The very basics:
Notes on Atomic structure
Atom: an electrically neutral combination of protons, neutrons and
electrons. The protons and neutrons are located in the center of the
atom in what is called a nucleus. The electrons are located in a region
around the nucleus in various energy levels or shells.
Each type of atom can be identified by an atomic number, which is
the number of protons that it contains.
Ion: an electrically charged atom formed by either adding or
removing electrons.
Cation: a positive ion
Anion: a negative ion
Isotope: varieties of atoms that have the same number of protons
but have different numbers of neutrons.
Or
Atom: The smallest building block of matter; a collection of protons,
neutrons and electrons.
An Element is a particular type of atom
Ion: a charged atom. An atom that has gained or lost electrons from its
normal state. Cations are positive. Anions are negative.
Isotopes: all atoms are some kind of isotope. All isotopes act the
same chemically but have slightly different masses (due to the number of
neutrons)
Teacher/student notes:
Construct a conceptual map with these terms: how are they related
Atomic structure
Electron
Proton
Atomic number
Mass number
neutron
Nucleus
isotopes
Remember that “all atoms are some particular isotope”
What is reported on the periodic table of elements is the weighted average
of all isotopes
- The atomic number (on periodic table) is the number of protons
- On the worksheet the Mass is really the mass number of a particular
isotope
- The sum of the protons and neutrons will be this mass number
- The # of electron is the sum of the electrons in each shell (and for a
neutral atom will equal the number of protons; the number of
protons and electrons will not be equal in an ION.
- When drawing a (2-dimensional) representation of an atom space
out the electrons into the 4 corners of the orbit doubling up when
there are more than 4 electron
- Remember these are only models the nucleus of an atom is 10,000
times smaller than the whole (really really, really, really dense stuff)
Isotopes are very common in the physical world,
Some are very radioactive and spontaneously undergo change into new
elements (radioactive decay, this is not “chemistry”) others are much
more stable and contribute to the “average atomic mass” of an element.
The value of 1 A.M.U. (or one atomic mass unit is defined as 1/12 the
mass of a carbon -12 atom)
There are several methods of describing a particular isotope learn them
now but for most of the year we will just use the simple chemical symbol
for an “average” element
There are 7 special element know as “the diatomic seven that you should
MEMORIZE (see list)
In chemistry we will calculate average atomic mass a little differently that
you may be used to; SEE Below
The seven diatomic elements
H2, N2, O2, F2, Cl2, Br2, I2
When in their natural state (gases since most have very low boiling points) these elements “pair up” into
molecules (non-metal to non-metal). They seem to do this to have a complete valence shell. You must
have these memorized in order to properly write balanced chemical equations. Here are some tricks I
have learned.
An oldie but a goodie
I Bring Clay For Our New House
( 4 paving 8 sidewalks; P4 S8 )
One teacher said all the elements got together to honor to honor a great
chemistry Ms.
(or Mr.) HOFBrINCl
(spell it just like it sounds)
One particularly energetic cheerleader noticed that it was
hockey stick (3 across and 4 down N,
H (hockey puck) and a
O, F, Cl, Br, I).
However you wish to memorize them JUST DO IT !!!
In chemistry Average Atomic Mass Analogy
Average cost of a can of pop
In chemistry you will calculate averages slightly differently than the simple add them all up and divide
method you may be familiar with. In Chemistry since we usually deal with large numbers of atoms we
will calculate the percentage first and then find what is called the weighted proportion and add these
weighted proportions up.
For example look at the data table for pop purchases
Cost for pop purchase
Where purchased
$0.50
machine (near home)
$0.65
Machine (at Mall)
$0.50
machine (near home)
$1.15
Embers restaurant
$0.50
machine (near home)
$0.65
Machine (at Mall)
$0.65
Machine (at Mall)
$0.50
machine (near home)
$1.25
Movies
$0.50
machine (near home)
A total of 10 cans were purchased: and using your grade school skills it is not hard to determine the
“average” price.
However in Chemistry there are usually so many “individual” atoms that we don’t even consider them
as individuals. Instead we use the percentage.
Analyzing the data we see that:
50% of the purchases cost $0.50
30% of the purchases cost $0.65
10% cost $1.15
10% cost $1.25
The “average can” is made up of the sum of these weighted portions.
.50 times 50 cents equals 25.0 cents of the “average”
.30 times 65 cents equals 19.5 cents of the “average”
.10 times 115 cents equals 11.5 cents of the “average”
.10 times 125 cents equals 12.5 cents of the “average”
Summing these weighted portions gives. 68.5 cents (the same answer you obtained using the
elementary school method)
Sample problems:
Example #1) Rubidium has two common isotopes, 85Rb and 87Rb. If the abundance of 85Rb is
72.2% and the abundance of 87Rb is 27.8%, what is the average atomic mass of rubidium?
Not all problems involve chemistry:
Example #2) What would the average income be for a country that has 20% of the population
earning $20,000 per year ($10.00/hour) 75 % of the population earning $50,000 and 5%
earning 2 million per year.
Answers:
Ex 1)
(.722 * 85) + (.278 * 87) = 85.556 or 85.56 rounded to the hundredths of an amu
Ex #2)
(.20 * 20,000) + ( .75 * 50,000) + ( .05 * 2,000,000) = 141,500
Much more than the “average” person)
John Dalton is given much of the credit with reviving the idea of the
atom causing it to become the well known theory it is today.
The postulates of Dalton's atomic theory were used to explain earlier
observations of the behavior of matter. His postulates were...
(a) All matter is composed of small, indivisible particles called
atoms.
(b) All atoms of a given element are identical in mass and
properties.
(c) Compounds are formed by a combination of two or more
atoms in definite arrangements in the ratio of small whole
numbers.
(d) Atoms are not created, destroyed or converted into other kinds
of atoms during chemical reactions. They are simply rearranged
into new compounds.
These postulates were accepted for most of the 1800's until new
experiments indicated atoms were made up of subatomic particles.
These new experiments provided evidence that postulates "a" and "b"
were not correct.
A Brief History of the Atom:
copied from web site: http://www.cerritos.edu/ladkins/a106/A%20Brief%20History%20of%20the%20Atom.htm
Antiquity
 (400 B.C.) Democritus of Adbera (northern Greece) asserted that all material things are
composed of extremely small irreducible particles called atoms. “Nothing exists except atoms
and empty space. Everything else is opinion”. The atomic theory was roundly rejected by
Aristotle, and, thus, by almost everybody else for the next two millennia.
17th through 19th Centuries
 (1627-1691) Robert Boyle (England) extended mathematics to chemistry and revived atomic
theory.
 (1777) Antoine Lavoisier (France) demonstrated the conservation of matter (matter can be
neither created nor destroyed) in a chemical reaction and defined the difference between an
element and a compound.
 (1780) Charles Coulomb (France) described the force between two electric charges with a
mathematical formula which looked very much like Newton’s law of gravity:
qq
F  1 22
r
where F is the force q1 and q2 are two charges and r is the distance between them.
The electrical force is the chief force involved in atomic reactions. This force is attractive when
charges q1 and q2 have opposite signs and repulsive when the charges have the same sign.
 (1803) John Dalton (England) formulated the modern version of the atomic theory. In his model
all atoms in a given chemical element are exactly alike, while the atoms of different elements
differ by atomic weight

(1831-1879) James Clark Maxwell (England) showed that electricity and magnetism are two
aspects of the same phenomena, and predicted that accelerating charges radiate waves traveling
at the speed of light. These waves are known generically as electromagnetic waves of which
visible light is one example.
 (1898) J.J. Thompson (England) discovered the electron, the component of the atom with
negative charge. His model of the atom had the negatively charged electron evenly distributed
throughout a sphere of positively charged material. This is known as the “plum pudding” model
of the atom.
20th Century
 (1900) Max Planck (Germany) introduced the quantum theory to explain the shape of the
temperature versus color curve of a glowing solid. Briefly, he found that light cannot be
converted into heat (energy) by any arbitrary amount, but only as discrete packets which he
called quanta (known as photons today). For light of wavelength λ, the energy per quanta is
given by:
E


hc

where h is a constant which we now call planck’s constant.
(1905) Albert Einstein (Germany, USA) published papers on special relativity which included
the famous equation relating energy E to mass m:
E  mc 2 Here, c is the speed of light. Thus, the mass of any particle has an equivalent energy
and a photon, viewed by Planck as a packet of pure energy, has an equivalent mass.
(1909) Ernest Rutherford (England) demonstrated that the atom is mostly empty space with a
small positively charged nucleus containing most of the mass and low mass negatively charged
particles (Thompson’s electrons) orbiting this nucleus. Rutherford could experimentally identify

nuclear particles with positive charge that he called protons. Although he could explain the
charge of atomic nuclei with the right number of protons, the mass of the nucleus for large atoms
was always larger than the sum of its protons. Therefore he postulated the existence of a neutral
particle with a mass nearly the same as the proton which, when added to the protons in the
nucleus, would give the right mass. Rutherford called this hypothetical particle the neutron.
Later (1930) Rutherford’s colleague James Chadwick was able to detect the neutron
experimentally.
(1913) Neils Bohr (Denmark) developed the first successful model of the atom. Since we still
use Bohr’s model to explain many aspects of physical phenomena such as the appearance of
spectra, it is worthwhile to spend some time describing it.
Bohr’s model of the atom builds on Rutherford’s basic conception. In detail, the nucleus
contains a number of relatively high mass particles with positive charge called protons
along (sometimes, not always) with some neutral particles of about the same mass
called neutrons. A chemical element is defined and distinguished from all other
chemical elements by the number of protons in its nucleus. Orbiting the nucleus, much
like planets orbiting the sun, are the electrons. This is pretty much the picture that pops
into most people’s heads when they think of atoms. They get this picture because that
is how atoms are usually illustrated in everything from comic books to textbooks.
Now according to Maxwell, accelerating charges, such as electrons traveling in circular orbits,
should radiate electromagnetic waves and, hence, energy. This loss of energy should make the
electrons spiral down into the nucleus. To get around this problem, Bohr proposed that the
electrons were confined to specific orbits that were quantized. As long as the electrons remained
in one of the allowed orbits, no electromagnetic radiation will be released. Under ordinary
conditions the electrons of most atoms are in the lowest orbit available; under such conditions
the atom is said to be in the “ground state” and cannot radiate energy. To move an electron from
the ground state to one of the higher orbits requires the input of energy exactly equal to the
energy spacing between the two orbits. Once at the higher level, the electron can then fall back
to a lower orbit, radiating a photon with an energy corresponding to the orbital spacing.
To summarize: To radiate energy, and atom must first be excited (electrons raised above the
ground state). The excited atom then returns to the ground state by emitting energy in the form
of electromagnetic radiation.
Bohr Hydrogen Atom
Photon in
-
+
Hydrogen atom in ground state
Hydrogen atom excited
Photon out
Hydrogen atom emits photon
Bohr set about explaining the visible spectrum of the hydrogen atom, i.e., the Balmer series of
lines familiar to just about everyone who has ever taken an astronomy lab. Bohr was able to
show that this set of violet, blue and red lines originated from an electron falling from higher
orbits down to the orbit immediately above the ground state. More precisely, if we designate
each orbit with a number beginning with n=1 for the ground state, the Balmer series represents
the transition of the electron from orbit n>2 to orbit n=2. The higher the originating orbit, the
greater the energy of the photon emitted. For example, the red line, representing the longest
wavelength (and, thus, the lowest energy photon), is produced by the electron falling from orbit
n=3 to n=2. The next blue line comes from the electron in n=4 falling to n=2, and so forth.

Now before we can obtain the Balmer spectrum from a hydrogen atom, two criteria must be
satisfied: (1) there must be an electron available and (2) it must be in an orbit greater than n=2.
Criterion (1) will not be satisfied if the atom has been stripped of its electron. An atom in this
condition is referred to as ionized and it occurs at elevated temperatures. On the other hand,
criterion (2) will not be satisfied if the hydrogen atom is in the ground state, i.e., its electron has
not been excited into a higher orbit. From this we can see why neither hot, blue O type stars nor
cool, red M type stars exhibit strong hydrogen lines. Type O stars are so hot that most of the
hydrogen atoms in their atmospheres have been ionized, and are, hence, unavailable to form
spectra. On the other hand type M stars are too cool to excite very many of the hydrogen atoms
above the ground state. Thus, for opposite reasons, neither type O or type M stars have strong
hydrogen lines in their spectra.
(1924) Louis de Broglie (France) hypothesized that the electrons in Bohr’s model were confined
to discrete orbits because they had the properties of standing waves. He proposed that any
particle with a momentum p (p = mv) has an equivalent wavelength λ given by
h
h
 
p mv . Here m is the mass of the particle and v is its velocity. Calculations based on the
assumption that matter at the atomic level can be viewed as waves agreed so well with
experiments that it became a cornerstone of quantum mechanics. The theory is known today as
the Principle of Complementarity :


Waves and particles represent complementary aspects of the same phenomenon.
In short, wave phenomena such as light can also have the properties of particles, and particle
phenomena such as the constituents of atoms can also have the properties of waves.
(1925) Cecilia Payne (England, USA), using the new model of the atom, showed that the sun and
stars are composed almost entirely of hydrogen and helium, with only trace amounts of more
familiar, heavier, elements. She came to this conclusion by studying the spectral data which had
been accumulated at Harvard Observatory over the past quarter century. As was pointed out in
the discussion of the Bohr atom, both hot stars (O and B) and cool stars (K and M) do not exhibit
strong hydrogen lines for opposite reasons. In the first instance (hot stars), most of the hydrogen
atoms are ionized and, thus, have no electrons available to be raised to a higher orbit. In the
second instance (cool stars), the stars do not produce many photons of the energy required to
raise electrons above the ground state. However, hydrogen lines are not entirely absent; faint
hydrogen lines are seen in both groups of stars. This is so because at a given instant a few
hydrogen atoms in hot stars do have electrons, and even the coolest stars produce some energy in
the range necessary to excite the hydrogen atom. The percentage of non-ionized atoms in hot
stars and the percentage of excitation photons in the energy output of cool stars can be calculated
statistically. When Cecilia Payne made these calculations, she came to the conclusion that the
fact that any hydrogen lines at all are visible in these stars implies that the number of hydrogen
atoms present must be enormous – over 90% of the total number of atoms and over 70 % of the
stellar mass. Similar reasoning led to the conclusion that most of the remaining mass was made
up of helium. The same statistical approach was applied to the spectra of the middle range stars
(A, F, and G) and a similar composition was found for these stars as well. For all stars, only a
tiny fraction of the stellar mass, typically no more than 2%, was comprised of the heavier
elements such as oxygen and silicon, the most common elements on Earth. When the hydrogen
emission hydrogen spectra found in gaseous nebulae was factored in, the message seemed clear:
The universe is mostly hydrogen and helium, the two simplest elements.
(1925) Sir Arthur Eddington (England) produced the first model of stellar structure based on
nuclear physics. The energy source of the sun and stars had been a mystery for centuries. It is
easy to determine the total energy output of the sun. Basically, you measure how much heat is
transferred to a square meter of water (or other calibrated material) in a given amount of time,
then multiply this number times the surface area of a sphere centered on the sun with a radius of
1 A.U. This yields a total energy output of about 4 x 1026 joules/sec (watts). From geological
considerations this amount of energy has been produced without interruption for about 5 billion
years. No energy source known before the 20th century could have produced energy at this rate
for this length of time.
Scientist studying the nucleus in the early twentieth century noticed that the atomic weight of the
helium nucleus was slightly less than the sum of the protons and neutrons that comprised it. The
implication was that when protons and neutrons were added together to make helium, energy
was produced equal to the mass loss in accordance with Einstein’s E=mc2 equation. However, in
order for two or more protons to come together, they had to overcome the couloumb barrier, the
electromagnetic repulsion between like charges. This requires the protons to be tremendously
energetic, which in turn requires that they be in a very high temperature environment.
Eddington showed that the core of the sun was an environment with the necessary temperature.
He did this by reasoning that the sun was neither getting smaller or larger. For a fluid substance
such a condition is known as hydrostatic equilibrium. The force trying to collapse the sun is
gravity. Since the sun’s dimensions are not changing, the force of gravity must be counter
balanced by force acting in the opposite direction. There is a simple equation that relates the
pressure (force per unit area), P, of a gas to its density, D, and its temperature, T: P =DT.
Knowing the mass and volume of the sun and the temperature at its surface, Eddington was able
to calculate the temperature required at any point in the sun’s interior to produce the outward
pressure necessary to counter balance the inward gravitational pressure. He found that at its
core, the sun’s temperature would have to be around 10 million K. When Eddington first
published his results it was felt that this was not hot enough. However, further understanding of
the behavior of matter at the quantum level showed that the temperature was sufficient. Today
we recognize that the conversion of hydrogen (one proton) into helium (two protons + two
neutrons) plus energy at the core of the sun is the basic process that makes the sun shine. This
process is known generically as thermonuclear fusion.
Copied from the web site: http://www.kennethsnelson.net/atom/Portrait2.html
A BRIEF HISTORY OF ATOMIC STRUCTURE
450 B.C.
1678
1684
1687
1864
1873
1887
1895
1900
1905
1911
1913
DEMOCRITUS, a Greek philosopher, proposed that all matter is made up of particles
called atoms, meaning indivisible.
CHRISTIAN HUYGENS, postulated that light is a wave which moves and acts like
waves in water.
SIR ISAAC NEWTON stated that "matter is formed of solid, massy impenetrable
particles", of some definite size which combine in various ways to produce substance.
SIR ISAAC NEWTON developed the "corpuscular theory of light." Light is thought to
be the result of "luminous corpuscles" or particles which produce the waves we see
as light.
CLERK MAXWELL developed a series of equations expressing the relationship
between electric and magnetic forces.
CLERK MAXWELL stated "we have strong reason to conclude that light itself is an
electromagnetic disturbance."
HEINRICH HERTZ discovered the photoelectric effect. If a beam of light falls on a
clean metal plate in a vacuum, the plate becomes positively charged.
SIR JOSEPH THOMPSON proved the existence of a negatively charged particle,
termed the electron, which existed as part of the atom.
MAX PLANCK developed the basis of modern Quantum Theory by finding that light is
emitted or absorbed by an atom in discrete amounts called quanta.
ALBERT EINSTEIN in his explanation of the photoelectric effect proposed that light
must have both the properties of particles as well as those of waves.
LORD ERNEST RUTHERFORD discovered that the atom's nucleus is very small in
relation to the entire atom. He proposed that the negatively charged electrons were
revolving around a heavier, charged nucleus.
NIELS BOHR synthesized Rutherford's discovery into a reasonable model of an
actual atom, using hydrogen as his example. Bohr proposed a positively charged
central nucleus with electrons moving about it in circular orbits. The important feature
in Bohr's theory was that electron orbits could occur only in specific, predetermined
paths. If an electron absorbs energy, it is moved to an orbit further from the nucleus.
Conversely, when it drops to an orbit nearer the nucleus, it gives off energy in the
form of light. Different colors of light are produced depending on which orbit the
electron starts from and to which orbit it drops.
1916
ARNOLD SOMMERFELD proposed elliptical orbits in addition to Bohr's circular ones.
Sommerfeld's ellipses altered Bohr's model by showing electrons moving inwardly
and outwardly without radiating or absorbing energy.
1923
LOUIS DE BROGLIE proposed that all objects have properties of waves. The lighter
the object, the more pronounced the wave effect. An object as small as the electron
would act very much like a wave, forming stationary waves around the nucleus.
1925
1925
1926
1927
1929
WOLFGANG PAULI developed the Pauli Exclusion Principle which states that no two
electrons within the same atom can have the same set of quantum numbers.
UHLENBECK & GOUDSMIT showed that the electron possesses a spin in either
direction upon its axis.
ERWIN SCHROEDINGER developed an equation, based on de Broglie's wave idea,
expressing the probable location of an electron. These probable regions of occupancy
were conceived as clouds of charge around the nucleus. Different shapes occurred
for different types of orbitals.
WERNER HEISENBERG derived his "Uncertainty Principle" which states that it is
impossible to determine simultaneously the momentum and position of an electron.
LINUS PAULING showed how 2 electrons could form a more stable wave
arrangement if their spins were antiparallel.
CURRENTLY ACCEPTED SCIENTIFIC
DESCRIPTION OF AN ATOM
1.
At the center of the atom is a small, dense positively charged nucleus consisting
primarily of protons and neutrons.
2.
Moving around the nucleus are negatively charged electrons which account for only
1/5000 of the atom's mass -- the rest of the mass being in the nucleus. Most of the
atom is empty space. The motion of the electrons is not described.
3.
The electrons in an atom are allowed to have only certain energies. The allowed
states are described by a set of "quantum numbers", which indicate their average
distance from the nucleus, their angular momentum and its direction, and the
electrons' spin direction.
4.
Light of a specific color is emitted or absorbed when electrons change from one state
to another.
5.
The "Heisenberg Uncertainty Principle" states that the position and momentum of an
electron cannot be simultaneously determined. The interpretation of the Heisenberg
principle is that the atom's structure and the interactions of its electrons are random
and can be discussed only statistically.
6.
Even though the electron's exact position cannot be determined, if its energy is
known, the theory predicts the probability that an electron could be at a particular
place.
7.
If the probability location of an electron of known energy is plotted in space, the plot
looks like a fuzzy cloud of varying density, the shape varying with differences in
angular momentum. It always has a definite symmetry about the nucleus. Some of the
clouds or orbitals are spherical, others are like dumbbells, while others are more
complex.
8.
In describing an atom with many electrons, the charge clouds of one shell are
superimposed in space with those of other shells.
Sub-atomic particles: Protons, Neutrons and electrons
mass (in a.m.u.’s)
Type
symbol
charge
Location
Proton
P
positive
1
nucleus
Neutron
N
neutral
1
nucleus
Electron
e-
negative
0
shells/orbits
Atoms (or elements) are neutral combinations of protons, neutrons and electrons
Ions are atoms (or groups of atoms) that have gained or lost electrons to obtain a net charge.
Cations are positively charged ions (they have lost electrons)
Anions are negatively charged ions (that have gained one or more electron)
Isotopes are varieties of the same element that have different masses due to more or less neutrons in
the nucleus. Isotopes of the same element react the same. (heavy water is just like water only more
dense).
Element
Symbol
Mass
# of protons
# of neutrons
# of electrons
Carbon
C
12
6
6
6
Helium
He
4
2
2
N
7
28
13
S
18
79
24
79
12
Sn
69
Hydrogen
1
2
(tritium)
Uranium
235
47
Ir
77
240
94
125
84
On the back of this sheet draw picture/models of 4 (or more) elements
- Answers:
Element
Symbol
Mass
# of protons
# of neutrons
# of electrons
Carbon
C
12
6
6
6
Helium
He
4
2
2
2
nitrogen
N
14
7
7
7
aluminum
Al
28
13
15
13
sulfur
S
34
16
18
16
gold
Au
197
79
118
79
magnesium
Mg
24
12
12
12
Tin
Sn
119
50
69
50
Hydrogen
H or 3H
3
1
2
1
Uranium
U
235
92
143
92
silver
Ag
108
47
61
47
iridium
Ir
192
77
115
77
Plutonium
Pu
240
94
146
94
Polonium
Po
209
84
125
84
(tritium)
IONS chart
Name: __________________
Element/
period: ____
Symbol
Mass
# of protons
# of neutrons
# of electrons
C+4
12
6
6
2
H+1
1
0
0
ION
Carbon cation
Nitrogen anion
N-3
7
33
36
S-2
35
Mg+2
24
36
12
Sn+2
U+1
45
69
9
10
47
46
20
18
235
Ir-2
Cesium
54
On the back of this sheet draw picture/models of 4 (or more) IONS (2
cations, 2 anions)
Try these mixed elements and Ions
Element/
Symbol
Mass
# of protons
# of neutrons
# of electrons
12
10
ION
Al+3
He
Sodium
8
10
16
18
Os
Answers:
IONS chart
Element/
Symbol
Mass
# of protons
# of neutrons
# of electrons
Carbon cation
C+4
12
6
6
2
Hydrogen cation
H+1
1
1
0
0
Nitrogen anion
N-3
14
7
7
10
Arsenic anion
As-3
75
33
42
36
Sulfur anion
S-2
32
16
16
18
Bromine anion
Br -1
80
35
45
36
Mag. cation
Mg+2
24
12
12
10
Tin cation
Sn+2
119
50
69
48
Fluorine anion
F-1
19
9
10
10
Uranium cation
U+1
235
92
143
91
Silver cation
Ag+1
109
47
62
46
Iridium anion
Ir-2error
192
77
115
79
Calcium cation
Ca+2
40
20
20
18
Cesium cation
Cs+1
133
55
78
54
ION
On the back of this sheet draw picture/models of 4 (or more) IONS (2
cations, 2 anions)
Try these mixed elements and Ions
Element/
Symbol
Mass
# of protons
# of neutrons
# of electrons
ION
Aluminium cation
Al+3
27
13
14
10
Helium atom
He
4
2
2
2
Sodium cation
Na+1
23
11
12
10
Oxygen anion
O-2
16
8
8
10
Osmium atom
Os
190
76
114
76
Sulfur anion
S-2
32
16
16
18