Download atom

Chapter 3
Atoms: The
Building Blocks
of Matter
Section 3-1
The Atom: From
Philosophical Idea to
Scientific Theory
Chemical Reaction
A
reaction is a change in which one
or more substances are converted
into different substances.
 The reactants are the starting
materials.
 The products are the ending
materials.
The Law of the
conservation of mass
 Matter
can neither be created
nor destroyed in ordinary
chemical reactions.
 This means that mass of
reactants equals mass of
products in a closed system.
Law of Conservation of
Energy
 Energy
cannot be created or
destroyed.
 It can be changed from one
form to another.
Law of Definite
Proportions
A chemical compound contains
the same elements in exactly the
same proportions by mass
regardless of how large a sample
or source of the compound.
Example:



NaCl
100. g sample
39.3% is sodium
60.7% is chlorine
25.0 g sample
39.3% is sodium
60.7% is chlorine
Law of Multiple Proportions
If two or more different compounds are
composed of the same two elements,
then the ratio of the second element
combined with a certain mass of the first
element is always a ratio of small whole
numbers.
Example: CO and CO2
In CO
1.00 g of C
1.33 g of O
Ratio of oxygens
1.33/2.66 = 1/2
In CO2
1.00 g of C
2.66 g of O
Concept of Atom
1. Democritus
Greek philosopher (400 B.C.)
called atomos which means
“indivisible”.
Aristotle thought matter was
continuous. Atoms were ignored
for 2000 years.
Concept of Atom
2. John Dalton
English School Teacher (1800’s)
Proposed explanation of the laws
that existed at that time.
Called Dalton’s Atomic Theory.
Dalton's Atomic Theory
a) All matter is made of atoms.
b) All atoms of a given element are
identical in size, mass and
properties.
c) Atoms cannot be subdivided,
created or destroyed.
Dalton's Atomic Theory
d)Atoms of different elements
combine in simple whole
number ratios to form chemical
compounds.
e) In chemical reactions, atoms
are combined, separated or
rearranged.
Dalton's Atomic Theory
Dalton turned Democritus’ idea into a
theory that can be tested.
Testing however showed that his
theory needed to be modified.
Section 3.2
The Structure
of the Atom
Atom
The smallest unit of an element that
maintains the properties of that
element
Parts of the Atom
1)The nucleus
Small region that is dense.
Contains protons and neutrons
2)The electron cloud
The region surrounding the
nucleus of an atom where
electrons are likely to be found.
Electron
 Negatively
charge particles found in
the electron cloud
 Mass: 9.109 x 10-31 kg
 Electric Charge: -1.6 x 10-19 coulombs
Proton
 Positive
charge particles found in the
nucleus
 Mass 1.673 x 10−27 kg (or about 1800
times the mass of an electron)
 Electric charge of +1.6 x 10−19 coulomb
Neutron
 Neutral
particles found in the nucleus
 Mass: 1.675 x 10-27 kg (slightly more
than a proton)
 Electric charge: no net electric
charge
Thomson's Cathode–Ray
Tube Experiment
 Time
frame was 1897
 The discovery was the electron
 The scientist was J. J. Thomson
Thomson's Cathode–Ray
Tube Experiment
Thomson's Cathode–Ray
Tube Experiment
Concluded that the negatively charged
particles in the cathode–ray tube were
fundamental particles which are present
in all matter.
Called the tiny particles electrons.
J.J. Thomson’s Atomic
Theory

3. Plum Pudding Theory
Atoms of elements are divisible.
Electrons are small particles
that can leave the atom.
Electrons are spread evenly
through a positive charge in
the rest of the atom like
plums in pudding.
Millikan's Oil Drop
Experiment
In 1909, American physicist Robert
Millikan determined the charge of the
electron with his famous oil drop
experiment.
Millikan's Oil Drop
Experiment
Rutherford's Gold Foil
Experiment
 Time
frame: 1911
 New Zealand
 Discovery: Nucleus is dense and
does not take up much space
Rutherford's Gold Foil
Experiment
Problem: Mass and charge were
believed to be uniformly distributed
throughout an atom.
Rutherford's team expected positively
charged alpha particles to pass
through a sheet of thin gold foil with
only a slight deflection.
Rutherford's Gold Foil
Experiment
Rutherford's Gold Foil
Experiment
One in 8,000 particles, however,
bounced back toward the source.
Why?
Rutherford's Gold Foil
Experiment
Rutherford's Gold Foil
Experiment
Conclusion: The rebounded particles
must have encountered a very densely
packed bundle of matter with a
positive electric charge.
After much thought, decided nucleus
existed.
Niels Bohr Atomic Theory
Student of Rutherford
 4. Called Planetary theory
Electrons travel around the
nucleus much like planets
travel around the sun.
Nucleus is very small, dense,
positively charged object.

Nuclear Forces
They are the interaction that binds
protons and neutrons, protons and
protons, and neutrons and neutrons
together in a nucleus.
Nuclear Forces
In stable nuclei, the nuclear force is
stronger than the electric repulsion
between the protons alone and holds
the nucleus together.

In unstable nuclei, the number of
protons and neutrons is not balanced,
causing the nuclear force to weaken
and the nucleus to decay.

Forces in Atoms
The forces acting between atomic and
subatomic particles.
 Gravitational
 Electromagnetic
 Strong nuclear
 Weak nuclear
Size of Atoms
Radius of an atom is the distance
from center to the outer portion of
the electron cloud.
 Use picometer (pm = 1 x 10-12 )
 Atomic radii range from 40 to 270
pm
 If nucleus were marble, atom would
be a football field.

Section Three
Counting
Atoms
Atomic Number
Z
- The number of protons in the
nucleus of an atom.
 The atomic number is the same for all
atoms of an element

Atomic No. = No. of protons = No. of
electrons
Mass Number
A - The sum of the numbers of protons
and neutrons in the nucleus of an atom.
 Because electrons have very little
mass, the mass number is very
close to the average atomic mass of
the element.
 Mass No. = no. of neutrons + no. of
protons

Isotopes and Nuclides
 Isotopes
are atoms of the same
element that have different masses.
 Nuclide is the general term for any
isotope of any element, based on the
number of protons and neutrons in its
nucleus.
Hydrogen Isotopes
Atomic No.
No. of
Mass No.
(No. of Protons) Neutrons
Protium
(Hydrogen-1)
Deuterium
(Hydrogen-2)
Tritium
(Hydrogen-3)
1
0
1
1
1
2
1
2
3
Nuclides (Isotopes)
Helium-3
Helium-4
Sodium-23
Carbon-12
Nuclear
Symbol
No. of
Protons
6
No. of
Electrons
8
No. of
Neutrons
8
9
Nuclides (Isotopes)

Relative Atomic Mass

Mass of an atom compared to a
standard.

Standard is atomic mass of carbon12 atom (12 amu)

1 atomic mass unit = 1/12 the mass
of a carbon-12 atom (1 amu)
Average Atomic Mass
 The
average atomic mass is the
weighted average of the atomic
masses of the naturally occurring
isotopes of an element.
 Atomic mass on periodic table is
average atomic mass.
Average Atomic Mass
Example:
25 marbles weigh 2.00 g each
75 marbles weigh 3.00 g each
25 x 2.00g = 50.0g
75 x 3.00g = 225 g
100 marbles =275 g
Average weight for each marble is
2.75 g
The Mole
A
mole is the amount of a substance
that contains as many particles as
there are atoms in exactly 12 g of
carbon–12.
 A mole is a quantity like a dozen.
Avogadro's Number
The number of particles in 1 mole
equals 6.022 x 1023 per mole
Molar Mass
 Molar
mass is the mass in grams of
one mole of a substance
 A mole of any other kind of atom or
molecule will always contain the
same number of particles.
Comparison of Terms
1 mole of Iron atoms = 55.85 g
 1 mole of Iron has 6.022 x 1023
atoms


1 atom of Iron = 55.85 amu
1 atom of Iron-55 = 55 amu
 1 mole of Iron-55 = 55 g

The End
(of terms that is.
We still have the
math problems to
go.)