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The structure of the Atom
Chemistry chapter 4
What makes up Matter?

In ancient times, people sought to organize the
world around the fundamental elements, such
as earth, air, fire, and water
Democritus (460–370 B.C.)
• The first person to propose the idea that matter
was not infinitely divisible, but made up of
individual particles called atomos.
Aristotle (484–322 B.C.)
 Disagreed with Democritus because he did not
believe empty space could exist.
 Aristotle’s views went unchallenged for 2,000
years until science developed methods to test
the validity of his ideas.
Alchemy (next 2000 years)
• Mixture of science and mysticism.
• Lab procedures were developed, but alchemists did not perform
controlled experiments like true scientists.
Daltons Atomic Theory 1808
All matter is composed of extremely small
particles called atoms
 Atoms cannot be subdivided, created,
or destroyed


Atoms of a given element are identical
in size, mass, and other properties;
Atoms of different elements differ in size,
mass, and other properties
Daltons Atomic Theory 1808
In chemical reactions, atoms are combined, separated, or rearranged
Compounds
 Atoms of different elements combine in simple whole-number ratios
to form chemical compounds
 Compounds are formed when atoms of elements combine
 Molecules are composed of atoms in definite proportions
 Law of Constant Composition
Billiard Ball Model
 atom is a uniform, solid sphere
Discovery of the Electron
Sir William Crookes (1879)
 Used evacuated tubes containing gas at low pressure
 Gasses encounter electricity in a sealed environment
• When an electric charge is applied, a ray of radiation travels from
the cathode to the anode, called a cathode ray
• Cathode rays are a stream of particles carrying a negative charge
Discovery of the Electron

Crookes discovered cathode rays had the following
properties:
 Travel in straight lines from the cathode
 Cause glass to fluoresce
 Are deflected by electric fields and magnets to
suggest a negative charge
Discovery of the Electron

In 1897, J.J. Thomson used a cathode ray tube (CRT) to deduce
the presence of a negatively charged particle.
 Used a fluorescent screen in CRT to measure deflection of beam
 Found that all particles in the beam had same charge and mass
 Proved that the beam, using magnets, was negatively charged
particles called Electrons
Problem with Negative model?



J.J. Thomson realized problem
 Using CRT again with hydrogen in tube
 Found 2nd beam of particles traveling
toward cathode. Theorized beam was
made of Protons
Modified Daltons atom model
 Plum-pudding model
 Atom is sphere with small electrons
embedded in a positively charged mass
Now model can be neutral, negatively or
positively charged atom
 An ion is an atom that has either lost or
gained electrons
 Cation – Positive ion
 Anion – Negative ion
Charge and Mass of Atom
Millikan (1909 - 1920)

Conducted his “oil drop” experiment which allowed him to
measure the charge on an electron

Difference in mass between Proton and Electron?

Proton approximately 2000x more than electron

Why important?
1.
Mass of proton is much much greater than an
electron
2.
The mass of an atom is know, If the proton accounts
for ½ the mass, what is the remaining “stuff”??
Rutherford
Credited with discovering that most of the atoms is made up of
“empty space”. In 1909 he conducted the “gold foil” experiment.
 Gold Foil thickness: 0.00006 cm
• By aiming the particles at a thin sheet of gold foil, Rutherford
expected the paths of the alpha particles to be only slightly altered
by a collision with an electron.

Rutherford Gold Foil Experiment

http://micro.magnet.fsu.edu/electromag/java/rutherford/
Rutherford

He established that the nucleus was:
 Very dense
 Very small and positively charged
 Also assumed that the electrons were located outside the
nucleus

Rutherford Proposed Planetary Model or Nuclear Model

Predicted the existence of the neutron in 1920

James Chadwick discovered the neutron in 1930’s. He found these
uncharged particles with essentially the same mass of the proton
Rutherford-Bohr Model
Niels Bohr (1922)
 Proposed improvements to Rutherford Atomic Model. For this
reason the planetary model of the atoms is sometimes called the
Rutherford-Bohr model
 Bohr added the idea of fixed orbits, or energy levels for the
electrons traveling around the nucleus
 His atomic model has atoms built up of successive orbital shells
of electrons
 Bohr Model or Orbital Model
Electron-Cloud Model


The charge-cloud model, also
called the quantum-mechanical,
does not attempt to describe the
path of each electron in a fixed
point.
Computers can calculate the
points in space that an electron
has the highest probability of
occupying.
• Scientists have determined that
protons and neutrons are
composed of subatomic particles
called quarks.
Chemistry Vocabulary
Atoms – Basic building block of
matter
 smallest unit of an element
that retains properties of that
element
Atom consists of 3 particles
 Protons: positive charge (+)
 Neutrons: neutral charge
 Electrons: negative charge (-)
Atom divided into 2 parts
1) Nucleus – center of atom,
contains protons and neutrons
 nucleus has positive charge
2) Electron cloud – region around
nucleus that contains electrons
 probable location where
electrons will be located
 this outer region is much
larger than nucleus
 diameter of nucleus =
1/100,000 of electron cloud
Electron cloud has several layers –
like an onion
 Electrons near the nucleus have
low energy
 Electrons further away have
higher energy
Chemistry Vocabulary



Symbol
Each element is identified by a symbol and
Atomic Number
Atomic number
 Atomic number is unique for each element
 # of protons in nucleus = atomic number
 number of protons for an element
NEVER changes
 # of electrons = atomic number
 As long as atom is neutral
Each element has a mass number
 sum of the number of protons and neutrons
in an atom
 Periodic table lists the Average mass
number

Mass number = # protons + # neutrons
8
O
Oxygen
15.99
Name
Chemistry Vocabulary
Examples of number of protons, electrons:
 Oxygen
 8 protons, 8 electrons
 Gold
 79 protons, 79 electrons
Difference between neutral atoms and ions:
 Neutral Atoms
 Same number of electrons & protons
 This what you read off the periodic table
 Ions
 Atoms that have either lost or gained electrons
 Positive ions are cations
 Negative ions are anions
Ca
Ca
O
+2
-2
Chemistry Vocabulary


Two atoms of the same element can have differing numbers of
neutrons. These are called isotopes
 Different atomic masses
Hydrogen has three isotopes:
 Hydrogen with 1 proton and 0 neutrons
 Deuterium with 1 proton and 1 neutron
 Tritium with 1 proton and 2 neutrons
Remember!!
Periodic table lists the
Average mass number
Chemistry Vocabulary

Isotopes written 2 different ways
1. Chlorine – 35
 means it has mass of 35
 How many neutrons does it have?
 From periodic table atomic number = 17 protons
 # neutrons = 35 (mass #) – 17 (atomic #)
 # neutrons = 18
2.
35
17
Cl
EXAMPLE:
How many protons, neutrons and electrons are found in an
atom of
133
55
Cs
Atomic number = protons and electrons
There are 55 protons and 55 electrons
Mass number = sum of protons and neutrons
133 – 55 = 78
There are 78 neutrons
Radioactive decay




Radionuclides undergo radioactive decay, meaning they eject small
nuclear fragments and sometimes high energy electromagnetic radiation
as well
This is called Radioactivity
 About 50 of the approximately 350 naturally occurring isotopes are
radioactive
 Few exist in nature—most have already decayed to stable forms
 Naturally occurring radiation consists principally of alpha, beta, and
gamma radiation
The goal for radioactive decay is for unstable nuclides to give off
radiation to become more stable
Nuclear reactions can change one element into another element
Alpha Radiation
• Alpha Radiation is made up of positively charged particles
called alpha particles

Each alpha particle contains two protons and two neutrons and
has a 2+ charge.
4
 Rapidly moving helium ions (no electrons)
2
He α
alpha particle
radioactive isotope
neutron
proton
Alpha Radioactive Decay (cont.)
•
The figure shown below is a nuclear equation showing the
radioactive decay of radium-226 to radon-222.
• The mass is conserved in nuclear equations.
Beta Decay
•
Beta radiation is radiation that has a negative charge and
emits beta particles
• Each beta particle is an electron with a 1– charge.
0
1
e
Gamma Radiation
• Gamma Rays are high-energy radiation with no mass and are
neutral.
γ

Pure energy; called a ray rather than a particle


Form of electromagnetic radiation
Gamma rays account for most of the energy lost during
radioactive decay.
241
96
Cm 
237
94
Pu  He  
4
2
Review
Conclusions from the Study of the Electron



Cathode rays have identical properties regardless of the element
used to produce them. All elements must contain identically charged
electrons.
Atoms are neutral, so there must be positive particles in the atom to
balance the negative charge of the electrons
Electrons have so little mass that atoms must contain other
particles that account for most of the mass
Modern Atomic Theory
Several changes have been made to Dalton’s theory.
Dalton said:
 Atoms of a given element are identical in size, mass, and other
properties; atoms of different elements differ in size, mass, and
other properties
Modern theory states:
 Atoms of an element have a characteristic average mass which is
unique to that element.
Dalton said:
 Atoms cannot be subdivided, created, or destroyed
Modern theory states:
 Atoms cannot be subdivided, created, or destroyed in ordinary
chemical reactions. However, these changes CAN occur in nuclear
reactions!
One atomic mass unit (amu) is defined as 1/12th the mass of a
carbon-12 atom
•One amu is nearly, but not exactly, equal to one proton and
one neutron
•Carbon 12 is assigned an atomic mass of 12.00 g
•12.00 is one atomic mass unit
The number of protons and neutrons in an atom is its mass number.
• Atomic numbers are whole numbers
• Mass numbers are whole numbers
• The atomic mass is not a whole number
• The atomic mass of an element is the weighted average
mass of the isotopes of that element.