Atomic Structure and Periodic Properties
Objectives:
 Describe the discovery, location, and properties of subatomic particles.
 Identify isotopes and calculate the average atomic mass of elements.
 Describe the arrangement of electrons within an atom of any element and
correlate it to the element’s location on the periodic table.
 Summarize the historical development of the periodic table and compare/contrast
the original and modern table.
 Identify location and general properties of groups/families within the table.
 Identify periodic and group trends within the periodic table.
Atomic Theory – Historical
Around 1800 John Dalton proposed a 5 point theory intended to explain the structure of
the atom. His theory, although revolutionary, was based on ideas from Democritus.
Dalton’s Theory:
1. All matter is composed of extremely small particles called atoms.
2. All atoms of a given element are identical, having the same size, mass, and
chemical properties. Atoms of a specific element are different from those of any
other element.
3. Atoms cannot be created, divided into smaller particles, or destroyed.
4. Different atoms combine in simple whole-number ratios to form compounds.
5. In a chemical reaction, atoms are separated, combined, or rearranged.
A few of Dalton’s points aren’t exactly correct. For example, we know atoms can be
split. Also, atoms of the same element can vary slightly. But for the most part, his theory
is sound and truly remarkable considering the primitive research tools available at that
time.
Subatomic Particles
At one time the atom was thought to be solid. Research by Ernest Rutherford led to the
discovery that the atom was really mostly empty space. You can read about his
experiment on page 94 and 95 in the textbook.
Since the atom was mostly empty space, it was concluded that the central portion of the
atom (the nucleus) contained most of the mass and was positively charged. In an area
outside the nucleus a minute amount of mass would be found and contained a negative
charge (electrons). Electrons had been discovered by J.J. Thompson in the late 1890’s,
but their location was unknown. Thompson concluded these negatively charged particles
were located inside the atom much like chocolate chips are found in cookie dough. This
model of the atom was called the “Plum Pudding” model. Rutherford, as we have said,
determined this model to be incorrect.
Chadwick discovered the neutron in 1932.
Here’s what we know:
 The proton is positive, found in the nucleus of an atom, and has a mass of 1.
 The neutron is neutral, also found in the nucleus, and has a mass of 1.
 The electron is negative, is found outside the nucleus and has a very small mass.
Its mass has been calculated to be 1/1840 of a proton or neutron. This mass is so
insignificant compared to the mass of the other subatomic particles (protons and
neutrons) we can ignore their mass.
Atoms differ in several ways.
 Their atomic number (number of protons) gives the atom its identity. If we
change the number of protons, we change the atoms to a totally different atom.
 Their mass number (total # of protons and neurons) can changed by gaining or
losing neutrons.
 They can gain or lose electrons becoming negative or positive ions, respectively.
It is important to note here that in an atom, unless it is indicated differently, the
number of protons (positive charges) is equal to the number of electrons (negative
charges). In other words, the atom is electrically neutral.
Definitions:
 Ion – an atom that has a positive or negative charge.
 Isotope – an atom that has gained or lost neutrons (changed its mass number).
Remember this formula:
Mass Number = Number of protons + Number of neutrons
How is the periodic table arranged?
The periodic table is arranged by increasing atomic number (number of protons). Note
that the periodic table is NOT one long string of elements. Instead it is arranged in
columns and rows. The columns are called groups or families and those elements in the
same group have similar, but not identical chemical and bonding properties. The rows
are called periods. One period will have the same sequence of similar properties as the
row above and below it.
So what does the number on the periodic table with all the decimals stand for?
This is the average atomic mass. If we were to average the number of students in each
of the classrooms here at WHS we would most likely come up with a number having a
decimal. As we have said before, atoms of the same element can vary their mass number.
In a lump of an element a certain percentage of those atoms will have a certain mass
number, another percentage of those atoms will have another and so on.
So how do I calculate the average atomic mass?
Simple!
Relative abundance of isotope #1 X the mass number of that isotope.
+
Relative abundance of isotope #2 X the mass number of that isotope.
Example:
Boron has two naturally occurring isotopes: boron-10 and boron-11.
Their relative abundance is 19.8% and 80.2% respectively.
First, convert the percentages to decimals,
Then multiply the relative abundance times their mass numbers.
Add the values together and you have the average atomic mass.
0.198 X 10 = 1.98
+
0.802 X 11 = 8.82
10.80 = the average atomic mass
Electrons in atoms.
It has been found that electrons exist in levels or shells surrounding the nucleus of an
atom. Their paths are very complex at times and the exact location of any electron can
never be strictly determined.
Here is the breakdown of the areas:
 Level – this can be thought of as a general distance away from the nucleus.
 Sublevel – although not exactly correct, but it works for us, a sublevel can be
thought of as a pattern within a level.
 Orbit – I like to think of orbits as pathways or directions of travel with the
sublevel.
There are 3 principles governing the location of electrons:
 Aufbau Principle – an electron will occupy the lowest energy level available.
This means the lower levels will fill first. Water fills a glass from the bottom up.
 Pauli Exclusion Principle – there will be a maximum of 2 electrons per orbital.
 Hund’s Rule – Each orbital a sublevel must have an electron before any orbit can
contain 2. I call this the “blackjack” rule.
Level Sublevel
# of Orbits
Max # e- per sublevel Max # e- per level
1
s
1
2
2
-----------------------------------------------------------------------------------------2
s
1
2
p
3
6
8
-----------------------------------------------------------------------------------------3
s
1
2
p
3
6
d
5
10
18
-----------------------------------------------------------------------------------------4
s
1
2
p
3
6
d
5
10
f
7
14
32
Life would be so simple if the electrons filled straight down the list. We couldn’t be so
lucky. Although sublevel 3d is closer to the nucleus than 4s, 3d requires more energy for
an electron to “live” there than to “live” on 4s. The diagram below will give you the
sequence or order with which the electrons fill.
1s
2s
2p
3s
3p
3d
4s
4p
4d
5s
5p
5d
6s
6p
7s
4f
Although all electrons are important, those found in the outside level, the valence
electrons, actually determine the bonding characteristics of that element.
Formation of Ions
Atoms will tend to try to gain a stable electron configuration. The will either gain
electrons to fill up the level or lose electrons to reveal the next lower full level.
Take oxygen for example. Oxygen has 6 electrons in its outside level. It is much easier
to gain 2 electrons than to lose 6 electrons. It will gain 2 electrons, thus filling the
outside level and have a negative 2 charge. O-2.
Magnesium has 2 electrons in its outside level. It is easier to lose 2 electrons than to gain
6 electrons. It will lose 2 electrons and have a positive 2 charge. Mg+2.
Periodic Properties
Periodic Properties are those properties that occur as we move from left to right across
the periodic chart within a particular period. Periodic (and Group) Trends are controlled
by the degree of “positiveness” of the nucleus and that attraction for the valence
electrons.
Atomic Radii
The actual size of the atom decreases as we move to the right across the periods. As we
move to the right we increase 1 proton each atom. The energy level occupied by the
valence electrons does not change. This results in an increased attraction by the nucleus
on the valence electrons causing them to be pulled slightly toward the nucleus. THIS
SAME TENDENCY IS RESPONISBLE AND CONTROLS THE REMAING
PERIODIC TREND AS WELL..
Electron Affinity
Electron affinity is an atom’s tendency to gain electrons. First, in order to gain electrons,
there must be a space available for adding an electron. This eliminates the Noble Gases
because they have no space available to gain electrons. Since the nucleus becomes more
positive as we move right and the attraction for the valence electrons increases this also
increases the attractive forces required to add and additional electron. For this reason
those elements on the right side of the Periodic Table are anions.
Electronegativity
Electronegativity is an atom’s pull on another atom’s electrons in a chemical bond. Let’s
look and the water molecule. Hydrogen and oxygen are bonded together and are sharing
one another’s electrons. Oxygen pulls harder on hydrogen’s electron than hydrogen pulls
on oxygen’s electron. This results in an unequal sharing. As we move up and to the
right on the Periodic Table the electronegativity increases.
Ionization energy
Ionization energy is the force needed to remove an electron from an atom. As we move
up and to the right on the Periodic Table the ionization energy increases.