Objective – To understand what atoms
are and how their characteristics
determine the periodic table.
 Throughout
history, scientists have tried to
explain what made up matter
 Democritus – Greek philosopher

First to propose “atoms”

Invisible, indestructible, fundamental units of matter
 Formulated
theory”
the first “atomic
Elements are made of tiny
particles called atoms.
 All atoms of a given element are
identical, but different from
atoms of any other element.
 Compounds are formed when
atoms of different elements
combine in fixed proportions.
 A chemical reaction involves the
rearrangement of atoms, not a
change in the atoms themselves.

Found that atoms are made of
smaller particles
 Used cathode ray tube to shoot a
beam of electrons that travel in a
straight line




Put magnets on sides of tube and the
ray bent towards the positive side.
Since opposites attract, electrons must
be negative
His model was that electrons
floated in a soup of positive
particles

“plum pudding” or “chocolate chip
cookie” model
 Millikan
– first
measured the
electrical charge of
an electron

Oil drop method – put
a charge on a drop of
oil and dropped it
between two charged
plates. He would
adjust the power of
the plates to suspend
the drop in mid-air,
defying gravity
 Goldstein
– found
the proton

Used a cathode ray
tube to observe canal
rays (protons)
traveling in opposite
directions of cathode
rays and were
attracted to the
negative end of the
magnet.
 Found
atom
the nucleus of the
Gold foil experiment – shot
positively charged Helium at
gold foil to see if atom was
same all the way through
 Most particles when straight
through, some were deflected
 Because H+2 is positively
charged and some were
deflected, he concluded there
must be a positively charged
mass in the atom.
 Atom is mostly empty space.

 Shot
alpha particles,
He+2, at an atomic
nucleus
 Found that mass
changed, but not
the charge.
 Had to be a particle
– the same mass as
a proton with no
charge

Essential discovery for
the fission of uranium

Necessary for nuclear
energy
Determined what
keeps electrons in
orbit around nucleus
 Proposed that
electrons have a set
amount of energy
putting them in
different energy
levels, orbits, around
the nucleus
 Electrons can change
energy levels; higher
levels are further
from the nucleus

 Atom
-a basic unit of matter that consists of
a dense, central nucleus surrounded by a
cloud of negatively charged electrons.
 Element - a pure chemical substance
consisting of one type of atom distinguished
by its atomic number


Found on the periodic table
Atomic Number – the number of protons in the
nucleus of an atom
 Isotope
– atoms of the same element with
different numbers of neutrons
 Protons
– positively charged
particles

Found in the nucleus
 Neutrons
– neutral particles
same average mass as protons
 Found in the nucleus

 Electrons
– negatively
charged particles
Found in orbits around the
nucleus
 Very small mass


Mass Number – the total number of protons and neutrons
Protons and neutrons have the same mass and are found in the
nucleus of an atom
 Electrons are approximately 2000 times smaller than protons
and neutrons
 The nucleus of an atom is relatively heavy since it holds most
of the atom‟s mass




Atomic Mass = the average mass of all of the isotopes of an
element





Protons and neutrons
Different isotopes have different mass numbers
Because electrons are so small, rounding the atomic mass will
give you the average mass number.
Atomic number = protons = electrons
Protons + neutrons = mass number
Mass number – atomic number = neutrons
Mass number – protons = neutrons
A
way to organize the 118 known elements
based on increasing atomic number
 Also organized based on other trends

To be discussed later
 Developed
by Dmitri Ivanovich Mendeleev
(with historical help from many others)



Late 1800‟s
First to develop a table that predicted
undiscovered elements based on gaps in size
Also first to recognize other trends in the table.
 Atomic
11
Na
Sodium
22.99
Number
 Symbol
 Name
(if included)
 Atomic Mass
 Shorthand
for the
box on the
periodic table
Mass Number
56
26Fe
Symbol
Atomic Number
 Using
your periodic
table…
How many protons
does helium (He)
have?
 How many neutrons
are in an atom of
carbon (C)?
 How many electrons
are in an atom of
lithium (Li)?

Name
Symbol Atomic
#
Mass #
Protons Neutrons
Electrons
Phosphorus - 31
5
84
6
36Kr
9
9

When atoms gain or
lose electrons



Cation – positively
charged, lost one or
more electron
Anion – negatively
charged, gained one or
more electron
Charge can be found
on the nuclide
 2713Al+3

If charge is positive, the
atom lost electrons. If
it‟s negative, it gained
them.

For the following,
how many electrons
can be found in the
atom?
 126C+2
 3517Cl 7934Se-2
Name
Nuclide
Atomic #
Mass #
Protons
Neutrons
Carbon - 14
Electrons
10
3 H+
1
6
3
9
2
10
10
 Protons
and Neutrons in an atom
are found clustered together in
the nucleus
 Electrons are found in orbits, or
energy levels, around the nucleus

If electrons move between energy
levels they absorb or emit energy

Moving away from the nucleus requires
energy, moving toward the nucleus
releases energy

Each energy level can only hold a certain number of
electrons

Octet rule – each orbital (energy level) is full once 8
electrons are found in it


Except the first, it only has two electrons
Example:
 Oxygen has an atomic number of 8.


It has 8 protons and 8 electrons.
 Two electrons in its first orbital, and 6 in the second
The number of electrons in the outermost orbital are
considered valence electrons

Helps determine the reactivity of the element


The closer to a full or empty orbital the more reactive the atom
 Ie Na only has one valence electron, it is very close to empty.
It is highly reactive.
Ions only gain or lose valence electrons

Atoms want full valence orbitals

They want to be like the closest Nobel Gas to them

The last column of the periodic table
This makes them the least reactive.
 The number of electrons they are likely to gain
or lose is based on emptying or filling an orbital

This is the Oxidation Number of the element
 The elements on the left side of the periodic table
are more likely to lose electrons, elements on the
right side are more likely to gain electrons
 The ones in the middle can become either anions or
cations
 Example: Na will likely be a ____ ion with a charge
of ___

 Use
the Bohr‟s model (planetary model) to
draw atoms


Protons, neutrons are found in the nucleus
Electrons are found the in energy levels
surrounding the nucleus

Don‟t go to the next energy level until the one before
it is full
-
Example: Be
Atomic #?
Mass #?
Protons?
Neutrons?
Electrons?
4p
5N
-
 Determine
the number of subatomic particles
for the following ions and atoms, and draw
the Bohr‟s model of the atom
Atom/
Ion
Li
S
Ne
Cl-1
Mg+2
Name
Nuclide
Atomic
#
Mass #
Protons
Neutrons
Electrons
Valence
Electrons
Draw
the following
atom/ions
+2
Ca
F
Ar
 The
Bohr‟s model is a very simple way to
draw atoms

Through technology we‟ve determined where
electrons are arranged within at atom or
molecule
 Electron
configuration is the arrangement of
electrons in an atom or molecule. They tell
you how many electrons are in each energy
level

Use the periodic table
as a map


Divide it into four parts
There are three parts of the electron configuration to
indicate where the electrons are around the atom


The big number: stands for the energy level
The letter (s, p, d, & f): the shape of the orbital


s – spherical etc
The exponent: the number of electrons in that orbital
Each orbital gets filled before you move on to the next
one.
 Example: O


1s2, 2s2, 2p4
H
 He
 Li
C
N
 Na
 Fe
 Elements
are
organized by
increasing atomic
number
 Also organized into

periods (rows)


Read from left to right
Groups/families
(columns)


All elements in a family
have similar trends
All have the same
number of valence
electrons

Metals






Alkali Metals
Alkaline
Earth metals
Transition
Metals
Inner
Transition
Metals
Metalloids
Non-Metals
Metalloids
 Gasses

 The
organization of the periodic table is not
just based on atomic number
 There are other trends that show up on the
periodic table
 Trend #1 – Atomic Radius



Definition – the average distance from the
nucleus to the outermost electron
As you travel to the left of the periodic table,
the elements have a larger radius
As you travel down the periodic table, the
elements have a larger radius

Trend #2 – Ionization energy




Definition: The amount of energy
required to remove one electron
As you travel to the right across a period,
the ionization energy increases
As you travel up a group, the ionization
energy increases
Trend #3 – Electron Affinity
Definition – the amount of energy gained
when an electron is added to it
 As you travel to the right across a period,
the electron affinity increases
 As you travel up a group, the electron
affinity increases

 Trend



#4 – Electronegativity
Definition – the ability of an element to
attract pairs of electrons in a covalent
bond
As you travel to the right across a
period, electronegativity increases.
As you travel up a group, the
electronegativity increases
 Metallic

Characteristics
Increase with lower valence electrons
and larger atomic radius
 Non-metallic

characteristics
Increase with higher valence electrons
and smaller atomic radius
 Place
the following elements in increasing
order based on each criteria

C, Na, Sr, Al, Ne




Atomic Radius
 Ne, C, Al, Na, Sr
Ionization Energy
 Sr, Na, Al, C, Ne
Electron Affinity
 Sr, Na, Al, C, Ne
Electronegativity
 Ne, Sr, Na, Al, C

Results from a loss of the forces of the nucleus
 Strong nuclear force: a super strong force that acts
between protons and neutrons in the nucleus, binding
them together



This attractive force is stronger than the force that repels
„like charges‟ and that attracts „opposite charges‟
Only acts at extremely small distances (10-15m)
When the nucleus becomes unstable, this
nuclear force becomes unbalanced, radioactive
decay occurs


Tends to occur in atoms with large proton to neutron
ratios
When they break down they emit radiation

Types
 Alpha α
 Beta β
 Gamma γ

Alpha α
 Most common form of radiation
 Alpha radiation consists of fast flying positively charged
particles
 Combination of protons and neutrons



Aka the nucleus of a Helium atom, atomic number 2
Beta β
 Medium strength of radiation
 Beta radiation consists of fast flying negatively charged
particles
 Each beta particle is an electron that is ejected by an atomic
nucleus
Gamma γ
 strongest form of radiation
 Occurs when an atom in an excited state releases energy
 Extremely short wavelength, much more energetic than visible
light
 Gamma radiation carries lots of electric charge and no mass
 When
atoms/elements break down, they
become another element
 This process emits radiation
 Types


Alpha particle emission
Beta particle emission
 Alpha

When an atom breaks down and emits an alpha
particle




particle emission
A 4He nucleus (2 neutrons and 2 protons)
Occurs with massive nuclei that have too large of
a neutron to proton ratio
To determine the products of alpha particle
emission, you subtract a He nucleus
Example:
U  23190Th + 42He
Just like a math equation: the top numbers have to
be equal and the bottom numbers have to be equal.
They symbol goes with the atomic number found on
the bottom
 23592

 Beta

When a atom breaks down and emits a beta
particle



Particle Emission
Energy converts a proton into a neutron (β+) and emits
a positive charge
OR energy can convert a neutron into a proton (β-) and
emits a negative charge
Example
Th 
 23190
231 Pa
91
+ 0-1e

The time required for half of the atoms in a sample of a
radioactive isotope to decay
Different isotopes decay at different rates
 The longer the half life, the greater the stability


Example




Radium-226 has a half life of 1620years
This DOES NOT mean that in 3240 years it will be gone!
This does mean that after another 1620 years, half of the
remaining half will be gone, leaving ¼ of the original sample
The half life of the sample continues like this: ½ will remain,
¼ will remain, 1/8 will remain, 1/16 will remain, and so on



Each half life cycle leaves 1/(2n) of the original
Half lives are VERY consistent and not affected by
environmental conditions
Half-lives can be measured by a radioactive detector and
by measuring how much decay occurs per year.

A 100 gram sample of 13C decays to 25 grams in
20.6 seconds. What is its half-life?
Original = 100g
Left = 25g
25/100 = ¼
 This means it when through 2 half life cycles.





20.6 seconds / 2 = 10.3s
The half life of 258Md is 2,800 years. If there are
33g of the sample left after 1,400 years, how
many grams were in the original?



Half life = 2,800 years
Time passed = 1,400 years
½ of a half life has passed, so


Only ¼ of a sample has decayed, so ¾ is left
33/(3/4) = 44
 There
are 5.0g of 210Bi left after 30.45 days.
How many grams were in the original sample
if its half-life is 6.09days?





Original = ?
Left = 5.0g
Half life = 6.09days
Time passed = 30.45 days
How many half life cycles?



30.45/6.09 = 5
5 cycles means 1/(25) of the sample is left
 1/32 left
5.0g /(1/32) = 160g

Nuclear fusion

Taking two atoms and making a new
one plus neutrons

Occurs mostly in lighter atoms
Releases radiant energy as gamma
radiation
 Found in stars and the hydrogen bomb


Nuclear Fission

Splitting atoms into two new ones


Creates two new smaller nuclei and
releases neutrons
Generally only occurs in heavier atoms
Releases gamma radiation
 Method behind nuclear power and
nuclear weapons
