Download Chapter 4: Atomic Structure

CHAPTER 4:
ATOMIC STRUCTURE
Intro Video!
(nothing about
Bohr, I
promise)
I. HISTORY OF ATOMIC THEORY
A. Highlights:
1. Democritus: suggested matter was made of tiny indivisible particles
2. Aristotle: rejected idea of the atom & suggested that all matter was
composed of one continuous substance called hyle (similar to clay);
accepted as true until 17 th century
3. Antoine Lavoisier (1770s): explained burning & law of conservation
of matter (mass remains constant)
4. Joseph Proust (1799): experimentally showed that matter does not
combine randomly; law of definite proportions
5. John Dalton (1803): experimented with many compounds,
carefully massing the products formed and comparing them
with the original reactants; led to law of multiple proportions
6. Dalton revived the particle theory of matter and made the
following statements based on experimental evidence:
a. All elements are composed
of atoms which are tiny
indivisible particles.
b. All atoms of the same
element are exactly alike.
c. Atoms of different
elements are different.
d. Compounds are formed by
joining of atoms from two
or more elements and
always in whole number
ratios.
7. Corrections to Dalton’s
theory:
a. Atoms can, to a small degree, be
altered by chemical change.
b. Atoms are not indivisible and can
be changed by nuclear reactions.
c. Atoms of the same element can
be different, isotopes.
 Copy into notes: Figures 4-2, 4-3, and 4-5 (4-2 & 4-3: do not
draw pictures, just copy text; 4 -5: draw pictures, copy text &
copy caption)
 Classwork: p. 91 #1-5 (will check tomorrow)
II. DISCOVERING SUBATOMIC PARTICLES
A. William Crookes (English physicist, 1870’s): worked with
cathode rays in discharge tubes & found that they could be
deflected by a magnet; concluded that the rays consisted of tiny
electrically charged particles
1. Due to Crookes’ work &
further research, the
scientific community
determined that these
particles contained a
negative charge, but there
was no concrete evidence
of this
B. J.J Thomson (English scientist, 1897): proved Crookes was right.
1. Wanted to determine the ratio of this particle’s negative charge
to its mass
2. Worked with cathode rays to show that both electric and
magnetic fields caused deflections & measured these deflections
3. Found a charge-to-mass ratio of the charged particle & compared
it to other known charge-to-mass ratios; discovered that the mass
of this particle was less than the mass of the smallest atom
(hydrogen)
4. Determined that this “charged particle” must be smaller than an
atom & therefore part of the atom; called these electrons
5. Calculated an e/m (charge/mass) ratio from data
e/m = 1.76 x 10 8 coulomb/g.
C. Robert Millikan (American physicist, 1909-11)
1. Performed the famous “oil drop experiment” &
determined the exact charge of the electron
2. All the charges were multiples of a smaller
charge; he calculated the smallest charge to be
1.60 x 10 -19 coulomb
3. He used Thomson’s e/m ratio & solved for the
mass of an electron; 9.11 x 10 -28 g
4. The discovery of the electron & its mass led to
J.J. Thomson’s plum pudding model – small
negative charges distributed throughout a
uniform positive charge
D. Lord Ernest Rutherford (English scientist, 1912) & squad (Niels
Bohr, Hans Geiger, Ernest Marsden)
1. Performed the famous “gold foil experiment” – shot alpha
particles (+) from a radioactive source at a thin sheet of gold foil.
Most passed right through, some were deflected, and a few
bounced back!
2. Made the following conclusions:
a. The foil was mostly empty space
b. The deflections were caused by a positive charge
c. The positive charge was in a very small area in the center of the atom
(nucleus)
d. The atom then consisted of a small dense positive core (nucleus)
orbited by electrons
e. The opposite charge of the positive core and the negative electrons
canceled each other out, resulting in a neutral atom.
3. He further surmised that the nucleus of hydrogen, the least
massive atom, is a single particle with a positive charge, the
proton.
4. Nuclei of larger atoms then contained a larger number of
protons.
5. Problem: he found that if he set the mass of the proton at 1.00,
then the protons that account for the positive charge do not
account for the total mass – only part of it.
6. Therefore, the atom had more mass than the amount contributed
by the protons.
7. Suggested tightly bound proton-electron pairs which he called
neutrons to account for the missing mass – no real evidence for
this.
E. James Chadwick - English scientist (1932).
1. Bombarded the element beryllium with alpha particles.
2. A new type of particle was given off which had no charge and
about the same mass as the proton.
3. Concluded that these particles were Rutherford’s proton -electron
pairs and kept the name neutron for these particles.
Copy Figures 4-9, 4-12, and 4-13 (include all text), & Table
4-1
Work in pairs or groups of 3. Create a poster addressing the following
models of the atom:
• John Dalton's solid sphere model
• J.J. Thomson's plum pudding model
• Ernest Rutherford's nuclear model
Poster must include:
• Who — Person(s) responsible
• What — A description and an illustration of the model; make sure to
address how protons, neutrons, and electrons are included in the
model
• Where — Country, state, university, etc.
• When — Date(s)
• Why — The prior knowledge at the time
• How — A summary of the technology or evidence used to develop the
model
• References — Page numbers from our textbook and at least two
other helpful resources
III. DIFFERENCES IN ATOMS
A. Atom: smallest amount of a chemical element that can still
take part in a chemical reaction. 3 parts:
1. Proton (p + ): positive particle found in the nucleus
Mass: 1.0073 amu
Charge: + 1.60 x 10 -19 coulomb
2. Neutron (n o ): neutral particle found in the nucleus
Mass: 1.0087 amu (about the same as p + )
Charge: 0
3. Electron (e - ): negative particle found orbiting the nucleus
Mass = 0.000549 amu (about 2000x less massive than p + & n o
Charge = - 1.60 x 10 -19 coulomb
sketch & label this
B. Atomic number: number of protons in an atom
1. Elements are composed of atoms which all have the same number of
protons
2. atomic number = number of protons = number of electrons
(in a neutral atom)
3. Henry Moseley (English physicist, 1913) found that the wavelength of
x-rays from a source depended on the number of protons present;
performed experiments to determine the atomic numbers of many
elements
C. Isotopes: atoms of the same element (therefore, same # of
protons) with dif fering numbers of neutrons
1. Isotopes of an element usually have the same chemical behavior
because chemical behavior is determined by electrons
2. Example: Isotopes of hydrogen (copy this diagram exactly, words
included)
D. Mass number: number of protons and neutrons in the
nucleus of a particular atom
1. mass number = #protons + #neutrons
2. Example: Find the number of neutrons in potassium-40
 Since it is called “potassium-40,” that means that the mass
number is 40
 Find # of protons (from periodic table -- # of protons is the same as
the atomic number) – potassium’s atomic number = 19
 mass number = #protons + #neutrons
40 = 19 + x
21 = x
 Potassium-40 contains 21 neutrons
 Copy Figure 4-14 (p. 98)
 Bookwork: p. 97 #6-10, p. 99 #11-13, p. 101 #14