
Energy levels in atoms are quantized.
Anything that comes in units such as: stairs
television channels, gears, and bookshelves are
quantized.
A turtle on stairs may only
be at specific heights. Its
potential energy is
quantized.
Quantization
Are you quantized right now?
Coming to class you
increase your
potential energy one
quanta (step) at a
time.
Evidence for Quantization of Energy
Evidence for Quantization of Energy
The circular orbits of the Bohr model are characterized by the
principle quantum number, n, which has positive integer values.
e.g.1, 2, 3….
Evidence for Quantization of Energy
In terms of the Bohr model absorption and emission looks like
this.
Evidence for Quantization of Energy
Electrons move between energy levels by absorbing and
emitting energy in the form of light.
We call the lowest energy level the ground state. The higher
energy level is called the excited state.
Evidence for Quantization of Energy
The Bohr model works well for the hydrogen atom which has
only one electron but performs poorly for more complex atoms.
This led to the development of the current quantum mechanical
model describing the arrangement of electrons in atoms.
Unfortunately for us this model is more complex than that
developed by Neils Bohr.
Quantum Numbers
The location of an electron in an atom can be described in order
of precision by its:
• shell (a positive integer given the symbol n = 1,2,3,….)
• subshell (Orbital) (designated by letters s, p, d or f)
• orientation (orbital symmetrical on x, y, or z axis)
•Spin (clockwise or counter clockwise)
+1/2
Like an electrons Social Security number:
4px
Quantum Numbers
As the value for n of a shell increases its energy and distance
from the nucleus increases. This is similar to the Bohr model.
Each shell has a number of subshells equal to its value for n (up
to a maximum of 4).
e.g. A shell with n = 1 will have one subshell (s)
A shell with n = 2 will have two subshells (s,p)
A shell with n = 3 will have three subshells (s,p,d)
A shell with n = 4 will have four subshells (s,p,d,f)
A shell with n = 5 will have four subshells (s,p,d,f,g)
Quantum Numbers
Each subshell is designated with the letter s, p, d or f.
The subshells are named by putting the value of n in front of the
symbol for the subshell.
e.g. a p subshell in the second shell is named 2p
The subshells vary in order of energy s < p < d < f.
The difference in energy between subshells is much smaller than
the difference in energy between shells. Just like the space
between buildings is smaller than the space between streets.
Quantum Numbers
As mentioned earlier each shell has a number of subshells equal
to its value for n.
Therefore for:
• n = 1 there will be one subshell 1s
• n = 2 there will be two subshells 2s and 2p
• n = 3 there will be three subshells 3s, 3p and 3d
• n = 4 there will be four subshells 4s, 4p, 4d and 4f
• n = 5 there will be four subshells 5s, 5p, 5d and 5f
Quantum Numbers
Each subshell is made up of one or more orbitals.
An orbital is a volume of space where an electron is likely to be
found.
What is the orbital for books called? Fish? Cars?
It is important not to confuse an orbit (a circular path on which
an electron moves in the Bohr model) with an orbital. They are
two very different things.
Quantum Numbers
An s subshell has one orbital which is spherically shaped.
If you were to measure where the electron was within an
s subshell many, many times and plot the results on a
graph you would get something like this.
Quantum Numbers
The p-orbital is next in energy after the s-orbital.
The p-orbital has a dumbbell shape with electrons located either
side of the nucleus in tear drop shaped lobes.
There are three types of p-orbitals all having identical shape
and energy directed along the x, y and z axis.
Quantum Numbers
Next in energy are the d-orbitals of which there are five all with
the same energy. The different d-orbitals do have different
shapes as well orientation. These are followed by the seven
f-orbitals.
Quantum Numbers
The way in which electrons are organized into shells, subshells
and orbitals in an atom is called the electronic configuration.
The electronic configuration of an atom can be determined
using the “Aufbau rule” also known as the “building up
principle”.
Aufbau comes from the German
meaning construction although it
was the Danish physicist Neils
Bohr who came up with the idea !!
Quantum Numbers
The Aufbau Principle states that:
“The orbitals of lower energy are filled in first with the
electrons and only then the orbitals of high energy are
filled.”
What is the lowest energy orbital of an atom?
1s orbital
What is the third lowest energy orbital of an atom?
2p orbital
Quantum Numbers
As we have seen previously for p, d and f subshells there are
multiple orbitals with the same energy, called degenerate
orbitals.
In particular:
• p subshells have three orbitals with the same energy
• d subshells have five orbitals with the same energy
• f subshells have seven orbitals with the same energy
Each of these orbitals may accommodate a maximum of two
electrons.
Quantum Numbers
If there are multiple orbitals with the same energy how do we
decide which orbital to put an electron?
We use Hund’s rule which states:
“Electrons fill degenerate orbitals one at a time
before doubling up in the same orbital”
Quantum Numbers
Using Hund’s rule how would we put three electrons in a p
subshell ?
px
px
py
py
p subshell

p subshell

pz
pz
Quantum Numbers
When we do put two electrons in one orbital then they obey the
Pauli exclusion principle.
“only electrons with opposite spin
can occupy the same orbital”
px
px
py
py
p subshell

p subshell

pz
pz
Quantum Numbers
How would we use our rules to “build up” the electron
configuration of a Li atom?
Li has Z = 3 so has 3 e-.
2s subshell
1s subshell
We can write this in shorthand as 1s22s1
Quantum Numbers
How would we use our rules to “build up” the electron
configuration of a N atom?
N has Z = 7 so has 7 e-.
2p subshell
2s subshell
1s subshell
We can write this in shorthand as 1s22s22p3
Quantum Numbers
Beyond the 3p subshell the orbitals don’t fill in an obvious way.
For example the 4s level lies lower in energy than the 3d .
Diagonal Rule
There is an easy way to
remember the sequence of the
energies of the subshells.
Quantum Numbers
So now we have everything we need to determine the electronic
configuration of any atom
What is the electronic configuration of a Na atom?
Start here
Z = 11, Na has 11 e1s2 2s2 2p6 3s1
Quantum Numbers
The electrons in the highest energy shell of an atom are (those
furthest from the nucleus) are called the valence electrons.
These electrons are very important as when atoms interact with
each other it is through their valence electrons.
Diagonal Rule Within the Periodic Chart
Periodic Chart
PERIODIC TRENDS
1. Atomic size
a. Increases from top to bottom (gaining a new outer
shell)
b. Decreases left to right in a period (Increase of protons
attracts electrons stronger, thus contracting the
element.
2. Ease of ionization (relative ease of losing electrons)
1. The larger the atom, the further electrons are from
nucleus, and the less they are held by nucleus.
3. Metals and nonmetals
a. Metallic character (increases going away from metals)
b. Nonmetallic character
Periodic Trends
Metals, Nonmetals, and Mettaloids
Chapter #9
Review
How many electrons fill the
following sublevels?
a. The 3rd level.
b. The 4th level.
th
c. The 5 level
How many electrons fill the following
sublevels?
a. The 3rd level. 2(n)2 =2(3)2=18
b. The 4th level.
th
c. The 5 level
How many electrons fill the following
sublevels?
a. The 3rd level. 2(n)2 =2(3)2=18
b. The 4th level. 2(n)2 =2(4)2=32
th
c. The 5 level
How many electrons fill the following
sublevels?
a. The 3rd level. 2(n)2 =2(3)2=18
b. The 4th level. 2(n)2 =2(4)2=32
th
2
2
c. The 5 level 2(n) =2(5) =50
How many orbitals are in an
atom with n=3?
How many orbitals are in an
atom with n=3?
Answer:
n = 3 refers to the third energy level, which contains the
s, p, and d sublevels. There is only one orbital in the s,
three in the p, and five in the d sublevels; for a total of 9
orbitals.
Identify the element with the following electron
configurations, also stating the group and period
the where the element where found.
a. 1s22s2p63s23p64s23d104p4
b. [Ar]3s23d7
Identify the element with the following electron
configurations, also stating the group and period
the where the element where found.
a. 1s22s2p63s23p64s23d104p4
Se, Period 4, Group VI (6)
b. [Ar]3s23d7
Identify the element with the following electron
configurations, also stating the group and period
the where the element where found.
a. 1s22s2p63s23p64s23d104p4
Se, Period 4, Group VI (6)
b. [Ar]4s23d7
Co, Period 4, Group 9
The End