Study Guides
Big Picture
The periodic table allows us to see trends in different groups of elements.
Key Terms
Period: Horizontal row on the periodic table.
Electron Shielding: The force caused by the inner
Group: Vertical column on the periodic table.
electrons of an atom that pushes valence electrons
Ionization Energy: The energy needed to remove an
electron from an atom.
Electronegativity: How strongly an atom attracts
electrons.
Electron Affinity: The energy released when an
electron is added to an atom.
Atomic Radius: One half of the distance between the
nuclei of two atoms of the same element when the
atoms are joined.
away from the center of the atom.
Effective Nuclear Charge: The combined effect of the
nuclear charge and electron shielding. The effective
nuclear charge affects all periodic trends.
Valence Electron: Electron in the outermost energy
level (highest n number). Number of valence electrons increases as you go across the periodic table.
Ion: An atom with a positive or negative charge.
Cation: A positively charged ion that contains
Nuclear Charge: The positive force exerted by the
nucleus on the valence electrons, pulling them
Chemistry
Atomic Trends
more protons than electrons.
Anion: A negatively charged ion that contains
towards the center of the atom.
more electrons than protons.
Important Periodicity Concepts
The periodic table contains many trends:
• Atomic
size and ionic size increase as you go from right to left across a period and from top to bottom within a
group.
• Ionization energy increases as you go from left to right and from bottom to top.
• Electronegativity increases as you go from left to right and from bottom to top.
• Electron affinity increases as you go from left to right and from bottom to top.
This guide was created by Steven Lai, Rory Runser, and Jin Yu. To learn more
about the student authors, visit http://www.ck12.org/about/about-us/team/
interns.
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v1.1.12.2012
Disclaimer: this study guide was not created to replace
your textbook and is for classroom or individual use only.
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Chemistry
Atomc Trends
cont .
Atomic Size
Atoms do not have a well-defined boundary, so the radius of an atom cannot be measured directly. There are several
ways to define atomic size, so there is some variations in sizes.
• One
definition of atomic radius is to take half of the distance between the nuclei of two atoms of the the same
element that are joined together.
• No
matter the definition, atoms are extremely small. The radius is often measured in picometers (10-12 m) or
angstroms (10-10 m), symbolized by Å.
In general, the atomic size increases as you go from right to left across a period and from top to bottom within a group.
Going down a group:
• The nuclear charge increases as the atomic number increases–a greater positive charge pulls the electrons closer
to the nucleus
• The number of core electrons increases with
each successive period–weakens the nuclear charge due to electron-
electron repulsion
• This is called the shielding effect and decreases the attractive force from the nuclear charge
• Due to shielding, the effective nuclear charge is less than the actual nuclear charge
• The increase in the shielding effect is greater than the increase in the nuclear charge, so
the atomic size
increases
Going to the right in a period:
• The nuclear charge increases as the atomic number increases
• The number of valence electrons increases, but the number of core electrons is constant
• The increase in effective nuclear charge causes the atomic size to decrease
Ion Size
An ion is an atom that has lost or gained electrons and is no longer electrically neutral.
• A cation is smaller than the original neutral atom.
• By losing an electron, the attraction between
the nucleus and the remaining electrons increases while the
number of shielding electrons remain the same.
• An anion is larger than the original neutral atom.
• By gaining an electron, the attraction between the nucleus and the remaining electrons decreases.
Cation sounds like cat, which is a nice thing (positive). Anion sounds like onion, which makes you cry (negative).
The trend for ionic size follows the trend for atomic size, as ions on the left side of the table are smaller and ions on
the right side of the table are larger than the neutral atoms.
Ionization Energy
To remove an electron, the atom needs to absorb enough energy to overcome the attraction of the positively charged
nucleus.
In general, this ionization energy increases as you go from left to right across a period and from bottom to top within
a group.
• Going
down a group, the atomic size increases as the effective nuclear charge decreases, making it easier to
remove an electron
• Going to the right in a period, the atomic size decreases as the effective nuclear charge increases, making it harder
to remove an electron
More than one electron can be removed, but each subsequent ionization energy requires more energy than the
previous because the nucleus becomes more positive. The ionization energy after removing the last valence electron
is much larger than the previous one because the atom behaves like a noble gas (noble gas has very high stability
because all the sublevels are filled).
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cont .
Electronegativity
Elements can chemically combine to form stable compounds by forming chemical bonds. Electronegativity characterizes
the ability for an atom in a chemical bond to attract an electron.
In general, electronegativity increases as you go from left to right across a period and from bottom to top within a
group.
The trends for electronegativity follow the trends for ionization energy. Ionization energy is one of the factors used to
determine electronegativity.
• The scale for electronegativity is arbitrary. The most widely used scale is the Pauling scale developed by Linus
Pauling. Fluorine is listed as the most electronegative element with an electronegativity of 4.0.
• Noble gases are typically listed as not having any electronegativity and are often left off of electronegativity charts.
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Electron Affinity
An atom gives off energy when it gains an electron. Electron affinity is a measure of the energy released when an
electron is added to a gaseous atom or ion.
In general, electron affinity increases as you go from left to right across a period and from bottom to top within a group.
The trends for electron affinity follow the trends for ionization energy.
• When atoms become ions, the atoms either release energy (electron affinity) or absorb energy (ionization energy).
Atoms that require a large amount of energy to release an electron will likely give off the most energy while
accepting an electron.
• The exception is noble gases. Noble gases have essentially zero electron affinity.
Notes
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Chemistry
Atomic Trends