Download Unit 5 Notes - Har

Unit 5 Notes
“The Modern Atom”
ATOMIC HISTORY:
DATE
Scientist
Major Discovery Toward Chemistry
1
2
3
4
5
ELECTRON CONFIGURATION
Electron Configuration:

Let’s color our periodic table…

Highest occupied level – the electron-containing main energy level with the highest principal
quantum number

Inner-shell electrons – electrons that are not in the highest occupied energy level

Octet – eight electrons
Reynolds pg.1
Lets learn the “rules” first:
He:
Li:
C:
Ti:
Write electron configurations for the following atoms (on your own):
 H

Ne

P

Cu
Reynolds pg.2
Noble-Gas Notation:
o
o
o
An outer main energy level fully occupied
Once you pass the first two periods of elements, noble-gas notation can be used.
The previous noble gas has the first same levels of electrons filled, so the symbol of the noble
gas is placed in brackets to represent those electrons



Sodium: Na 1s22s22p63s1
[Ne] 3s1
Elements in the fourth period would use Argon [Ar]
Fifth period would use Krypton [Kr], etc.
Examples (Write both the electron configuration and noble gas notation):
Fe:
Ag:
Sr:
Orbital Notation
These show the actual placement of electrons in individual orbitals. Each orbital is represented as a
box, and only orbitals after the last filled noble gas are shown. To do these diagrams, you need to know
one more principle. Hund’s rule states that when electrons are being placed in a set of orbitals, they
must be placed singly with the same spin before beginning to pair them up. A set of orbitals is also
referred to as degenerate orbitals because they are of the same energy. If you follow Hund’s rule,
Pauli’s exclusion principle and the Aufbau principle, this will give the ground state of the atom.
Let’s write orbital diagrams for the following atoms:

N

O

Mn
Reynolds pg.3
For the following elements, write the following notations (electron configuration, noble gas notation, and
orbital notation)
Rb:
Si:
Cr:
QUANTUM NUMBERS & ORBITALS
I.
DeBroglie’s equation: DeBroglie found one of the first hints that the electron didn’t follow
Newtonian physics. Einstein had earlier used the equation E=mc2 to relate the energy of an
electron to its mass, meaning the electron was a particle. Planck had earlier used his equation E=h
to relate the energy of an electron to its frequency, meaning the electron was a wave. DeBroglie
combined the two equations to derive his formula:

h
mv
h=Plank’s const., m = mass, and v = velocity
This meant that anything that had mass, had a wavelength and vice versa. Therefore an electron
was both a particle and a wave. This is known as the particle-wave duality. Strange stuff, huh?
II.
Heisenberg’s uncertainty principle: Heisenberg came up with an idea that became the death of
Bohr’s idea of the atom. He said that it was impossible to know both the position and velocity of an
electron at the same time. Therefore it is impossible to know the exact location of an electron at
any given moment. This lead to the idea that we don’t have Bohr orbits, but rather, now we define
an orbit as an area where the probability of finding an electron is high.
Reynolds pg.4
III.
Schrödinger’s equation: Schrödinger applied the principles of wave mechanics to derive
information of about the electron. He used Heisenberg’s concept of the say that the probability of
finding the electron was equal to the square of the amplitude of the wave of the electron, ||2.
The solution to his equation consisted of four quantum numbers. You can think of these quantum
numbers as an electron’s address. The first three numbers give the specific orbital an electron can
be found in and the fourth tells how the electrons are placed in that orbital. Each number
determines the possible allowable values for the next.
Assigning Quantum Numbers:
What are quantum numbers? – Quantum numbers are a set of four numbers that scientists use to identify
the properties of an electron in an atom. A good analogy is to think of the quantum numbers as an
"address" of the electron in the atom.
o
Principal quantum number (n) – indicates the main energy level occupied by the electron
 Values are positive integers
 As n increases, so does the electron’s energy and its distance from the nucleus
 Called the electron shell
 The total number of orbitals that can exist in a given shell is n2
 number n
o
Angular momentum quantum number (l)
 EXCEPT for the first main energy level, orbitals of different shapes (sublevels) exist for a
given value of n
 l identifies the shape of the orbital
 the number of orbital shapes possible is equal to n
 values of l are zero to all positive integers less than or equal to n-1
 the orbitals are assigned a shape letter
Main Level
Sublevel
l value
letter shape
0
s
1
p
2
d
3
f
The angular momentum numbers occupy blocks on the periodic chart:
Magnetic quantum number (ml)
Orbitals



Indicates the orientation of an orbital around a nucleus
See pages 102-103
See Table 4-2 on page 104
Reynolds pg.5
o
Spin quantum number (ms) (+ ½ , - ½ )
 Indicates the two fundamental spin states of an electron in an orbital
 Only two electrons can exist in a single orbital and they must have opposite spins
 Go by the last draw arrow in orbital notation
Name
Principal Quantum
Number
Angular Momentum
Quantum Number
Magnetic Quantum
Number
Symbol
Spin Quantum number
ms
n
l
ml
How to find it
Row # in periodic table
Block (s, p, d, f)
-l to +l
(if n = 1, it is -1, 0, +1)
Last arrow drawn in orbital
notation
Meaning
Energy level occupied by
electron
Shape of the orbital
Orientation around the
nucleus (x, y, z planes)
Spin state (+1/2 or -1/2)
A) Hydrogen
B) Fluorine
C) Calcium
D) Titanium
Reynolds pg.6
E) Bromine
Rules for governing electron configurations:
1.
Aufbau principle – and electron occupies the lowest-energy orbital that can receive it. (sometimes it
takes less energy to fill a higher orbital)
Example:
2.
Pauli exclusion principle – no two electrons in the same atom can have the same set of four quantum
numbers (remember they must have different spins, but everything else may be the same)
Example:
3.
Hund’s rule – orbitals of equal energy are each occupied by one electron before any orbital is
occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin
Example:
Some Odds & Ends:
Ions:
-
Elements with a charge (due to the loss or gain of electrons)
Energy & Stability:
-
The further away from the nucleus the more energy the electron will have. However, the less stability.
Other ways to represent e- configuration (Quest):
Reynolds pg.7
Valence electrons
-
the electrons present in the highest occupied energy level. These electrons are important because
they are the electron involved in chemical reactions and basically define the chemical properties of
the element.
Write the electron configuration for the following elements and determine the number of valence electrons:
K:
As:
Zr:
Br:
*elements are placed in columns on the Periodic Table according to similar chemical properties. Is there any
relationship in the number of valence electrons for the atoms in the same family?
MODEL OF THE ATOM
Octet Rule- compounds form so that each atom either has __________ electrons (ionic) or feels like (covalent) it has
eight electrons in its ____________ shell.

Valence Shell Electrons- the _______________ electrons in an atom
Group 1- __________
Group 2- __________
Group 3-12 ________
Group 13- _________
Group 14- _________
Group 15- _________
Group 16- _________
Group 17- _________
Group 18- _________
Lewis Dot Notation- shows the number of ____________ electrons of an atom. The pattern of placing electrons is
shown below: These show only the electrons with the highest principal quantum number. They are placed as
dots around the atom, with the s electrons going as dots on top of the element symbol, and the p electrons being
placed around the remaining 3 sides. These electrons are shown because they are the electrons which normally
participate in chemical bonding. These diagrams will be very important as we start to talk about chemical
bonding.
*You have to first draw the orbital notation (from noble gas) in order to draw Lewis Dot Notation.*
Reynolds pg.8
Examples (Write the Lewis dot notation for the following elements):
A) Hydrogen
D) Oxygen
B) Fluorine
E) Nitrogen
C) Calcium
Diamagnetic vs Paramagnetic:
-
An element is said to be diamagnetic if there are no unpaired electrons. An element is said to be
paramagnetic if unpaired electrons are present.
Which of the following elements are diamagnetic and which are paramagnetic?
Mg
P
Reynolds pg.9
Extra Practice
Reynolds pg.10