Download atomic radii

Trends of the
Periodic Table
GPS Standards:
SC4. Students will use the organization of the Periodic Table to predict properties of elements.
a. Use the Periodic Table to predict periodic trends including atomic radii, ionic radii, ionization
energy, and electronegativity of various elements.
1/21/10
1. Atomic Radii
• atomic radii (atom size) is determined by
measuring the distance between the nuclei of
2 identical atoms that are chemically bonded
together, then dividing this distance by 2.
• measured in pm (picometers, or 10-12 m).
• across a period, atomic radii tends to
decrease from left to right.
• this is due mostly from the increase in proton number.
• as electrons are added to the s and p sublevels, they are pulled
closer and closer to the highly positive nucleus, which results in a
decrease in atomic radii.
• down any group on the table, the atomic radii will increase.
• this is due to an increase in total number of particles (protons,
neutrons, and especially electrons) in the atom. The number of
orbitals needed increases, so the overall size increases.
• this is due to an increase in total number of particles (protons,
neutrons, and especially electrons) in the atom. The number of
orbitals needed increases, so the overall size increases.
• also, the electrons that are added to these larger elements are
being added in orbitals that lie farther and farther away from the
nucleus, and thus, father away from any positive charges that
would bring them closer inside and decrease the radii.
• thus, the overall trend line for atomic
radii looks like the figure to the right.
• the placement of the arrowhead in
the bottom left corner of the table
indicates that you will find the largest
elements there.
• note that fluorine (F) is one of the smallest elements.
• let’s take a look at the actual values for atomic radii........
2. Ionization Energy
• ionization energy is the energy required to remove one electron
from an atom, forcing the creation of a positive ion.
• for example, if you were to remove the one valence electron
that sodium (Na) has in its outermost s and p orbitals, you would
create an Na+ ion.
• any process that refers to the creation of an ion is called
ionization.
• IE is measured in kilojoules per mole (kJ/mol), which is a unit that
represents energy required per amount of substance.
• across a period from left to right, IE will generally increase.
• this is due mostly to the number of valence electrons that can
be found in their outermost s and p orbitals.
• the lower the number, the easier it is to lose them.
• therefore, the alkali metals will lose electrons very easily
(only one to lose) and have low IE, and.......
• therefore, the alkali metals will lose electrons very easily (only
one to lose) and have low IE, and.......
• the noble gases (with 8 valence e-) will hold onto those electrons
for dear life, and have very high IE.
• this explains why the alkali metals are very reactive with most
substances, and noble gases are the most unreactive elements.
• the reason that lower valence electron number equals lower IE is
because of a lower number of protons are present in the nucleus on
the left side of the table when compared with the right side.
• therefore, there is less positive charge in the nucleus to hold those
electrons in towards the nucleus, and they are more easily lost for
elements on the left side of the table.
• down a group, IE will tend to decrease.
• just as atomic radii was affected by the electrons being added to s
and p orbitals that were farther and farther away from the nucleus, IE
is affected in the same way, leading to an easier removal for
electrons that lie in the outermost, most distant s and p orbitals.
• also, you must realize that electrons that lie in between the nucleus
and those outermost s and p electrons will tend to shield the effect
of the nuclear charge from those outermost electrons.
• thus, the overall trend line for IE looks
like the figure to the right.
• the placement of the arrowhead in the
top right corner of the table indicates
that you will find the elements with the
largest IE there.
• note that fluorine is the element that has the highest IE (other
than the noble gases).
• let’s take a look at the actual values for IE........
3. Ionic Radii
• ionic radii (ion size) is the new size of an atom after it has either
gained or lost one or more electrons, forming an ion.
• there are two kinds of ions: cations and anions.
• positive ions are known as cations.
• the formation of a cation (by the loss of one or
more electrons) always leads to a decrease in
atomic radius because the removal of the
highest-energy-level electrons results in a
smaller electron cloud.
• also, the remaining electrons are drawn closer
to the nucleus by its unbalanced positive charge.
• negative ions are known as anions.
• the formation of an anion (by the addition of one or more
electrons) always leads to an increase in atomic radius.
• negative ions are known as anions.
• the formation of an anion (by the addition of one or
more electrons) always leads to an increase in
atomic radius.
• this is because the total positive charge of the nucleus
remains unchanged when an electron is added to an
atom or an ion.
• so the electrons are not drawn to the nucleus as
strongly as they were before the addition of the extra electron.
• the electron cloud also spreads out because of greater
repulsion between the increased number of electrons.
• within each period of the periodic table, the metals at the left tend
to form cations and the nonmetals at the upper right tend to
form anions.
• within each period of the periodic table, the metals at the left tend
to form cations and the nonmetals at the upper right tend to
form anions.
• cationic radii decrease across a period because the electron
cloud shrinks due to the increasing nuclear charge acting on
the electrons in the same main energy level.
• anionic radii decrease across each period for the elements in the
upper p-block. The reasons for this trend are the same as the
reasons that cationic radii decrease from left to right.
• as they are in atoms, the outer electrons in both cations and anions
are in higher energy levels as one reads down a group.
• therefore, just as there is a gradual increase of atomic radii
down a group, there is also a gradual increase of ionic radii.
• thus, the overall trend line for ionic radii
looks like the figure to the right.
• the placement of the main arrowhead
in the bottom left corner of the
table indicates that you will find the
elements with the largest ionic radii there.
• note that fluorine is an element with a very small anionic radii.
Aluminum is an element with a very small cationic radii.
4. Electronegativity
• electronegativity is a measure of the ability of an atom in a chemical
compound to attract electrons.
• remember that valence electrons are the only electrons
involved in bonding.
• some elements are better able than others to attract the
valence electrons involved in the bonding process.
• an uneven concentration of charge in a compound affects the
chemical properties of that compound.
• if electrons within a compound are being pulled unevenly in one
direction, the molecule is said to be “ionic” or “polar”
(depending on what elements are involved).
• if there is an equal distribution of electrons, and thus, charge, on a
molecule, then it is said to be “non-polar.”
• Linus Pauling, one of America’s greatest chemical
minds, created an electronegativity chart to relate
elements to each other.
• he assigned fluorine, the most electronegative
element, a value of 4.0.
• values were assigned to all the other elements
in relation to this value.
Pauling
• across a period, electronegativity values increase
from left to right., and down a group, they tend to decrease.
• thus the trend line for electronegativity looks like the figure to the right.
• the placement of the arrowhead in
the top right corner of the table
indicates that you will find the elements
with the largest electronegativities there.
• notice fluorine is the element with the most electronegativity.
• Let’s take a look at the actual values for electronegativity........