Chapter 4 –Atomic Structure
Honors Chemistry
(Student edition)
Chapter 4 problem set:
Useful Diagrams:
4.1
36, 40, 50, 51, 59, 65, 80, 81
4.2, 4.4, 4.5, 4.7, 4.10 and tables 4.1, 4.3
Early Models of the Atom
Atoms: are the fundamental building blocks of matter.
Atoms: the smallest piece of an element that still retains the properties of that element.
Some Historical Background:
Democritus’s Atomic Theory:
In approximately 460 BC, Democritus (Greek) coins the term “atom” (means
). Before that matter was thought to be one continuous piece
- called the
theory of matter. Democritus creates the
theory of matter.
Solid
Sphere
In the 18th century, experimental evidence appears to support the idea of atoms.
Law of Conservation of Mass: Antoine Lavoisier (French) - 1770’s (see Ch2)
Law of Definite Proportions: In 1799, Joseph Proust (French) states:
“The proportions of masses of chemicals in reactions are always the same.”
Examples:
8g O
+
1 g H yields
9 g H2 O
16g O
+
2 g H yields 18 g H2O
Law of multiple proportions: In 1803, John Dalton - English school teacher states:
“The mass of one element combines with masses of other elements in simple, whole # ratios.”
Examples:
2H
+
1O
yields
H2O
2H
+
2O
yields
H2O2
1
Dalton’s Atomic Theory : Dalton put together the laws of conservation of mass,
definite proportion, and multiple proportion to create his own atomic theory.
*1. All matter is made up of atoms which are indivisible (can’t be destroyed)
*2. Atoms of the same element are identical
3. Atoms combine in simple whole number ratios to form compounds.
4. Chemical reactions are the result of the separating or combining of atoms


#1 is wrong as an exception is
#2 is wrong because of the existence of
(
)
Sizing up the Atom: Individual atoms are observable with instruments such as a scanning
tunneling microscope.
4.2 Structure of the Nuclear Atom
Subatomic Particles: Atoms are made up of
Electron:
charge,
.
amu (
)
In the 1870’s, English physicist William Crookes studied the behavior of gases in
vacuum tubes.
Crookes tubes - forerunner of picture tubes in TVs
Crookes’ theory was that some kind of radiation or particles were traveling from
the cathode across the tube. He named them
.
20 years later, J.J. Thomson (English) repeated those experiments and devised
new ones. In 1897, JJ Thomas discovered the
.
2
Thomson used a variety of materials, so he figured cathode ray particles must be
fundamental to all atoms. J.J. Thomson gets credit for discovering the electron.
Plum
Pudding
Model
+
+
Thomson and Milliken (oil drop experiment) worked together (their data, not
them) to discover the
and
of the electron
Electron charge:
Electron mass:
coulomb this is the smallest charge ever detected
grams this weight is pretty insignificant
Proton: In 1886, Eugen Goldstein found evidence for the proton. It has
charge and a mass of
amu.
Neutron: In 1932, James Chadwick confirmed the existence of the
It is
, slightly more than
.
amu.
NIB - Quarks:
They are make up protons and neutrons. There are 6 types:
ups +2/3 charge
downs - 1/3 charge
so....
2 ups, 1 down =
2 downs, 1 up =
Particle accelerators: miles long with magnets that propel particles along
the chamber. The particles smash into each other at high speeds. This
results in nuclear disintegration.
Other simpler particles:
The Atomic Nucleus:
The Rutherford Gold Foil Experiment: In 1911, Rutherford (New Zealand) …
3
1 – the positive portion of the atom is in the middle
2 – most of the atom is empty
3 – most of the mass is in the middle
4 – electrons orbit the nucleus
Conclusions:
Analogy: if an atom is the size of the Eagle’s stadium, then the nucleus is the size
of a tennis ball floating in the middle of the stadium.
+
-
Shortcomings of the Rutherford Model: According to gravity, electrons should
move towards the nucleus eventually - they don’t. So… more work needs to be
done to understand the structure of the atom.
Technology and Society: In 1931, Ernst Ruska and Max Knoll built the first
electron microscope. It uses an electron beam and “lenses” that consist of magnetic or
electric fields. Objects can be magnified over 100,000 times and projected on to a
monitor. Biochemistry, Microelectronics, and Biology all use this technology.
4.3 Distinguishing Among Atoms
Atomic Number:
: particles that make up the nucleus.
Proton and Neutrons make up most of the
Protons: 1 amu,
of atoms.
grams, positive charge, determines
Atomic # (Z): Always a
periodic table. # of
if the element is
Neutron: neutral, determines the
of the atom.
number,
number on the
in the nucleus, also indicates the # of electrons
charged
of the atom, mass is slightly more than 1 amu
Electrons: not a nucleon, negative charge, orbit nucleus, determine
of an atom
Mass Number:
4
Originally it was thought that all atoms of the same element had the same mass (Dalton)
Hydrogen was observed to have the lowest mass (assigned a weight of “1”).
Original periodic table listed elements in order of their atomic
This is not true today (see
).
.
Mass number: Represented by the letter “A” . It is the sum of the protons and
neutrons in a nucleus. This number is rounded from atomic mass due to the fact
that there are isotopes.
# neutrons = A - Z
Example - # of neutrons in Li =
Isotopes: atoms of the same element with different masses because they have different
numbers of neutrons.
Hydrogen
Isotopes
Protons
Neutrons
Electrons
Mass
% Abundance
Protium
Deuterium
Tritium (artificial
and radioactive)
Some isotopes occur naturally. Most isotopes are produced artificially.
Ions: Atoms that have lost or gained electrons
+ Charge indicates the
- Charge indicates the
Isotope
40 +1
K
Protons
Electrons
12
91
-1
36
10
of an electron (
of an electron (
Neutrons
Atomic #
)
)
Mass #
40
12
53
19
C
S-2
Na+1
Br
14
5
Atomic Mass:
Mass of Cl thought to be 35.5 times that of hydrogen.
Today we know this isn’t true. It’s the weighted average of 2 isotopes:
75%
35
Cl
and
25%
37
Cl
Average atomic mass = [(%)(mass of 1st isotope)] + [(%)(mass of 2nd isotope)]......
Sample problem: find the average atomic mass of B:
11
B = 80.20%
10
B = 19.80%
Sample problem: find the %’s of 2 isotopes of Carbon given the following information:
average atomic mass = 12.0111 isotope 1 = 12 C , isotope 2 = 13 C
History lesson - originally H was the basis of all atomic masses and was given
the mass of 1.0. Later, chemists changed the standard to oxygen being 16.000 (which left H =
1.008). In 1961, chemists agreed that 12C is the standard upon which all other masses are based.
1/12 of the mass of 1 atom of 12C = 1 amu
Periodic Table Preview: see chapter 6
NIB Technology, and Society: The mass spectrometer
The mass spec (chemist lingo) is used to detect, analyze, and identify unknown
chemicals. Samples are vaporized, bombarded with electrons (in order to create + ions) and
placed in electrical and magnetic fields.
Due to differences in mass (# of neutrons) the paths of the molecules curve based on their
individual mass. Heavier particles curve less. This change in curvature causes the particles to
land on different places on a detector.
Draw picture of mass spectrometer below
The mass spectrometer was invented in 1912. By 1922 it was discovered that there were
300 naturally occurring isotopes that existed out of the 92 elements known at that time.
Now used for identifying components of mixtures, analyzing pollution, and dating works of art.
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