Download THE ATOM

THE ATOM
Objectives: Understand the
experimental design and conclusions
used in the development of modern
atomic theory, including Dalton’s
Postulates, Thomson’s discovery of
electron properties, Rutherford’s
nuclear atom, and Bohr’s nuclear
atom.
Democritus
• Made his discovery
around the year 250
B. C.
• This was the first
discovery about the
atom, the next would
come in another 2000
years.
The First Atom
• Democritus took a sea shell and broke it in
half.
• Than he broke it in half again.
• When the pieces got to small he use a
mortar and pestle to crush the shell.
• He finally believed he got to the smallest
piece possible and called it the ATOM;
which in Greek means INDIVISIBLE.
John Dalton (1766-1844)
A New System of Chemical
Philosophy (1808)
Dalton’s Atom Model
1. All matter is made on atoms; and atoms are
indivisible.
2. Atoms of the same element are all identical.
3. Compounds are formed by a combination of
two or more different atoms and they always
have the same proportion of elements. THE
LAW OF DEFINITE COMPOSITION
4. A chemical reaction is a rearrangement of
atoms and the atoms are neither created nor
destroyed. THE LAW OF CONSERVATION
OF MATTER
J. J. Thomson (1856-1940)
Joseph John Thomson
• English physicist who in 1897 discovered a particle
smaller than the atom ; the electron.
• Particle has a negative charge and is much smaller than
the atom so must come from the inside of the atom.
• Electrons are scattered around the atom like raisins in
pudding. (THE PLUM PUDDING MODEL)
Thomson and Rutherford
Rutherford’s Gold
Foil Experiment
Rutherford’s Gold
Foil Experiment
Ernest Rutherford (1871-1937)
• New Zealand born physicist;
worked in England
• 1911 conducted the “Gold Foil
Experiment” the proved the
existence of a small positively
charged center of the atom.
• Disproved the “Plum Pudding
Model”
• THE NUCLEAR MODEL
• Discovered the proton.
• Thought that the electrons
orbited the nucleus like planets
orbited the sun.
Millikan’s Oil Drop Experiment
• A fine mist of oil
droplets is introduced
into the chamber.
• The oil is ionized by
x-rays.
• The electrons adhere
to the oil drops.
• The value for the
charge of the electron
can be calculated.
Niels Bohr (1885-1962)
• Danish physicist,
produced his model in
1911.
• Saw problems with
Rutherford’s model.
• If electrons “orbit” than
they are changing
direction so they are
accelerating.
• That would require
energy.
The Orbital Model
• Electrons do not
“orbit” but are in
allowable ENERGY
LEVELS.
• When the electrons
stay in these levels,
which are at specific
distances from the
nucleus, they do not
give off energy.
Bright Line Spectrum
• But, if the electron moves
from one level to another
it gives off or absorbs
energy.
• These Bright Line
Spectrums are produced
when the electrons “fall
back” to a lower energy
level and give off energy.
• Every element has a
unique Bright Line
Spectrum.
The Subatomic Particles
THE PROTON
•
•
•
•
•
•
•
p+
positively charged
located in the nucleus
relative mass = 1 atomic mass unit
mass = 1.673 x 10-24 grams
equal to atomic number
number of protons “defines” the atom
The Subatomic Particles
THE NEUTRON
•
•
•
•
•
•
•
•
n0
neutral (no electrical) charge
located in the nucleus
relative mass = 1 atomic mass unit
mass = 1.675 x 10-24 grams
equal to mass number minus atomic number
mass number is protons + neutrons
James Chadwick proposed the existence of the
neutron.
The Subatomic Particles
THE NEUTRON
• Isotopes – different atoms of the same element that
have the same number of protons but different numbers
of neutrons
• some isotopes are radioactive – they emit energy when
the nucleus of the atom breaks down spontaneously
• most radioactive isotopes are not dangerous
• to determine if an isotope is radioactive calculate the
proton to neutron ratio
• if ratio is greater than or less than 1:1 for “small” atoms
the isotope is unstable (smaller than Ca)
• if ratio is greater then 1:1.5 for “large” atoms the isotope
is unstable
The Subatomic Particles
THE ELECTRON
• e- (negative electrical charge)
• located in the electron cloud which is divided into energy
levels, sublevels, orbitals, and spins
• relative mass = 0 atomic mass units
• mass = 9.11 x 10-28 grams
• equal to the number of protons if atom is neutral
• atom becomes a charged ion if electrons are gained or
lost
• positive ion = CATION
• formed by the loss of electron, happens to metals
• negative ion = ANION
• formed by the gain of electron, happens to nonmetals
Location of Electrons
•
•
•
•
•
•
•
•
Energy Levels
Discovered by Niels Bohr
# electrons = 2n2
“n” is the energy level
1st level can hold 2 e2nd level can hold 8 e3rd level can hold 18 e(eight if the outside energy
level)
• 4th level can hold 32 e• (eight if the outside)
• The outside level is called the
valance level and can never
hold more than 8 electrons.
NUCLEAR SYMBOLS
mass number
ion charge
23
+1
Na
p+ = 11
11
n0 = 12
atomic number
e- = 10
name
symbol
calcium
atomic
number
mass
number
ion
charge
20
42
+2
number
of
protons
number
of
neutrons
number
of
electrons
atomic
mass
40.08
19 -1
F
9
10
238 0
U
92
10
10
20.18
name
symbol
potassium
atomic
number
mass
number
ion
charge
19
40
+1
number
of
protons
number
of
neutrons
number
of
electrons
atomic
mass
39.098
amu
15 -2
O
8
18
56 0
Fe
26
22
18
39.948
amu
Objective
• Use isotopic composition to
calculate the average atomic
mass of an element.
Mass Number vs. Atomic Mass
• mass number is given for an individual
atom
• mass number is given in nuclear symbols
• atomic mass is an average mass for all
isotopes for the element
• atomic mass is the number on the periodic
table
• if you round the average atomic mass you
will have the mass number of the most
common isotope
Average Atomic Mass