Download Chapter 5 The Structure of the Atom

Chapter 3
Atoms: The Building Blocks of Matter
Objectives:
Students should be able to:
• Summarize the essential points of Dalton’s
atomic theory.
• Describe the particle theory of matter.
• Use the Bohr model to differentiate among
the three basic particles in an atom
• Compare the Bohr atomic model to the
electron cloud.
Atomic Models
• The Bohr Model
The nucleus is orbited
by electrons. The
electrons are in
different energy levels.
 Early theories of matter
Philosophers (ideas)
1.Democritus (460-370 B.C.)
a. First to propose the idea that matter
was not divided into smaller parts
b. Atomos: Atoms: tiny particles
could not be divided
2. Aristotle (384-322 B.C.)
a. Rejected Democritus
b. believed in 4 elements
air, earth, water and wind
(everything came from these 4
elements)
3. Alchemist
a. Changed metal into gold
b. Tried to prolong life
c. Through their mistakes they
learned chemistry
d. First to do actual experiments
• Dalton’s Model
- Performed a number of experiments that
eventually led the idea of atoms
- Theory:
1. All elements are composed of atoms.
Atoms are indivisible and indestructible
particles.
2. Atoms of the same element are exactly
alike.
3. Atoms of different elements are different.
4. Compounds are formed by the joining of
atoms of two or more elements.
This theory became the foundation of modern
chemistry.
 J.J. Thomson
- Discovered the electron
- Plum-pudding model of the atom
Stated atoms were made from a
positively charged substance with
negatively charged electrons scattered
about
His Experiment
1. Passed an electric current through a gas.
2. As the current passed through the gas, it
gave off rays of negatively charged
particles
3. Passed a magnet over the gas. The beam
bent away or moved towards the magnet.
Conclusions:
1. The negative charges came from within the
atom.
2. A particle of matter smaller than the atom
had to exist.
3. The atom was divisible.
4. Called the negatively particles “corpuscles”
(now called electrons)
5. Since the gas was known to be neutral,
there had to be positive charged particles in
the gas.
• His experiment:
 Rutherford
- Discovered the nucleus and it’s charge,
positive
- His experiment: Gold-Foil experiment
Results:
1. Most of the positively charged particles
passed straight through
2. A few of the alpha particles were slightly
deflected
3. A very small number of alpha particles
hit and bounced back off the foil
Conclusions:
1. Most of matter is empty space
2. Scattered throughout- not everywhereare small areas of positive charge
3. Scattered throughout are areas of very
dense matter
4. He called the center the nucleus; it is
positive and tiny compared to the rest of
the atom
• Rutherford’ model
Chadwick: discovered the neutron
Completing the atom
1. Atomic number (Z)
- tells the number of protons
- in a neutral atom: protons = electrons
2. Mass number (A)
protons + neutrons
3.
A
z x
nucleotide
ex.
40
20
Ca
4. Protons and neutrons make up 99.7% of
the atoms mass
5. Chemical name
Hydrogen
atomic number
1
chemical symbol
H
average atomic mass
1.008
Examples:
 How many protons, neutrons and electrons
in the following:
1.
2.
3.
4.
Properties of subatomic particles
Isotopes
1. atoms with the same number of
protons but different number of
neutrons
2. Identified by the mass
3. Example: Chlorine-35
Chlorine-37
mass number
Examples of neutral atoms and isotopes
Atomic # Mass #
p
n
e
27
Al3+
13
13
27
70
Zn
30
30
70
13
30
14
40
10
30
• Relative Atomic Masses
1.Atomic mass unit (u)
: Is exactly 1/12 the mass of a carbon-12
atom
2. Atomic mass: expression of atomic mass unit
3. Atomic mass unit– weighted average mass
of the isotopes of that element
Average atomic mass formula
(mass)(% of element) + (mass)(% of element) = (total mass)
Examples:
1. A carbon isotope has an abundance of 98.90%
with a amu of 12.0 g and the other isotope has
an abundance of 1.10% with a amu 13.003 g.
What is the average atomic mass of the two
isotopes?
2. An element, E, has an atomic mass of 18.40
and consists of two isotopes:
E-17 with a mass of 16.95 and E-20 with a
mass of 19.35. How much E-20 does this
element contain?
(amu)(%) + (amu)(%) = total mass (100)

Unstable Nuclei and Radioactive Decay
Radioactivity
1. Substances spontaneously emitting
radiation
2. radiation – ray and particle emitted
by
radioactive materials
3. Elements (atoms) can change into
atoms of another element
Radioactive elements
1. Nuclei are unstable
2. Gain stability by losing electrons
3. Radioactive decay – process of
becoming stable by losing energy
Type of radiation
1. alpha radiation (α)
a. Deflected towards (-) charged plate
b. alpha particles have 2p+ and 2n0
c. α particles equivalent to He-4
d. Nuclear equation
(Shows how decay occurs)
226
Ra 
88
4
+
α
2
222
Rn
86
2. Beta radiation
a. Deflected towards the positive plate
b. Beta particles (
c.
)
14
0
14
C 
β + N
6
-1
7
 Half-Life (t1/2)
1. Time required for one-half of a
radioisotope’s nuclei to decay into it
products
2. Examples:
Iron-59 half-life is 44.5 days. How
much of a 2.000 g sample will remain
after 133.5 days.
t1/2 = 44.5 days
half-life