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U N I T
L E C T U R E
1
N O T E S
LECTURE 11
Atomic Models
(3.3, 3.5 & 11.1)
GENERAL CHEMISTRY
Fall 2009
Updated: 5/3/2017
3.3 Dalton’s Atomic Theory
John Dalton (1766 – 1844) explained observations such as the law of constant
composition (a compound always has the same composition) using his theory. The
predictive value of the theory led to its eventual acceptance.
A. Defining the Atom
1. Atomic Theory
a. All matter is made up of very tiny particles called atoms
b. Atoms of the same element are chemically alike
c. Individual atoms of an element may not all have the same mass.
However, the atoms of an element have a definite average mass that is
characteristic of the element
d. Atoms of different elements have different average masses
e. Atoms are not subdivided, created, or destroyed in chemical reactions
3.5 The Structure of the Atom
1e. Students know the nucleus of the atom is much smaller than the atom yet contains most of its mass.
The volume of the hydrogen nucleus is about one trillion times less than the volume of the hydrogen atom, yet the nucleus contains almost all the mass in
the form of one proton. The diameter of an atom of any one of the elements is about 10,000 to 100,000 times greater than the diameter of the nucleus. The
mass of the atom is densely packed in the nucleus.
The electrons occupy a large region of space centered around a tiny nucleus, and so it is this region that defines the volume of the atom. If the nucleus
(proton) of a hydrogen atom were as large as the width of a human thumb, the electron would be on the average about one kilometer away in a great
expanse of empty space. The electron is almost 2,000 times lighter than the proton; therefore, the large region of space occupied by the electron contains less
than 0.1 percent of the mass of the atom.
2. Sizes of Atoms
a. Atomic radius
i. 40 to 270 picometers (pm)
1. 1 pm = 10-12m
ii. Most of the atomic radius is due to the electron cloud
b. Nuclear radius
i. 0.001 pm
ii. density is 2x108 metric tons/cm3
1. 1 metric ton = 1000kg
Updated: 5/3/2017
B.
Models of the Atom
Scientist
Year
Democritus
Dalton
~ 400
B.C.E.
1808
Model
Thomson
1897
Plank
1900
Rutherford
1911
Bohr
1913
Einstein
Schrödinger
1905
1926
Energy emitted in discrete
quantities
Nuclear Atom; also called the
planetary model
Bohr Model, electrons travel in
discrete orbits
Wave mechanical model
Wave mechanical model
Heisenberg
1929
Wave mechanical model
Solid sphere, tiny, indivisible,
indestructible particles
Plum pudding
Experiment
Focus
None
Suggested Atom
Weather data
Cathode Ray Tube;
also invented mass
spectrometer
Radiation from solids
Electrons
Gold foil
Nucleus
Spectrum of Hydrogen
Excited and Ground state
Photoelectric Effect
Schrödinger cat;
thought experiment
Photons
Schrödinger equation
Quanta
Heisenberg uncertainty
principle
Quarks and leptons
(matter)
Murray Gell1970s
Standard Model
Mann; George
Zweig
* There were many other models developed during this time period but we’ll only focus on these particular ones.
C.
Contributed to the Models of the Atom
Maxwell
1873
Planck
1900
Chadwick
1932
Visible light consists of
electromagnetic waves
Energy emitted in discrete
quantities
Radiation from solids
Identified subatomic
particle
Provides description of
light
Quanta; Plank’s constant
Neutron
Updated: 5/3/2017
11.1 Rutherford’s Atom
1a. Students know how to relate the position of an element in the periodic table to its atomic number and atomic mass.
An atom consists of a nucleus made of protons and neutrons that is orbited by electrons. The number of protons, not electrons or neutrons, determines the unique
properties of an element. This number of protons is called the atomic number. Elements are arranged on the periodic table in order of increasing atomic number.
Historically, elements were ordered by atomic mass, but now scientists know that this order would lead to misplaced elements (e.g., tellurium and iodine) because
differences in the number of neutrons for isotopes of the same element affect the atomic mass but do not change the identity of the element.
D. Structure of the Nuclear Atom
1. The Electron
a. Discovery
i. Joseph John Thomson (1897)
1. Cathode ray tube produces a ray with a constant charge to
mass ratio
2. All cathode rays are composed of identical negatively
charged particles (electrons)
ii. Plum-pudding model
b. Inferences from the properties of electrons
i. Atoms are neutral, so there must be positive charges to balance
the negatives
ii. Electrons have little mass, so atoms must contain other particles
that account for most of the mass
2. The Nucleus
a. The Rutherford Experiment (1911)
b. Alpha particles (helium nuclei) fired at a thin sheet of gold
i. Assumed that the positively charged particles were bounced back if
they approached a positively charged atomic nucleus head-on (Like
charges repel one another)
Updated: 5/3/2017
Results from gold foil experiment
1. Very few particles were greatly deflected back from the gold sheet
a. nucleus is very small, dense and positively charged
b. most of the atom is empty space
2. Structure of the Nucleus
a. Protons
i.
Positive charge, mass of 1.673x10-27kg
ii.
The number of protons in the nucleus determines the atom's
identity and is called the atomic number (Z)
b. Neutrons
i.
James Chadwick (1932)
ii.
No charge, mass of 1.675x10-27kg
c. Nuclear Forces
i.
Short range attractive forces:
a. neutron-to-neutron, proton-to-proton,
neutron
proton-to-
Unanswered Questions
What are the electrons doing? How are the electrons arranged? How do
electrons move? Why aren’t electrons (negatively charged) attracted to the
positive nucleus?
Updated: 5/3/2017