Download Section 6.2 Notes - oologah.k12.ok.us

4-1 Notes
Defining the Atom
Early Models of the Atom
 All matter is composed of atoms
 Atoms are the smallest particles of an
element that retains their identity in a
chemical reaction
 Greek philosopher Democritus
(460B.C.-370B.C) was the first to
suggest the existence of atoms
He believed that atoms were indivisible
and indestructible
His ideas did not explain chemical
behavior
Dalton’s Atomic Theory
2000 years later John Dalton used
experimental methods to transform
Democritus’s ideas on atoms into a
scientific theory
Dalton studied the ratios in which
elements combine in chemical
reactions.
He developed the atomic theory
which had 4 main points
Dalton’s Atomic Theory
1. All elements are composed of tiny indivisible
particles called atoms.
2. Atoms of the same element are identical. The
atoms of any one element are different from
those of any other element.
3. Atoms of different elements can physically mix
together or can chemically combine in simple
whole number ratios to form compounds.
4. Chemical reactions occur when atoms are
separated, joined, or rearranged. Atoms of one
element are never changed into atoms of
another element as a result of a chemical
reaction
Sizing up the Atom
 Atoms are extremely small
The radii of most atoms fall within the range of
5 x 10-11m to 2 x 10-10m.
 Despite their small size, individual atoms are
observable with instruments such as scanning
tunneling microscopes.
 By using this microscope, individual atoms can be
moved around and arranged in patterns
 This technology is important in future
applications in medicine, communications, solar
energy, and space exploration
4-2 Notes
Structure of the Nuclear Atom
Subatomic Particles
Dalton’s atomic theory is still
accepted today except for 1 thing,
that atoms are not indivisible.
An atom can be broken apart into 3
different particles – electrons,
protons, and neutrons.
Electrons
 Discovered by English physicist J.J.
Thompson
 His experiment included passing electric
current through gases while they were
sealed in a tube (fig. 4.4)
 This produced a cathode ray (glowing
beam) that traveled from the cathode to
the anode
 Since the ray was attracted to the + plate
and repelled by the – plate, he concluded
that the particles had a negative charge
 Named them electrons
 Robert Millikan later concluded that each
electron carried a -1 charge.
Protons
 There were four things that were known to be
true about atoms
 atoms have no net electric charge
 electric charges are carried by particles of matter
 electric charges always exist in whole-numbered ratios
 when a given number of negatively charged particles
combine with the same number of positively charged
particles, an electrically neutral particle is formed
 Therefore, there had to be a positive particle in the
atom.
 In 1886 Eugen Goldstein found evidence of a positive
particle he named a proton
 He also determined the mass was 1840 times that of
an electron
Neutrons
In 1932, English physicist James
Chadwick (1891-1974) discovered
the neutron
Neutrons are subatomic particles
with no charge but a mass
approximately equal to that of a
proton
The Atomic Nucleus
 Based on experiments Ernest Rutherford
performed in 1911 he concluded the following
 The majority of an atom is empty space.
 All the positive charge & most of the mass is
concentrated into a small region
Called the nucleus
Composed of protons and neutrons
 The electrons are distributed around the
nucleus and occupy almost all the volume of
the atom.
 Rutherford’s model was an improvement of
Thompson’s but it turned out to be
incomplete
4-3 Notes
Distinguishing Among Atoms
Atomic Number
Elements are different because they
contain different numbers of protons
The atomic number = the number of
protons in an element
Find the # of protons for H, C, K
Mass Number
Protons + neutrons = the mass
number
The number of neutrons in an atom is
the difference between the mass
number and the atomic number
# of neutrons = mass # - # of protons
Find the # of neutrons for C, Nb, Ag,
The composition of an atom is often
written in shorthand notation using
atomic number and mass number
Ex. Gold - 197
Isotopes
Isotopes are atoms that have the
same number of protons but different
numbers of neutrons.
Because isotopes of an element have
different numbers of neutrons, they
also have different mass numbers
Hydrogen has three isotopes
What are they?
Atomic Mass
Since the actual masses of elements
are extremely small, they are difficult
to work with.
Therefore, we find the relative masses
of atoms using a reference isotope as
a standard
The isotope chosen was Carbon-12,
and it was given a mass of exactly 12
atomic mass units (amu)
An atomic mass unit is defined as
1/12th of the mass of carbon 12
Atomic Mass
 Carbon 12 has 6 protons and 6 neutrons,
therefore each proton or neutron is 1 amu
 Due to the existence of isotopes, you can’t
accurately predict the atomic mass of any
element.
 Most elements occur as a mixture of 2 or
more isotopes, each with a specific mass
and natural percent abundance
 The atomic mass of an element is the
weighted average mass of the atoms in a
naturally occurring sample of the element
Calculating Atomic Mass
To calculate the atomic mass of an
element:
1. multiply the mass of each isotope by
its natural abundance
2. Add the products
The sum is the weighted average
mass of the atoms
Sample Problem 4.2
The Periodic Table
A periodic table is an arrangement of
elements based on a set of repeating
properties
Each element is identified by its
symbol in a square
Each horizontal row of the periodic
table is called a period
Each vertical column of the periodic
table is called a group or family
Element of the same group have the
same physical and chemical properties