Download Honors Chemistry Chapter 4 Student Notes

Honors Chemistry
Chapter 4 –Atomic Structure
(Student’s 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
Defining the Atom
Atoms: are the
.
Atoms: the smallest piece of an element that still retains the
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
piece - called the
theory of matter. Democritus creates the
theory of matter.
Solid
Sphere
In the 18th century ,
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
of masses of chemicals in reactions are always the
Examples:
8g O
+
1 g H yields ____ g H2O
16g O
+
2 g H yields ____ 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,
Examples:
2H
+
1O
yields
2H
+
2O
yields
ratios.”
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.
*2.
3.
4.

#1 is

#2 is
Sizing up the Atom: Individual atoms are observable with instruments such as a
microscope.
4.2 Structure of the Nuclear Atom
Subatomic Particles: Atoms are made up of
Electron:
charge, 1/1837 amu (.0005)
In the 1870’s, English physicist William Crookes studied the behavior of gases in
tubes.
Crookes tubes - forerunner of picture tubes in
e- e- e- e-
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
Thomson used a variety of materials, so he figured cathode ray particles must be
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:
this is the smallest charge ever detected
this weight is pretty insignificant
Proton: In 1886, Eugen Goldstein found evidence for the
charge and a mass of
amu.
. It has
Neutron: In 1932, James Chadwick confirmed the existence of the
slightly more than
amu.
NIB - Quarks:
They are made up
ups +2/3 charge
and
. It is
. There are 6 types:
downs - 1/3 charge
so....
2 ups, 1 down = proton
charge
2 downs, 1 up = neutron
charge
Particle accelerators: miles long with
propel particles along
the chamber. The particles
into each other at high speeds. This
results in nuclear
.
Other simpler particles:
.
The Atomic Nucleus:
The Rutherford Gold Foil Experiment:
In 1911, Rutherford (New Zealand) …
- The Experiment:
towards a thin sheet of
the
screen.
Concluded:
particles from
(in the lead box) were released
foil. Most of the particles went through and were seen on
alpha particles bounced back.
1–
2–
3–
4–
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:
Nucleons: particles that make up the
.
Proton and Neutrons make up most of the
Protons: 1 amu,
Atomic # (Z): Always a
# of
element is
of atoms.
, positive charge, determines
of the atom.
number,
number on the periodic table.
in the nucleus, also indicates the # of electrons if the
charged
Neutron: neutral, determines the
of the atom, mass is slightly more than 1 amu
Electrons: not a nucleon, negative charge,
nucleus, determine
of an atom
Mass Number:
Originally it was thought that all atoms of the same element had the
Hydrogen was observed to have the
mass (Dalton)
mass (assigned a weight of “1”).
Original periodic table listed elements in order of their atomic
not true today (see Iodine).
. This is
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 =
Example - # of neutrons in Li =
Isotopes: atoms of the same element with different
numbers of
.
because they have different
Hydrogen
Isotopes
Protons
Neutrons
Electrons
Mass
% Abundance
Protium
Deuterium
Tritium (artificial
and radioactive)
1
1 amu
99.85%
1
2 amu
.15%
1
3 amu
0%
Some isotopes occur naturally. Most isotopes are produced artificially.
Counting protons, electrons, and neutrons:
Mass # =
Neutrons =
Protons =
Atomic # =
Electrons =
Isotope
40 +1
K
+ Charge indicates the
of an electron
- Charge indicates the
of an electron
Protons
Electrons
12
Neutrons
Atomic #
Mass #
40
12
53
36
10
19
C
S-2
Na+1
Br
14
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
Analogy for weighted averages:
If your homework grade is 80.0 and your test grade is 95.0, then what is
your average? Note: homework is worth 50.0% of your grade.
If your homework grade is 80.0 and your test grade is 95.0, then what is
your average? Note: homework is worth 20.0% of your grade.
Average atomic mass = [(%)(mass of 1st isotope)] + [(%)(mass of 2nd isotope)]......
Sample problem: find the average atomic mass of B
B11 = 80.20%
B10 = 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
Another example: Antimony consists of two naturally occurring isotopes:
Antimony-121 (57.2500% abundance and an exact mass of 120.9038) and some
other isotope. Calculate the percent and mass of the other isotope.
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
[positively charged particles due to a loss of 1 electron]), 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.
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.
Used for identifying components of mixtures, analyzing pollution, and dating
works of art.