I. Atomic Concepts
1.1 The modern model of the atom has
evolved over a long period of time through
the work of many scientists. (3.la)
Atomic structure models
 The model for atomic structure of the
elements has passed through many stages of
development.
In the Bohr model, electrons were considered
to revolve around the nucleus in one of several
concentric circular orbits.
 The orbital model differs from the Bohr
model in that it does not represent electrons
as moving in planetary orbits around the
nucleus.


The term orbital refers to the average region
of most probable electron location. Electrons
occupy orbitals that may differ in size, shape,
or orientation in space.
 Rutherford’s gold foil experiments indicated
the atom to be mostly empty space with the

size of the nucleus very small compared to the
size of the atom.
http://www2.biglobe.ne.jp/~norimari/science/JavaApp/e-Scatter.html
1.2 Each atom has a nucleus, with an overall
positive charge, surrounded by one or more
negatively charged electrons. (3.lb)
 Atoms differ in the number of protons and
neutrons in the nucleus and in the
configuration of electrons surrounding the
nucleus
 The mass of the atom is concentrated almost
entirely in the nucleus.
 The nature of the forces holding nuclear
particles together is not adequately
understood and is the subject of much current
research.
1.3 Subatomic particles contained in the
nucleus include protons and neutrons. (3.lc)
The atomic number indicates the number protons
in the nucleus.
The mass number indicates the total number of
protons and neutrons.
1.4 The proton is positively charged, and the
neutron has no charge. The electron is
negatively charged. (3.ld)
1.5 Protons and electrons have equal but
opposite charges. The number of protons
equals the number of electrons in an atom.
(3.le)
1.6 The mass of each proton and each
neutron is approximately equal to one atomic
mass unit. An electron is much less massive
than a proton or a neutron. (3.lf)
1.7 In the wave-mechanical model (electron
cloud model), the electrons are in orbitals,
which are defined as the regions of the most
probable electron location (ground state).
(3.lh)

http://winter.group.shef.ac.uk/orbitron/AOs/1s/index.html
1.8 Each electron in an atom has its own
distinct amount of energy. (3.li)
 Electrons in orbits near the nucleus are at
lower energy levels than those in orbits
farther from the nucleus.
1.9 When an electron in an atom gains a
specific amount of energy, the electron is at a
higher energy state (excited state). (3.lj)
1.10 When an electron returns from a higher
energy state to a lower energy state, a specific
amount of energy is emitted. This emitted
energy can be used to identify an element.
(3.1k)
 Electrons can absorb energy only in
discrete amounts called quanta.
 When electrons return to a lower energy
level, energy is emitted in quanta.
 The excited state is unstable, and the electrons
fall back to lower energy levels.
 When electrons in an atom in the excited state
return to lower energy levels, the energy is
emitted as radiant energy of specific
frequency, producing characteristic spectral
lines which can be used to identify the
element.
 The study of spectral lines has provided much
of the evidence regarding energy levels within
the atom.

http://www.colorado.edu/physics/2000/quantumzone/index.html
1.11 The outermost electrons in an atom are
called the valence electrons. In general, the
number of valence electrons affects the
chemical properties of an element. (3.11)
 The term “kernel” is sometimes used to refer
to the atom exclusive of the valence electrons.
The kernel consists of the nucleus and all
electrons except the valence electrons.
 The valence electrons may be represented by
electron dot symbols in which the kernel of
the atom is represented by the letter symbol
for the element and the valence electrons are
represented by dots.
1.12 Atoms of an element that contain the
same number of protons but a different
number of neutrons are called isotopes of that
element. (3.lm)
 The atomic number identifies the element.
The difference in the number of neutrons
affects the mass of the atom.
 For a given element the number of protons in
the nucleus remains constant but the number
of neutrons may vary.
1.13 The average atomic mass of an element is
the weighted average of the masses of its
naturally occurring isotopes. (3.ln)
 http://www.colorado.edu/physics/2000/isotopes/index.html
 The atomic mass of an element given in the
Reference Tables for Chemistry is the
weighted average mass of the naturally
occurring isotopes of that element. This
average is weighted according to the
proportions in which the isotopes occur.
 Most elements occur naturally as mixtures of
isotopes. This accounts for fractional atomic
masses found in reference tables. In general,
the mass number of the most abundant
isotope of an element can be determined by
rounding off the atomic mass of the element to
the nearest whole number.
 The atomic mass (weight) of a single atom
(isotope) such as neon-20 is 19.992 amu while
the atomic mass (weight) of the element neon
(which is the weighted average of all its
natural isotopes) is 20.183 amu.