Download 2.9 Periodic Trends

Unit 2: Structure of Matter
Content Outline: Periodic Trends associated with Electrons and Coulombs Forces (2.9)
I.
Coulomb’s Forces
A. These are the attractive forces that exist between two oppositely electrically charged particles, such
as positive protons and negatively charged electrons.
1. Remember, in the natural state, an atom has equal numbers of protons and electrons; therefore,
they are electrically neutral.
2. As the protons are larger in mass, if an atom has more protons than electrons, the electrons are
pulled inward and thereby possess less Potential Energy. See below 1 & 2
3. If the electron has more electrons than protons, the electrons are able to be farther from the
nucleus and thereby possess greater Potential Energy. See below 1 & 2
B. The formula for Coulomb’s Force is:
l q1q2 l
This equation is affected by 1) magnitude (number) of charges (q)
F = ke x r2
q1 would be the number of Protons; q2 number of electrons
2) Distance – r2
9
ke – is Coulomb’s constant = 8.99 x 10
II. Effective Nuclear Charge
A. This is symbolized as ZEFF
B. This is the net positive charge experienced by a single outer shell electron of an atom.
1. The Pull the electron “feels” by the nucleus.
2. Electron shielding
a. This term refers to the inner energy level of electrons helping to reduce (shield) the
outermost electrons from the full pull of the nucleus. Hence the term electron shielding.
i.
This allows the outermost electrons to remain farther away from the nucleus.
ii.
The more electrons an atom has, the greater the distance away from the nucleus
because of greater shielding.
III. The Periodic Trends in the Periodic Table
A. There are trends across a Period (row), referred to as Period Trends. There are trends within a
Family (column), referred to as Group Trends or Family Trends.
B. Atomic Radii Trends
1. Period Trends
a. Atomic Radii decreases as you go across (left to right) a period. This is due to greater
Coulombs forces or greater ZEFF. (Nucleus has greater pull on the electrons.)
2. Group Trends
a. Atomic Radii increases as you gown down a group. This is due to greater numbers of
electrons in higher energy levels being present and greater electron shielding. (Nucleus has
less pull.)
b. Exceptions 1) – Gallium. It is because Gallium has all the Group D protons, whereas Aluminum
does not. Therefore Gallium is small due to greater Coulombs forces or ZEFF. 2) At the hallway
point through the “d” and “f” blocks as each sublevel is half full. It resumes when the next
electron is added into a half filled orbital.
C. Ionization Energy Trends
1. Ion
a. This is a charged particle that is the result of gaining or losing an electron.
b. Cation
i.
These are positively charged particles/atoms that have lost electrons.
ii.
Because of the Law of Conservation of Matter, the electrons must have gone to
another location.
α. The Law states: Matter is neither created nor destroyed; just transferred and
transformed.
c. Anion
i.
These are negatively charged particles/atoms that have gained electrons.
d. Ionization
i.
This is the process of creating ions by gaining or losing electrons. (“tion” means
“process of”).
e. Ionization Energy (IE)
i.
This is the energy required to remove one electron from a neutral atom.
ii.
It is measured in kilojoules/mole (kJ/mol).
2. Period Trends
a. Group 1 Metals have the lowest Ionization Energies. (They lose electrons easily, hence their
being very reactive.)
b. Group 18 (Noble Gases) have the highest Ionization Energies on the whole Periodic Table.
(They do not lose electrons easily, hence their being unreactive.)
c. The Ionization Energies increase as you go across the period due to greater Coulomb Forces
or increased ZEFF. (The nucleus has a greater pull…requires more energy to remove.)
3. Group Trends
a. Ionization Energies generally decrease as you down a group due to lesser Coulomb Forces or
weaker ZEFF. (Electrons are farther away (more shielded) from the nucleus, so less energy is
needed to remove electrons.)
4. To remove more than one electron, even greater amounts of energy are needed than for the first.
This information can be found on an ionization table. With each electron removed, the greater
the Coulombs forces or increased ZEFF become within the atom.
D. Electron Affinity
1. This term refers to the energy change (positive or negative) to add an electron to an atom.
a. If the change is negative, energy has been released. (This is easy to do.)
b. If the change is positive, energy has been invested into the atom. (This is hard to do.)
i.
It requires a greater pushing force be put on the atom to overcome the negative
repulsive forces, thus making the atom very unstable. (This is what happens in your
cell phone… we use the greater electrical force to reverse the electron flow back to a
“charged state”.)
2. It is measured in kilojoules/mole (kJ/mol) too.
3. Period Trends
a. Group 17 (Halogens) has the greatest affinity as shown by large negative numbers.
b. In general, Electron Affinity increases as you go across a period.
c. Exception to the rule: Carbon – adding an electron allows the atom to reach a half filled p sublevel. That makes adding the fourth into a half filled sub-orbital more difficult as seen in N.
4. Group Trends
a. In general, Electron Affinity decreases (become more difficult) as you go down a group
because of greater Atomic radii and increased electron repulsive forces.
5. Adding a second electron is generally more difficult, unless it is a Group 16 element, such as
Oxygen. It needs 2 electrons to fill the valence shell.
E. Ionic Radii
1. Cations have decreased atomic radii due to greater Coulombs Forces or greater ZEFF. (The
electron cloud gets smaller.)
2. Anions have increased atomic radii due to greater electron shielding and no increase in ZEFF.
(The electron cloud gets larger.)
3. Period trends
a. Ionic Radii decreases until Group 15 and then begins to increase.
b. This is because Groups 1 – 14 are more Cationic. See above in F.1
c. Groups 15 – 17 are more Anionic. See above in F.2
4. Group Trends
a. Ionic Radii increases as you go down a group due to greater energy levels of electrons.
F. Electronegativity
1. This can be thought of as the desire to acquire negative electrons from another atom.
2. Fluorine (F) is the most electronegative element. (It has the strongest desire to be “like” a noble
gas…if it can just get that last electron in the valence shell.
3. Francium (Fr) is the least electronegative element. (It has very little desire to acquire an
electron. It would rather get rid of the solo electron in its valence shell and then be “like” the
Noble Gas Radon (Rn).)
4. Period Trend
a. Electronegativity generally increases going across a period.
5. Group Trend
a. Electronegativity decreases as you go down a group because the valence electrons are
farther from the nucleus. (There is more electron shielding, so it is less desirous to acquire
electrons.)