Atomic Concepts (Unit 2) – Chapter 1
I: Subatomic Particles: An atom is the smallest unit of an element. It consists of three
major particles:
Particle
Mass
Location
Charge
Proton
1 amu
nucleus
+1
Neutron
1 amu
nucleus
No charge
Electron
1/1836 amu
orbitals
-1
Note:
a.m.u. = atomic mass unit
II: Atomic Models:
Atomic Theory: explains the structure of matter in terms of different combinations of very
small particles called atoms.
1) Dalton’s Theory (Cannonball)
a)
All elements are composed of indivisible atoms
b)
All atoms of a given element are identical
c)
Atoms of different elements are different; that is, they have different masses
d)
Compounds are formed by the combination of atoms of different elements
2) JJ Thomson: used a __cathode ray tube_ to show smaller units that make up an atom.
The ray was deflected a certain way by a magnetic field, so he concluded that the ray was
formed by particles and that the particles were negatively charged. The only source available
for the particles was the atoms present. Therefore, Thomson theorized that an atom contains
small, negatively charged particles called _electrons_. This theory is referred to as the
_Plum Pudding Model__. In this model, the mass of the rest of the atom was evenly
distributed and positively charged, taking up all of the space not occupied by the electrons.
3) Rutherford’s Gold Foil Experiment: If electrons are present in atoms, what makes up
the rest of the atom? Ernest Rutherford (with other scientists) directed positively charged
alpha particles at a piece of gold foil. If the plum pudding theory were correct, all of the alpha
particles would pass through the foil with just a few being slightly deflected. What they had
expected did in fact occur, but some of the alpha particles were greatly deflected and some
even bounced back. Rutherford concluded that:
a)
Majority of the volume of an atom is empty space
b)
Atoms have a dense positively charged central core
4) The Bohr Model
1 or K-shell = __max 2 e-____________
2 or L-shell = __max 8 e-____________
3 or M-shell = __max 18 e-__________
4 or N-shell = __max 32 e-__________
5) The Orbital Model:
*Important Definitions:
a) Principle Energy Level: Region around the _nucleus_ (the dense positively
charged central core of an atom) in which _electrons_ can be found. (The closer
to the nucleus, the lower the energy)
b) Quanta: Small amount of energy that a(n) _electron_ can absorb or release as _it
moves through principle energy levels__
c) Ground State:
all electrons fill lowest energy levels before higher energy levels
begin to fill
d) Excited State:
one or more electrons fills a higher energy level before the lower
ones are filled
e) Spectral Lines: As electrons at __higher_ energy levels (excited electrons) fall
back to their normal energy levels (ground state) they __release___ energy in the
form of the spectrum. ROYGBIV
IV: Electron Configurations
Looking at the periodic table of elements, you will notice numbers at the bottom of each
element. These numbers represent the electron configuration of the element (the address of
the electrons). The _Period___ represents the number of principle energy levels (orbitals)
present. The _Group_ represents the sublevel for each principle energy level.
Period 1 = 1 shell
Group 1 & Group 2 = ‘s’ sublevel (max – 2 electrons)
Period 2 = 2 shells
Group 13 – Group 18 = ‘p’ sublevel (max – 6 electrons)
Period 3 = 3 shells
Group 3 – Group 12 = ‘d’ sublevel (max – 10 electrons)
Lanthanum & Actinum Series = ‘f’ sublevel (max – 14 e)
Period 4 = 4 shells
Period 5 = 5 shells
Period 6 = 6 shells
Example: Write the electron configuration for the following:
Na
2-8-1
S
2-8-6
Kr
2-8-18-8
1) Valence Electrons and Electron Dot Diagrams:
a) Valence Electrons are electrons that fill the outermost principle energy level of an atom
Example: Mg 2-8-2 has __2__ valence electrons
Ne 2-8 has _8_ valence electrons
Valence electrons are largely responsible for an element’s chemical and physical
properties.
b) Electron Dot Diagrams: the term Kernel refers to all of the _non-valence_ electrons and
the __nucleus_ of an atom. The Kernel is represented by the element’s symbol, valence
electrons are represented by dots.
Example:
Na
Mg
N
Ne
V: Differences between atoms:
1) Atomic Number: the number of __Protons_ in the nucleus of an atom. (Atom by
definition is an electrically neutral particle, so this must also be equal to the number of
_Electrons).
2) Mass Number: the number of _protons_ plus _neutrons_
Question: Why are there fractional mass numbers on the periodic table? (ex: Na, O 2, …)
Answer: Due to the existence of _isotopes_.
Note: Atomic Symbols: One or two letters, 1st is always capital, the 2nd is always lower case.
3) Isotope: Atoms if the same element having the same number of protons, but different
number of neutrons
Example: Isotopes of Hydrogen
Particle
Protons
Neutrons
Mass Number
Symbol
Protium
1
0
1
1 H
1
Deuterium
1
1
2
2 H
1
Tritium
1
2
3
3 H
1
Calculating Isotopes (weighted atomic mass)

1) Take the percent of each isotope and convert it back to a decimal ( 100)

2) Multiply the decimal by the mass number

3) Add the numbers together to get the Weighted Atomic Mass
a.
C-12 99%
C-14
1%
(.99 x 12) + (.01 x 14) = 12.02
or
.99 x 12 =
11.88
.01 x 14 =
0.14
12.02
b.
Mg-26 1.75 %
Mg-24 98.25%
(0.0175 x 26) + (.9825 x 24) = 24.035
c.
Cl-35
75%
Cl-37 25%
(.75 x 35) + (.25 x 37) = 35.5
**
This is why we can’t round Chlorine to 35
4) Ion: Atoms of the same elements having
the same number of protons, but different
number of electrons
Example: Na  Na+1
F  F-1
a) Cations: positive charge formed by losing electrons
Example:
b) Anions:
Example:
negative charge formed by gaining electrons
THE HISTORY OF THE MODERN ATOMIC THEORY
400 BC
Democritus
Proposes the idea of the atom
1807
John Dalton
Founder of modern atomic theory
His model survived for almost 100 years
1885
J.J. Thomson
Discovery of the “proton”
1896
Henri Becquerel
Discovery of Radioactivity
1897
J.J. Thomson
Discovery of the “electron”
1903
J.J. Thomson
“Plum-pudding” model
1909
Ernest Rutherford
The Gold Foil Experiment
1911
Ernest Rutherford
Rutherford’s model of the atom
1913
Niels Bohr
Electrons exist in discrete energy levels
1923
Robert Millikan
The Oil Drop Experiment
Discovered the charge of an electron
Discovered the mass p+ = 1836 the mass of e-
1923
Erwin Schrodinger Wave Particle duality
1925
Erwin Schrodigner Quantum Mechanical Model of the Atom
1932
James Chadwick
Discovery of the “neutron”
1945
Wolfgang Pauli
Pauli Exclusion Principle