Transcript Periodic Trends - The Woodlands College Park High School

Periodic Trends
Predicting Periodic Trends
• A number of physical and chemical properties of
elements can be predicted from their position in the
periodic table. Among these properties are Ionization
Energy, Electronegativity, and Atomic/Ionic Radius.
• These properties all involve the outer shell (valence)
electrons as well as the inner shell (shielding)
electrons.
• Electrons closest to the nucleus experience the full
nuclear charge and are held most strongly.
• As the number of electrons between the nucleus and
valence electrons increases, the nuclear charge
decreases, due to the “shielding” of these inner shell
electrons.
Predicting Periodic Trends
Cont.
• The charge felt by the valence electrons is called the
effective nuclear charge, Zeff.
• Going down a group increase the value of n, and
increases the number of inner shell electrons. This
leads to better shielding and a weaker attraction
between the nucleus and the outer shell electrons.
• Going across a period leads to a larger nuclear
charge, as the number of protons increases. There
is also increase in the number of valence electrons,
but electrons in the same shell are poor at shielding
each other. Going across a period generally leads
to a stronger interaction between the nucleus and
valence electrons.
Atomic Radius
• One-half the distance between the two
nuclei of two identical atoms.
• Period trend: decreases as you move
across the period (horizontal). Due to
the increase in atomic charge (Z) that
draws in the electron cloud.
• Group trend: increases as you move
down the group (vertical). As n
increases, the sizes of the orbitals
increases.
Trends in atomic radius in
Periods 2 and 3
Atomic Radius
• Arrange the following atoms in order of
decreasing atomic radius.
Na
Al
P
Cl
Mg
• Which is the largest atom in Group 4?
• Which is the smallest atom in Group 7?
• Which is the smallest atom in Period 5?
Ionic Radius
• Ion: an atom that has a positive or negative charge. Atoms lose
or gain electrons to become ions.
• Ionic radius is the radius of an anion or cation. When neutral
atoms are ionized, there is a change in their sizes.
• If the atom forms an anion the size increases, because the
nuclear charge (Z) is unchanged but the electron-electron
repulsion increases, due to the added electron, enlarging the
electron cloud.
• If the atom forms a cation, the size decreases. The nuclear
charge is unchanged and the decreased electron-electron
repulsion shrinks the electron cloud.
Ionic Radius Cont.
• Periodic trend: decrease across period
– Groups 1-13: decrease across the period (protons
pull in the electron cloud)
– Group 14: decrease (protons pull in the electron
cloud) or increase (added repulsion in electron
cloud increases volume)
– Group 15: large increase
– Groups 16-18: decrease across the period
Ionic Radius Cont.
• Which of the two species is larger?
– N3- or F– Mg 2+ or Ca 2+
• Which of the two species is smaller?
– K+ or Li+
– P3- or N3-
Ionization Energy
• Energy required to remove one electron. Related to how
“tightly” the electron is held by the nucleus. The higher
the IE, the more difficult it is to remove the electron.
• Period trend: increases across a period. Electrons in
the same set of orbitals do not shield each other very
well but the nuclear charge increases, making the
electrons more difficult to remove.
• Group trend: decreases down a group. Electrons
become increasingly easy to remove since they are at an
increasing distance from the nucleus, with increasing
numbers of shielding inner electrons.
Ionization Energy
• Ionization energies are measured in kJ
mol-1 (kilojoules per mole). They vary in
size from 381 (which you would consider
very low) up to 2370 (which is very high).
Ionization Energy
• The first ionization energy is the energy
required to remove the most loosely held
electron from one mole of gaseous atoms to
produce 1 mole of gaseous ions each with a
charge of 1+.
• This is more easily seen in symbol terms.
H(g)  H+(g) + e• All elements have a first ionization energy even atoms which don't form positive ions in
test tubes. The reason that helium (1st I.E. =
2370 kJ mol-1) doesn't normally form a positive
ion is because of the huge amount of energy
that would be needed to remove one of its
electrons.
Patterns of first Ionization Energies
in the Periodic Table:
The first 20 elements
First ionization energy shows periodicity. That means that it
varies in a repetitive way as you move through the Periodic
Table. For example, look at the pattern from Li to Ne, and then
compare it with the identical pattern from Na to Ar.
Ionization Energy Cont.
• The amount of energy required to
remove a second electron from a 1+ ion
is called the second ionization energy.
• The amount of energy required to
remove a third electron from a 2+ ion is
called the third ionization energy and so
on.
• The energy required for each
successive ionization always increases.
Ionization Energy Cont.
• 1st IE: Be + 900 kJ  Be+1 + e• 2nd IE: Be+1 + 1760 kJ  Be+2 + e• 3rd IE: Be+2 + 14850 kJ  Be+3 + e-
Electronegativity
• Measure of the ability of an atom in a
chemical compound to attract a bonding
pair of electrons. Depends on: the number
of protons in the nucleus, the distance from
the nucleus, and the amount of screening
by inner electrons.
– The Pauling scale is the most commonly
used. Fluorine (the most electronegative
element) is assigned a value of 4.0, and
values range down to cesium and
francium which are the least
electronegative at 0.7
Electronegativity
• Period trend: increases across the
period
• Group trend: decreases down the group
Electron Affinity
• Amount of energy released when an atom
gains an electron.
–
Nonmetal + e-  X- + energy
• Halogens have a very high electron affinity
(because they need only one electron to
be stable).
• The unstability of the halogens makes
them very reactive.
• When they become stable, they release
energy (exothermic).
Electron Affinity
• Period Trend: increases across a
period.
• Group Trend: decrease down a group.
The greater the distance from the
nucleus, the less the attraction.