Periodic Table Trends
AP Chemistry
Ms. Grobsky
• Each student will address the forces that create each trend
seen in the periodic table
Objectives
• A trend is an observation, NOT an explanation!
• It is fine to state the trend in your answers, but you must
also go further by explaining what causes the observed
trend
• Be sure to mention BOTH the atoms or ions in the
question asked
Some Things to Keep in
Mind
• Generalizations
• Nuclear effective charge (Zeff) justifies trends across a
period
• Zeff is defined as the amount of positive charge on the nucleus
perceived by an electron
• Electrons intervening between the nucleus and an outer electron are
said to “shield” or screen the outer electron from the nucleus so that
the outer electron does not experience the full nuclear charge
• Increased distance from nucleus (greater value of n)
justifies trends down a group
Justification of the Trends
• Distance between the nucleus and the outer edge of the
electron cloud
• Measured by dividing the distance between two atomic
nuclei by 2 since the outer edge of the electron cloud cannot
be deduced
• Influenced by the nuclear pull and the number of energy
levels
What is Atomic Radius?
• Effective nuclear charge, Zeff, increases the attraction of
the nucleus and therefore, pulls the electron cloud closer
to the nucleus
• Smaller atomic radius across a period
• Left to right
Atomic Radius Trend
Across a Period
• Increased number of energy levels (n) increases the
distance over which the nucleus must pull and therefore,
reduces the attraction for electrons
• Larger atomic radius down a group
• Note:
• Full energy levels provide some shielding between the
nucleus and valence electrons
• Does NOT explain the trend across a period
• Only full energy levels, not full sublevels, are of concern in a
shielding argument
Atomic Radius Trend
Down a Group
• Ionization energy (IE)
• Energy needed to remove an electron from a gaseous atom
or ion
• An isolated atom or ion, NOT part of a solid, liquid, or
molecule
• Always endothermic
• Stepwise process
• Highest-energy electron (one bound least tightly) is always
removed first
• First ionization energy
What is Ionization
Energy?
• Effective nuclear charge, Zeff, increases the attraction of the nucleus
and therefore, pulls the electron cloud closer to the nucleus
• Holds electrons more tightly
• Increases left to right across a period
• Exceptions:
• A drop in IE occurs between groups 2 and 13
• p electrons do not penetrate the nuclear region as greatly as s electrons do
and are therefore not as tightly held
• Does NOT mean that p electrons are farther away from nucleus
• A drop in IE occurs between groups 15 and 16
• Increased repulsion created by the first pairing of electrons outweighs the
increase in Zeff
• Thus, less energy is required to remove the electron
• Does NOT mean group 15 is more stable because they have a half-filled sublevel
Ionization Energy Trend
Across a Period
• Increased number of energy levels (n) increases the
distance over which the nucleus must pull and therefore,
reduces the attraction for electrons
• Smaller ionization energy down a group
• Note:
• Full energy levels provide some shielding between the
nucleus and valence electrons
Ionization Energy Trend
Down a Group
• Electron affinity
• Energy change associated with the addition of an electron to a gaseous atom
or ion
• Can be exothermic or endothermic
• NOT the opposite of ionization energy
• If energy change is exothermic, consider the absolute value of the
energy term
• Negative sign is simply indicating the direction of energy flow out of the
system
What is Electron Affinity?
• Measures the attraction of an atom of the pair of outer
shell electrons in a covalent bond with another atom
What is
Electronegativity?
• Effective nuclear charge, Zeff, increases the attraction of
the nucleus and therefore, strengthens the attraction for
electrons
• Holds electrons more tightly
• Increases left to right across a period
Electronegativity Trend
Across a Period
• Increased number of energy levels (n) increases the
distance over which the nucleus must pull and therefore,
reduces the attraction for electrons
• Smaller electronegativity down a group
• Note:
• Full energy levels provide some shielding between the
nucleus and valence electrons
Electronegativity Trend
Down a Group
• Distance from the nucleus to the other edge of the electron
cloud in a charged ion
• Same radii trends apply once you divide the table into metal
and non-metal sections
• Metal section
• Positive ionic radii decrease from left to right with only minor
changes in the transition metals
• Non-metal section
• Negative ionic radii decreases from left to right
• Ionic radii increases down all groups because of the additional
energy levels (n)
What is Ionic Radius?
• Positive metal ions result from the loss of valence
electrons
• In many cases, this means that the farthest electrons are now
in a smaller principal energy level (n) than the original
neutral atom
• Address the ratio of protons to electrons
• As electrons are lost, the ratio of p+/e- increases
• Thus, the electrons are held closer and with more strength
Positive Ions and Ionic
Radii
• Larger than cations
• Negative non-metal ions result from the addition of
valence electrons
• Results in change in the proton to electron ratio
• As electrons are added, the p+/e- ratio decreases and the
electrons are not as closely held
• Increased electrons and electron repulsions also play a
role in expanding the electron cloud
Negative Ions and Ionic
Radii
• Depends on whether the element reacts by losing
electrons (metals) or gaining electrons (non-metals)
What is Reactivity?
• Due to the fact that metals react by losing electrons, a
loosely held electron will result in a more reactive metal
• Directly tied to ionization energy
• Remember, ionization energy is measured using gaseous atoms,
whereas most chemical reactions for metals occur in an
aqueous solution or between solids
• Therefore, only looking at ionization energies allows us to make
only general predictions about the reactivity of metals
• With an increased number of energy levels (n) comes
increased distance from the nuclear attraction
• Thus, a more loosely held electron for reaction
Metals Reactivity
• Due to the fact that non-metals tend to gain electrons, a
strong nuclear attraction will result in a more reactive
non-metal
• An atom with the highest Zeff and the least number of
energy levels should be the most reactive non-metal
• F because its nucleus exerts the strongest pull!
Non-metals Reactivity
• Follow these three steps EVERY time you answer a
periodicity question involving two or more species:
• Locate all elements on the periodic table and state the
principal energy level (n) and the sublevel containing the
valence electrons for each element
• Do they have the same or different n values?
• If same n, argue periodicity using Zeff
• If different n, argue periodicity with n vs. n distances from
nucleus
Final Thoughts