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Models of the Atom
Section 13.1
The story of how the atomic
theory has evolved over time….
John Dalton


The atom is a solid indivisible mass.
He had 4 key ideas: Page 107

All elements are composed of tiny indivisible
particles called the atom.
Atoms of the same element are identical.
Atoms of different elements are different.
Atoms combine chemically with one another in

simple whole-number ratios.
During chemical reactions, atoms are separated,



joined, or rearranged. Atoms are never created
nor destroyed.
JJ Thomson

Plum-pudding model.



Sometimes called the Chocolate Chip Cookie
Model
The atoms has negatively charged
electrons stuck into a lump of positively
charged material, like chocolate chips
stuck in a cookie dough.
Did not address protons and neutrons.
Ernest Rutherford




Discovered the nucleus.
Showed that most of an atom’s mass is
concentrated in a small, positively
charged region called the nucleus.
Electrons resided on the outside.
Did not address how electrons were
arranged.
Neils Bohr



Electrons are arranged on concentric
circular paths, or orbits around the
nucleus.
Solar system model or planetary model.
Gave us the idea of definite energy
levels.
Quantum Mechanical Model –
Our Currently Accepted Model

Erwin Schrodinger




Primarily a mathematical model using quantum
mechanics
It addresses “probabilities” of finding an electron
at any instant in an area called “electron clouds”.
Introduced the ideas of Principal Energy Levels
and Sublevels of energies.
The electron clouds take certain shapes,
represented by the s,p,d,f subatomic orbitals.
Principal Energy Levels

Just like the Bohr model, the Quantum
Mechanical Model designates energy levels
of electrons by means of principal
quantum numbers called (n).

Principal Energy Levels refers to a
major region where electrons are most
likely to be found.

They are assigned values in order of
increasing energy: n=1, n=2, n=3, etc.
Sublevels

Within each principal energy level, the
electrons occupy energy sublevels.

The number of sublevels within each
principal energy level is the same as
the principal quantum number.

How many sublevels does the 4th
principal energy level have?
Atomic Orbitals


The regions in which electrons are likely
to be found are called atomic orbitals.
Letters denote the atomic orbitals




S-shape orbitals are spherical
P-shape orbitals are hour-glass shapes
D-shape orbitals have clover-leaf shapes
Draw an example of each into your
notes.
Exploring further…


The lowest principal energy level (n=1)
has only one sublevel, called 1s.
The second principal energy level (n=2)
has 2 sublevels, the 2s and 2p. The 2p is
higher in energy and consists of 3 p
orbitals.

Let’s show an easy way to remember this
order of filling electrons…
The Electron Pyramid




The s orbitals have 1 spatial orientation,
therefore can hold 2 electrons
The p orbitals have 3 spatial orientations,
therefore can hold 6 electrons
The d orbitals have 5 spatial orientations,
therefore can hold 10 electrons
The f orbitals have 7 spatial orientations,
therefore can hold 14 electrons.
Electrons Fill following 3 simple
rules


Aufbau principle: Electrons enter the
lowest energy level first.
Pauli Exclusion Principle: An atomic
orbital may describe at most 2 electrons,
both spinning in opposite directions.

Hund’s Rule: When electrons occupy
orbitals of equal energy, one electron
enters each orbital until all the orbitals
contain one electron with parallel spins.
Classifying Elements by Electron
Configuration



Of the three major subatomic particles, the
electron plays the most significant role in
determining the physical and chemical
properties of an element
The arrangement of elements in the periodic
table depends on these properties
Elements can be classified into four categories
according to their electron configurations
1.
2.
3.
4.
The noble gases
The representative elements
The transition metals
The inner transition metals
The Noble Gases


These are elements in which the outermost s
and p sublevels are filled
Belong to Group 0 (Group 8A)
Configurations
Helium
Neon
Argon
Krypton
1s2
1s22s22p6
1s22s22p63s23p6
1s22s22p63s23p64s23d104p6
The Representative Elements


In these elements, the outermost s or p sublevel is only
partially filled
Usually called the Group A elements and include the noble
gases in some definitions
Configurations
Lithium
Sodium
Potassium
Carbon
Silicon
Germanium
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p2
1s22s22p63s23p2
1s22s22p63s23p64s23d104p2
The Transition Metals


These are metallic elements in which the outermost s
sublevel and nearby d levels contain electrons
Called the Group B elements
Configurations
Zinc
1s22s22p63s23p64s23d10
Tungsten
1s22s22p63s23p64s23d104p65s24d105p66s24f145d4
Lead
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p2
The Inner Transition Metals


These are metallic elements in which the
outermost s sublevel and nearby f sublevel
generally contain electrons
Characterized by the filling of f orbitals
Configurations
Uranium
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f4
Curium
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f8
S1
S2
1
2
S2
3
Period
5
d1
d2
“S” Block
4
p1
d3
d4
d5
d6
d7
d8
d9
p2
p3
p4
p5
d1
0
“P” Block
“D” Block
6
7
Period
F1
6
7
F2
F3
F4
F5
F6
F7
F8
“F” Block
F9 F10
F11
F12
F13
F14
p6
Practice

Write the electron configurations for the following
elements:
Nitrogen
1s22s22p3
Nickel
1s22s22p63s23p64s23d8
Vanadium
1s22s22p63s23p64s23d3
Practice

What are the symbols for all the elements that have the following
outer configurations?
S2
Helium (He)
Beryllium (Be)
Magnesium (Mg)
Calcium (Ca)
Strontium (Sr)
Barium (Ba)
Radium (Ra)
S2P5
Fluorine (F)
Chlorine (Cl)
Bromine (Br)
Iodine (I)
Astatine (At)
S2D2
Titanium (Ti)
Zirconium (Zr)
Hafnium (Hf)
Rutherfordium (Rf)
Exceptional Electron
Configurations




Chromium and Copper have exceptional
electron configurations.
They fill their d sublevel completely,
leaving their 4s partially filled.
Much more stable this way!
Write them correctly into your
notepacks…
Cr: 1s22s22p63s23p63d54s1
Cu: 1s22s22p63s23p63d104s1
Physics and the Quantum
Mechanical Model



This section studies the electron as a
property of light.
Electrons travel as waves and are made
of particles of light called photons
According to the wave model, light
consists of ELECTROMAGNETIC
RADIATION.
Electromagnetic Spectrum

This form of energy includes










Gamma rays
X-rays
Ultraviolet rays
Visible light
Infrared rays
Radar
FM
TV
Shortwave
AM
Electromagnetic Spectrum

Every element emits light when it is
excited by the passage of electric
discharge through its gas or vapor.

The atoms first absorb energy, then
lose the energy as they emit light.
Electromagnetic Spectrum

Electrons are said to move from their
GROUND STATE (lowest energy level)
to and EXCITED STATE (higher energy
level).

When the electron falls back to its lower
energy, it emits a PHOTON of energy,
and can be seen in the visible spectrum.
Electromagnetic Spectrum

Passing the light emitted by an element
through a prism gives the ATOMIC
EMISSION SPECTRUM of the element.

Because each atom has a unique
electron arrangement, each atom emits
a unique wavelength during this
process. This wavelength falls within
the visible spectrum.
Kernel Structures

The kernel is a structure used to
shorten an electron configuration.

A kernel is an inert gas symbol in
brackets that stands in place of all of
the filled orbitals contained in the inert
gas.

Examples
Section 14.1
Classification of the
Elements By Electron Configuration
Classifying Elements by
Electron Configuration

Of the three major subatomic particles,
the ELECTRON plays the most
significant role in determining the
properties of an element.

The arrangement of elements in the
PERIODIC TABLE depends on these
properties.
Elements can be classified into 4
categories:

The Noble Gases :
These are elements in which the
outermost s and p sublevels are filled.
Write for Helium, Neon, Argon, Krypton


The representative elements:
In these elements, the outermost s
and p sublevel is only partially filled.
Write for Lithium, Sodium, Potassium,
Carbon, Silicon, Germanium

Elements can be classified into 4
categories:

The transition metals:

These are metallic elements in which
the outermost s sublevel and nearby d
sublevel contain electrons.


Write for Zinc and Zirconium.
The inner transition metals:

These are metallic elements in which
the outermost s sublevel and nearby f
sublevel generally contain electrons.
14.2 Periodic Trends

Atomic radius – ½
the distance
between the nuclei
of two like atoms in
a diatomic molecule.
Group
Trends

Atomic size generally increases as
you go down a group on the
periodic table.

Adding additional energy levels!
Periodic
Trends

Atomic size generally decrease as you
move from left to right across a period.

Same energy level increasing nuclear charge pulls electrons
closer to nucleus.
Ionization Energy

An ion – a charged atom that results
from either losing or gaining an
electron.

Ionization Energy – The energy
required to overcome the attraction of
the nuclear charge and remove an
electron from a gaseous atom.

(The ease of losing an electron and
forming a +1 charge)
Ionization Energy

First ionization energy – the energy
needed to remove the first electron
from an atom.

Second ionization energy – the energy
needed to remove the second electron
from an atom, etc.
Ionization Energy

Group Trends: The first ionization
energy generally decreases as you
move down a group on the periodic
table.

The size of the atoms increases, so the
outermost electron is farther from the
nucleus and will be more easily
removed.
Ionization Energy

Periodic Trends – For the representative
elements, the first ionization energy
generally increases as you move from
left to right across a period.

Increasing nuclear charge makes it
more difficult to remove an electron.
Ionic Size

The atoms of METALLIC elements have
low ionization energies. They form
POSITIVE ions easily.

By contrast, the atoms of
NONMETALLIC elements readily form
NEGATIVE ions.
Trends in Ionic Size

Positive ions are always smaller than
the neutral atoms from which they
form.
 They lose their outer shell electrons

Negative ions are always larger than
the neutral atoms from which they
form.
 This is because the effective nuclear
attraction is less for an increased number
of electrons.
Trends in Electronegativity

Electronegativity – the tendency for the
atoms of the element to attract electrons
when they are chemically combined with
atoms of another element.

Electronegativity generally DECREASES as
you go down a group.

As you go across a period from left to right,
the electronegativity of the representative
elements INCREASES.
Electronegativity

The electronegativity of cesium, the
least electronegative element is 0.7

The electronegativity of fluorine, the
most electronegative element, is 4.0

Electronegativity values help predict the
type of bonding that can exist between
atoms in compounds, either IONIC OR
COVALENT bonds.
Summary of Periodic Trends

Using page 406, create a summary of
periodic trends into your notes.