Topic 2: Atomic Theory
2.1 The atom
2.1.1. State the relative mass and relative charge of protons, electrons and neutrons.
Relative Mass
Charge
Proton
1
+1
Neutron
1
0
Electron
1/1840
-1
2.1.2. State the position of protons, neutrons and electrons in the atom.
The simplest view of the atom is that it consists of a tiny nucleus (with a diameter of 10-13
cm( and electrons that move about the nucleus at an average distance of about 10-8 cm
from it. Since chemistry is mainly focused on the electrons, we accept a relatively crude
view of the nucleus. Basically, the nucleus is assumed to contain protons, which have a
positive charge equal to the magnitude to the electron’s negative charge, and neutrons,
which have virtually the same mass as a proton but no charge. The electrons comprise
most of the atomic volume and thus are the parts that “intermingle” when atoms combine
to form molecules.
2.1.3. Define the terms mass number (A), atomic number (Z) and isotope.
The mass number is the total number of nucleons, or particles in the nucleus. These
include neutrons and protons. Since both neutrons and protons have a mass of roughly 1
amu, the atomic mass number is assumed to be the integer closest to the atomic mass
shown on the periodic table. Atomic number is just the number of protons in the nucleus,
and this is how we identify different elements. An atom with 12 protons in the nucleus is
ALWAYS carbon, regardless of how many electrons or neutrons the atom has. This is
also on the periodic table. Isotopes are atoms with the same number of protons but
different numbers of neutrons.
2.1.4. State the symbol for an isotope given its mass number and atomic number.
Use the notation AZX, eg 126C.
An isotope is written with the mass number on top of the atomic number, then the symbol
for the element, as you can see from Carbon 12 written above.
2.1.5. Explain how the isotopes of an element differ.
Isotopes have the same chemical properties but different physical properties. Examples
such as 11H, 21H, 31H, 126C, 147C, 3517Cl and 3717Cl should be considered.
Isotopes are chemically basically identical, however they have slightly different physical
properties, and this is how they are separated from each other. For example, hydrogen is
slightly lighter then deuterium (2H) and tritium is actually radioactive.
2.1.6. Calculate and explain non-integer atomic masses from the relative abundance
of isotopes.
The reason atomic masses are non-integer on the periodic table is because in natural
samples of an element, there are different isotopes and different isotopes have different
atomic masses. Take hydrogen for example. 1H is the normal, regular isotope of
hydrogen that know so well, its atomic mass is 1, and in a sample of hydrogen from the
natural world, this isotope would make up 99.98% of the hydrogen in the sample.
Deuterium however would make up the other .02%, and its atomic mass is 2. Then there
is tritium which is one in every 10,000, so it is neglected in this case. In order to find the
true atomic mass of hydrogen, we have to take into account the fact that protium (regular
hydrogen) is not the only isotope in nature. So, the way you do this is you multiply the
atomic mass of the isotope by the percentage in decimal form and you add them together.
So for protium, the percentage is .9998(1 amu) + protium which is .0002(2 amu)= 1.0079
amu. This is the non-integer atomic mass which we find on our periodic table.
2.1.7. Calculate the number of protons, electrons and neutrons in atoms and ions
from the mass number, atomic number and charge.
The number of protons, electrons and neutrons in an atom are all closely related through
the mass number, atomic number, and charge values. First of all, if you subtract from the
mass number the number of protons, you will get the amount of neutrons because the
mass number is the protons and neutrons added together. In a normal atom, the number
of protons and electrons are equal because an atom has an overall neutral charge, which
means the amount of protons must be equal the amount of electrons so that their charges
cancel out. Yet sometimes you have substances or things that are ionic, and thus do not
have a neutral charge. When you have such a molecule (it’s not really a molecule but it’s
easy to call it that), the numbers of protons and electrons don’t add up. Say you have an
atom that has a negative 2 charge and has 6 protons in the nucleus. Since it’s charge is
negative, that means the atom has more electrons then protons, and since it is negative
two, that means it has two more electrons then protons. So that means that its 6 electrons
plus two more to equal the negative two charge, so you have a negative 8 charge. Say
you have a atom that has a positive 3 charge, and you have 6 protons in the nucleus. You
know that you have three less electrons then protons since you have a positive three
charge, so you subtract three from six to find that the atom has 3 electrons.
2.2 Electron Arrangement
2.2.1. Describe and explain the difference between a continuous spectrum and a line
spectrum.
A continuous spectrum is created when white light is passed through a prism. This
spectrum, like the rainbow produced when sunlight is dispersed by raindrops, contains all
the wavelengths of visible light. In contrast, a line spectrum is when only a few lines are
emitted rather then the entire spectrum.
2.2.2. Explain how the lines in the emission spectrum of hydrogen are related to the
energy levels of electrons.
Students should be able to draw an energy-level diagram, show transitions between
different energy levels and recognize that the lines in a line spectrum are directly related
to these differences. An understanding of convergence is expected. Series should be
considered in the ultraviolet, visible and infrared regions of the spectrum. Calculations,
knowledge of quantum numbers and historical references are not required.
When a sample of hydrogen gas received a high-energy spark, the H2 molecules absorb
energy, and some of the H-H bonds are broken. The resulting hydrogen atoms are
excited; that is, they contain excess energy, which they release by emitting light of
various wavelengths to produce what is called the emission spectrum of the hydrogen
atom. This is a line spectrum. This line spectrum indicated that only certain energies are
allowed for the electron in the hydrogen atom. In other words, the energy of the electron
in the hydrogen atom is quantized. This observation ties in perfectly with the postulates
of Max Plank. Changes in energy between discrete energy levels in hydrogen will
produce only certain wave-lengths of emitted light. For example, a given change in
energy from a high to a lower level would give a wavelength of light that can be
calculated from Plank’s equation (but which we don’t have to know.) The discrete line
spectrum of hydrogen shows that only certain energies are possible; that is, the electron
energy levels are quantized. In contrast, if any energy level were allowed, the emission
spectrum would be continuous. This means, electrons can only be at certain energy
levels around a nucleus. Convergence is a mathematical term for a series of numbers that
gradually decrease, but never quite get to 0, and add up to infinity (but technically reach
some limit). Not completely sure how this applies to this, the energy levels are an
example of convergence I guess.
2.2.3. Describe the electron arrangement of atoms in terms of main energy levels.
Students should know the maximum number of electrons that can occupy a main energy
level (up to Z = 18). No knowledge of sublevels s, p, d and f is required. The term
valence electron is used to describe the electrons in the highest main energy level.
Energy levels are the levels around the nucleus that were discussed and shown by the
different wavelengths of light in the light spectrum. There are 7 energy levels around an
atom, although they are not always nor nearly always filled. Up until you reach the point
where you have more then 18 protons in the nucleus, the maximum number of electrons
you can have in an energy level is 8 (the exception to this rule is the first energy level
which can only have two electrons in it). When one energy level is filled, the electrons
move on to the next highest energy level. The highest energy level that is not yet filled
with 8 electrons is known as the valence energy level, and the electrons in that energy
level are known as valence electrons. So, say you had the first two energy levels filled
and then had 3 more electrons in the third energy level. Those three electrons would be
considered the valence electrons for that atom. Energy levels are arranged around the
nucleus in varying distances (the lower the energy level, the closer it is to the nucleus).
Energy levels are represented on the periodic table by rows. All of the elements in the
first row are atoms that only have electrons in the first energy level. All of the elements
in the second row are elements that have filled the first energy level and have electrons in
the second. All of the elements in the third row are elements that have filled the first two
energy levels and have elements in the third.
2.2.4. Determine the electron arrangement up to Z = 20.
For example, 2,8,7 or 2.8.7 for Z=17.
An electron arrangement can be organized by writing how many electrons are in each
energy level, and the energy level is separated by either a comma or a period. So,
hydrogen, which has one electron in the first energy level, is 1. Helium is 2. Lithium has
the first energy level filled and then has one in the second, so it is 2,1. Berylium is 2,2.
This pattern continues until electrons have filled the second energy level. Then, you go
onto sodium which is 2,8,1 because it has two electrons in the first, 8 in the second, and
then one in the third. This continues all the way to 2,8,8,2 for Calcium. Then something
special happens that will not be discussed until later in the syllabus.