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Transcript
Early Atomic Theory
5.1
Dalton’s Model of the Atom
5.2
Electric Charge
5.3
Subatomic Parts of the Atom
5.4
The Nuclear Atom
5.5
Isotopes of the Elements
5.6
Atomic Mass
Early Theories on the Structure of Matter
Early models of the atom were developed by the Greeks.
Empedocles proposed matter was composed of four basic
elements: earth, air , water and fire.
Democritus proposed matter was composed of small,
indivisible particles he called atoms.
Atoms could combine in different ways,
giving rise to the diversity of compounds we observe.
Aristotle, an influential philosopher, supported Empedocles’
theory, so atomic theory was not fully accepted until 2000 years
later.
Dalton’s Model of the Atom
Elements are composed of small, indivisible particles called
atoms.
• Atoms of the same element are identical in mass and
size.
• Atoms of different elements differ in their mass and
size.
• Compounds are formed by combining two or more
atoms of different elements.
• Atoms combine to form compounds in simple whole
number ratios.
• Atoms of two elements may combine in different
ratios, leading to formation of different compounds.
Dalton’s Model of the Atom
Atoms are individual particles which are different for
each element.
!
Atoms combine in fixed ratios to form compounds.
Two elements can combine in varying ratios to give
different compounds.
H 2O
H2O2
Most of Dalton’s theory remains valid today.
Dalton’s Model of the Atom
Revisions to Dalton’s Theory
1. Elements can be decomposed under certain
conditions.
2. Not all atoms of the same element have identical
mass. Atoms of the same element with different
atomic mass are called isotopes.
3. Atoms are not indivisible. They are composed of
subatomic particles.
Electric Charge
Properties of Electric Charge
1. Charge may be either positive or negative.
2. Opposite charges (positive and negative) attract
while like charges (i.e. negative and negative)
repel.
3. Charge may be transferred from one object to
another, by contact or induction.
4. The force of attraction between charges (F) is related to
the distance between charges by:
F=
kq1q2
r2
where q1 and q2 are the charges,
r is the distance between charges, and k is a constant.
Subatomic Parts of the Atom
Because atoms are so small, determining the presence
of subatomic particles was very difficult.
A single atom is tiny (diameter of 0.1 to 0.5
nm).
New instruments in the early 1900s permitted detection of these
particles.
A scanning tunneling
microscope (STM)
image shows an array of
Cu atoms.
Subatomic Parts of the Atom
A Crooks tube permits generation of cathode rays,
which are streams of electrons.
A Crooks (cathode) ray tube. The stream of electrons passes between the
electrodes.
The electron beam is deflected by both electric
and magnetic fields, indicating it has charge.
Electrons and Protons
Electrons (e–):
A particle with negative electrical charge
(assigned a relative charge of –1).
Electrons have a very small mass
(9.110 x 10–28 g) and size (<10–12 cm).
Protons (p):
A particle with positive electrical charge
(assigned a relative charge of +1).
Protons have a much larger mass
(~1837 times the mass of an electron).
The Effect of Subatomic Particles
Thomson’s work demonstrated the atom is composed
of smaller, charged particles. Dalton’s theory of the
atom then had to be revised.
Thomson’s Model of the Atom
Electrons are negatively charged particles which are
embedded in a positively charged atomic sphere.
+ charged
sphere
Electrons
Thomson’s “plum p
model of the atom.
The Effect of Subatomic Particles
Atoms can become ions by gaining or losing electrons
from this sphere.
Electrons are lost
from atoms to give
cations.
Electrons are
gained from
atoms to
give anions.
Neutrons
The last subatomic particle was discovered by Chadwick in
1932.
Neutrons (n): A particle with no electrical charge.
Neutrons have a mass similar to that of a proton.
Summary of Subatomic Particles
Atoms are composed of three smaller, subatomic
particles: electrons, protons and neutrons.
Chemical properties of atoms can be described based on
the electrons, protons and neutrons.
Though other subatomic particles are now known, the
theories of atomic structure are based only on these 3
subatomic particles.
Nuclear Model of the Atom
In 1911, Ernest
Rutherford
established the
nuclear model of
the atom.
Most of the particles passed through the gold foil, but some were
deflected and some even bounced back!
This suggested the gold atoms must have a densely, positively
charged nucleus to affect the path of an α particle (a positively
charged He atom).
Nuclear Model of the Atom
Because most of the particles
were not deflected, this
suggested most of the atom
is empty space.
Protons and neutrons are located in the nucleus.
Electrons are dispersed throughout the remainder
of the atom (mainly open space).
Neutral atoms contain the same number of protons and
neutrons to maintain charge balance.
Atomic Number
Atomic Number: Number of protons in the nucleus of an atom.
The atomic number determines the identity of the atom.
Atomic numbers for every element are above the
element’s symbol in the periodic table.
27
Co
Atomic Number
Isotopes of the Elements
After discovery of the nuclear model of the atom,
the mass of almost all atoms was found to be larger than
expected, based on the number of protons and electrons.
This led to the discovery of neutrons.
Though all atoms of the same element have the same
number of protons, atoms of the same element may have
different numbers of neutrons.
Isotopes: atoms of an element with the same atomic
number but different numbers of neutrons.
Isotopes of the Elements
Example: Isotopes of Hydrogen
Protium
Deuterium
1 proton
0 neutrons
1 proton
1 neutron
Tritium
1 proton
2 neutrons
Standard Isotopic Notation
Mass Number
Atomic Number
A
E
Z
Element Symbol
Mass number: Total number of protons and neutrons for an element.
Isotopes of the Elements
Practice:
How many protons, neutrons, and electrons are found
in each of the following isotopes? 64
Cu
29
Atomic Number: 29 protons (therefore 29 electrons)
# Neutrons = Mass Number – Atomic Number
64 – 29 = 35 neutrons
Let’s Practice!
Which isotope corresponds to an element with
24 protons and 28 neutrons?
a. 28 Cr
52
Solution:
b. 52 Cr
24
# protons = Atomic Number = 24 Element: Cr
c. 52 Ni
28
Mass Number = protons + neutrons
128 Te
d. 52
e. 24 Cr
52
!
= 24 + 28 = 52
Atomic Mass
Because the mass of a single atom is so small, it is
inconvenient to use this as a mass unit.
Instead, relative atomic mass units (amu) are used.
Using carbon-12, as a standard, 1 atomic mass unit
is equal to 1/12th the mass of a carbon-12 atom.
1 amu = 1.6606 x 10-24 g
All periodic tables use atomic masses
based on the carbon-12 isotope.
Atomic Mass and Isotope Distribution
Since most elements are a mixture of isotopes, the atomic mass
for an element is the weighted average of all naturally occurring
isotopes of the element.
Example:
The atomic mass of Cu is 63.546 amu. Cu exists as 2
major isotopes, Cu-63 and Cu-65.
Cu-63 is more abundant IS more abundant than Cu-65
The average atomic mass is equal to the sum of the atomic
mass of each isotope multiplied by its % abundance.
Atomic Mass and Isotope Distribution
Average atomic mass of Cu:
(62.9298) x (0.6909) + (64.9278) x (0.3091) = 63.55 amu
Atomic Mass % Abundance
Atomic Mass % Abundance
Atomic Mass Practice
Silver exists as two isotopes with atomic masses of
106.9041 and 108.9047 amu. Determine the average atomic
mass for silver if the % abundance for each isotope is 51.82
and 48.18%, respectively.
Average atomic mass of Ag:
(106.9041) x (.5182) + (108.9047) x (0.4818) = 107.8680 amu
Atomic Mass
% Abundance
Atomic Mass
% Abundance
Let’s Practice!
Chlorine exists as two isotopes, Cl-37 (36.96590 amu)
and Cl-35. If the percent abundance of each isotope is
24.47 % and 75.53 %, what is the atomic mass of
Cl-35 if the average atomic mass is 35.46 amu?
Solve for a:
(36.96590) x (.2447) + (a) x (0.7553) = 35.46 amu
a.
b.
c.
d.
e.
36.95690
34.96885
36.57823
33.56438
35.64544
Solution:
9.046 + (a) x (0.7553) = 35.46 amu
(a) x (0.7553) = 26.41 amu
a = 34.97 amu