Download History of the Atom

History of the Atom
A Look at Some Famous Scientists
History and Experiments
•
•
•
Comes from the time
around 400 B.C.
Ancient Greek philosopher
He performed no
experiments during his time
Conclusions or Discoveries
Came up with the word
“atomos”
• He said that the atom was a
small indestructible
molecule
• It was the smallest piece of
matter
•
DEMOCRITUS
History and Experiments
Conclusions or Discoveries
384-322 B.C
 Greek Philosopher
 Used reason rather than
experiments



Stated that all matter was
made up of fire, wind, earth
or water
Views were unchallenged
for 2000+ years
Aristotle
History and Experiments
Conclusions or Discoveries
1.
•
English school teacher who
developed the Atomic Theory in
1803
•
Studied the atom in 1803
•
Studied the masses of elements and
compounds and their ratios
3.
•
Used experimental data to support
his theory
4.
2.
5.
All matter is made of
indivisible and indestructible
atoms
Atoms of the same element
are identical
Atoms of different elements
are different
Atoms of different elements
combine in whole number
ratios and form compounds
Chemical reactions consist
of the combination,
separation, or
rearrangement of atoms.
John Dalton
1. Law of Conservation of Mass

Mass (matter) is neither created or destroyed,
just rearranged.

Mass of reactants = mass of products
C
+ O2
12 g
2(16) g
CO2
44 g
History and Experiments
•
•
Researched the atom back
in 1897
He used a cathode ray tube
and observed the cathode
ray is attracted to a positive
(+) charge
Conclusions or Discoveries
Discovered electrons
• Plum pudding model of the
atom
• (Chocolate Chip Cookie)
•
J.J.Thompson
Observations
Inferences
When using a magnetic field, the
rays bent away from the field
The rays must have a charge
When using electric fields the rays
were attracted to the (+) field.
Support for rays having (-) charge.
When a small paddle wheel was
placed in the path of the cathode
rays the wheel was set in motion.
Rays must consist of tiny particles
that have mass (electrons).
Cathode Ray Tube
History and Experiments
•
•
•
Some of the most
important discoveries
related to the atom were
completed by Rutherford
His work was completed in
1908
Gold Foil Experiment
Conclusions or Discoveries
Discovered the nucleus
• The nucleus is very small
and dense
• The atom is mainly empty
space
•
Ernest Rutherford
James Chadwick

In 1932, James Chadwick discovered the neutron, a
neutral subatomic particle with relatively the same mass
as a proton.

The masses of all atoms except hydrogen were known
to be greater than the combined masses of their
protons and electrons.

Theorized that there must a 3rd particle

Neutron – neutrally charged particle
(no charge)

Nucleus – made up of protons and neutrons
History and Experiments
•
•
Conclusions or Discoveries
Bohr’s research on the
atom was completed in
1913
Still highly relevant today
•
Electrons orbited the
nucleus
Niels Bohr
History of the Atom - Timeline
1766 – 1844
Antoine Lavoisier
Thomson
makesJ.J.
a substantial
the
number discovers
of contributions
electron
and
to the
field of
proposes the
Chemistry
Plum Pudding
Model 1871
in 1897
– 1937
Niels Bohr
proposes
the Bohr
Model in
1913
1887 – 1961
James
Chadwick
discovered
the neutron
in in 1932
1700s
1800s
1900s
460 – 370 BC
0
Democritus
proposes
the 1st atomic
theory
1743 – 1794
Erwin
John Dalton
Ernest Rutherford
Schrodinger
proposes performs
his
the Gold Foil
describes
1891 – 1974
atomic theory
Experiment
in
in 1909
the electron
1803
cloud in 1926
1885 – 1962
Click on picture for more information
1856 – 1940
Definitions

Atom
◦ The smallest part of an element that upholds
the chemical properties of that element

Element
◦ A substance that can’t be broken down into
simpler chemical substances.
◦ What’s smaller: an atom or element?
The 3 Subatomic Particles
Symbol
Charge
Location
Mass (amu)
Electrons
e-
negative
Electron Cloud
Protons
p+
positive
Nucleus
1
Neutron
n0
Nucleus
1
neutral
1/1836 (= zero)
PERIODIC TABLE OF ELEMENTS
•
Just like the 26 letters of the
alphabet combine to form all
the words in the English
language, the 100 or so
elements combine to form
everything that exists in the
world
•
About 90 of the elements on
the PT are found naturally in
nature, the rest have been
created in a laboratory
•
Any material made of one type
of atom is classified as an
ELEMENT
PERIODIC TABLE OF ELEMENTS
•
The Periodic Table is organized into
ROWS and COLUMNS
•
Each vertical column is called a group or
family
•
Elements within the same group have
similar properties
• Au, Ag, Cu
•
Each horizontal row is called a period
•
Properties of the elements gradually
change when you move through a period
• Elements get smaller when you
move from LEFT to RIGHT
How to Read the Periodic Table

Each element is
designated by its atomic
symbol
◦ Some symbols do not
match actual element name
because they were named
under their Latin name

The first letter of an
atomic symbol is
capitalized and the
second letter is
lowercase
How to Read the Periodic Table
The atomic number is
the special I.D. of the element

◦ It can tell you how many protons are in the
element
Protons have an electric charge of +1
 Electrons are equal to the number of
protons in an element

◦ Electrons have an electric charge of -1
◦ Atomic # = # Protons = # electrons
How to Read the Periodic Table

The atomic mass of an element is the
number at the bottom of the square

Atomic mass is the sum of the protons and
neutrons in the atom
◦ Neutrons have an electric charge of zero
◦ Atomic mass = protons + neutrons

The nucleus of an atom contains all the
protons and neutrons
Let’s practice
ISOTOPE
MASS #
ATOMIC #
35
17
7
3
# PROTONS
# NEUTRONS # ELECTRONS
17
35-17
18
17
3
7-3
4
3
Isotopes

Atoms of the same element that contain
different numbers of neutrons are called
Isotopes.
We identify isotopes by their mass
number. The mass number is the total
number of protons and neutrons it
contains.
 Mass # = # Protons + # Neutrons

Isotopes
One way to express isotopes
 Mass Number
◦ OVER
 Atomic Number

The total number of neutrons
in an isotope is to subtract its
atomic number from the mass
number.
Mass number
-Atomic Number
Number of Neutrons
Examples
How many Protons?
 How many Neutrons?

◦ 6 protons, 6 neutrons
How many Protons?
 How many Neutrons?

◦ 6 protons, 8 neutrons
Atomic Mass
Atomic mass is the average atomic mass
of an element and its various isotopes.
 Measured in Atomic Mass Units (amu)

◦ Example: Imagine 100 students take a test
worth 100 points. If all students scored an 80
then the average would be an 80%. But what
would happen if one student gets a 100%? It
would raise the average slightly.
◦ Same thing happens with isotopes
Are these elements the
same, different or isotopes
1. Element X has 45 protons and 50
neutrons, while Element Y has 50 protons
and 45 neutrons
 Same, Different, Isotope?

Element A has 13 protons and 15 neutrons,
while Element B has 13 protons and 14
neutrons
 Same, Different, Isotope?

Electromagnetic Spectrum

All types of electromagnetic radiation
share common characteristics:
◦c =λxν
 c - Speed of light (in a vacuum) = 3.0 x 108 m/s
 λ – wavelength
 ν - frequency
Atomic Line Spectrum
Atomic Line/Emission spectrum-set of
frequencies of the electromagnetic wave
emitted by atoms of the element.
 Each element is unique.
Flame gives off colors

Quantum Theory

The quantum theory describes the
probability of finding an electron in a
given region of space called an orbital.
Principal Quantum Number (n)
n = principal quantum number
 Major energy levels
 Larger n, the further away electron is
from nucleus and the higher its energy
 Values of n are whole numbers from 1 
infinity

N= Periods- Horizontal rows
Sublevel (l)
l = sublevel within a major energy
level
 Designates the shape of the orbital
 Values of l are whole numbers from

l = 0 to n – 1
N=2
L= 0 n-1
L=0,1
Sublevel (l) values = orbital type
L value = 0 = s
 “spherical”

L value = 1 = p
 “dumbbell”

Sublevel (l) = orbital type
L value = 2 = d
 “Daisy”

L value = 3 = f
 “Fantastic”

Magnetic quantum number (ml)
Designates the orientation of the orbital
in space
 Values of m are whole numbers:
-L  0  +L
 For example, if l = 1 calculate m
 -1  0  +1 Therefore, m = -1, 0, +1

Spin quantum number (ms)
Describes the direction of an electron’s
spin
 Electrons spin on their axis either
clockwise or counterclockwise
 Only two possible values: +1/2 or -1/2

Quantum Theory Summary
Quantum numbers (n, l, ml, and ms)
attempt to locate where electrons in an
atom “should” be
 Quantum numbers are analogous to the
information give on a ticket stub: Stadium,
Section, Row, and Seat

Practice Problems

n=4
◦ l=
4-1=3
◦ ml=
-3,-2,-1,0,1,2,3
-2,-1,0,1,2
-1,0,1
0
◦ ms=
+1/2, -1/2

0,1,2,3
n=2
◦ l=
2-1=1
◦ ml=
-1,0,1
0,1
 0
◦ ms=

+1/2, -1/2
n=1
◦ l=
1-1=0
◦ ml=
0
◦ ms=
+1/2, -1/2
0
Electron Shells
We identify elements based on how the
electrons are arranged in shells.
 At least 7 shells and each shell can hold a
certain amount of electrons.

Electron Configuration
For full configuration you must have all prior parts
to it as well!
 Fluorine: 1s2 2s2 2p5
 All adds up to 9!


Francium

1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1

All adds up to 87!
s-block
Subshell Orbital blocks
p-block
d-block
f-block
Quick Review
Quick Review
Orbital Notation







Orbital Notation is a way to show how many
electrons are in an orbital for a given element.
They can be shown with arrows.
Take what we just learned with electron
configuration but take it a step further.
Each line can hold up to 2 electrons.
S= ____
P= ____ ____ ____
D= ____ ____ ____ ____ ____
F=____ ____ ____ ____ ____ ____ ____
Pauli Exclusion Principle

A max. of two
electrons can occupy
each orbital and
these electrons must
have opposite spin
Pauli Exclusion Principle
Aufbau Principle
Electrons will occupy
the lowest energy
levels possible
 Aufbau (German) =
“building up”

Hund’s Rule
Electrons enter
orbital's of the same
energy one at a time
before pairing up
 It is more STABLE to
have the max.
number of unpaired
electrons

Hund’s Rule
Mole Conversions

1 mole of any substance has 6.02 x 1023
atoms
◦ Avogadro’s number

1 mole of any substance is equal to its
mass on the periodic table in grams.
◦ Molar mass
Lewis Dot Structure


Depending on what group they are in on the
periodic table determines how many valence
electrons they have.
You cannot double up valence electrons until
each side of the element has an electron
◦ Example:
◦
C
C
Lewis Dot Structure Practice


Try creating Lewis Dot structures for the
first 10 elements!
H
He
Li


B
Be
Lewis Dot Structure Practice
C
N


O
F
Ne
Answers