Download Chapter 5 Section 3: Electron Configuration and Periodic Properties

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Chapter 5 Section 3: Electron Configuration and Periodic Properties
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Valence Electrons – The electrons available to be lost,
gained, or shared in the formation of chemical
compounds. Chemical compounds form because
electrons are lost, gained, or shared between atoms.
Main-Group Elements
Main-group elements are the s-block and p-block elements
(also known and the representative elements).
Atomic Size (Atomic radius): is one-half the distance
between the nuclei of identical atoms that are bonded
together.
Period Trend: atoms get smaller across a period This is
caused by the increasing positive charge of the nucleus.
Group Trend: In general, the atomic radii (size) of the MainGroup Elements increase down a group. The cause is that
electrons occupy orbitals in higher energy levels
Ion – an atom or group of bonded atoms that has a
positive or negative charge.
+
Example: Sodium forms a Na ion. (cation)
Ionization – Any process that result in the formation of an
ion
Ionization energy
Ionization energy, IE (or first ionization energy, IE1) – is
energy required to remove one electron from a neutral
atom of an element. Ionization energy is measured in
kilojoules per mole (kJ/mol)
Period Trend: In general, ionization energies of the maingroup elements increase across each period. This
increase is caused by increasing nuclear charge.
Group Trend: among the main-group elements, ionization energies
generally decrease down the groups. Electrons removed
from atoms of each succeeding element in a group are in
higher energy levels, farther from the nucleus. Plus there
are electrons between the valence electrons and the
nucleus that shield the outer electrons. The shielding
makes valence electrons easier to remove.
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2nd IE (IE2), or 3rd (IE3) Each successive electron removed from
an ion has an increasingly stronger effective nuclear charge; therefore,
they require more energy to remove.
Electron Affinity – The energy change that occurs when an electron is
added to a neutral atom; Energy released is represented by a negative
number. Some atoms must be “forced” to gain an electron by the addition of
energy. When energy is absorbed it is represented by a positive number and
the ion is unstable and usually decays (breaks apart back to the element).
Periodic Trend for Electron Affinity – It is easier to add
electrons to the p block. An exception occurs between
Groups 14 and 15. The ease with which halogen atoms
gain electrons is a major reason for the high reactivities
of the Group 17 elements.
Group Trends for Electron Affinity – Trends for electron affinities
within groups can be unpredictable;
as a general rule,
electrons are more difficult to add down a group. This
pattern is a result of two competing factors: nuclear
charge and size. In general, the size effect predominates
and lowers the electron affinity.
Ions
A cation is a positive ion. An anion is a negative ion.
Ionic Radii
A cation forms by the loss of one or more electrons;
This leads to a decrease in atomic radius because the
removal of the highest-energy-level electrons causes a
smaller electron cloud.
An anion forms by the addition of one or more
electrons; which always causes an increase in electron
cloud size.
Period Trends
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.
Group Trends
Just as there is a gradual increase of atomic radii down
a group, there is also a gradual increase of ionic radii.
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Sublevel Shielding – the
inner electrons (d and f orbitals)
shield the valence electrons causing the nucleus to have
less pull on the valence electrons (Weakens the atoms
hold on the valence electrons.)
Electronegativity
Electronegativity – measure of the ability of an atom in a chemical compound
to attract electrons.
Linus Pauling devised a scale of numerical values
reflecting the tendency of an atom to attract electrons.
The most electronegative element, fluorine, is arbitrarily assigned
an electronegativity value of four. The rest of the
element’s values are then calculated relative to fluorine
Periodic Trend: Electronegativities increase across each
period, although there are exceptions.
Group Trend: Electronegativities either decrease down a
group or remain about the same.
The combination of the period and group trends in
electronegativity results in the highest values belonging
to the elements in the upper right of the periodic table.
The lowest values belong to the elements in the lower
left of the table. Noble Gases do not have assigned
values – they don’t react readily.
Periodic Properties of the d- and f-Block Elements
The properties of the d-block elements vary less and with less regularity
than those of the main-group elements.
The d- and f-Block
Elements
Atomic Radii - The atomic radii of the d-block elements generally decrease
across the periods, this decrease is less than that for the main-group elements
because the electrons added to the (n − 1)d sublevel shield the valence
electrons
Atomic Radii
Ionization Energy - ionization energies of the d-block and f-block elements
generally increase across the periods. The first ionization energies of the dblock elements generally increase down each group.
Ionization Energy
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Ion Formation and Ionic Radii - The order in which electrons are Ion Formation
removed from all atoms of the d-block and f-block elements is exactly the
and Ionic Radii
reverse of the order given by the electron-configuration notation. In other
words, electrons in the highest occupied sublevel are
always removed first. For the d-block elements, this
means that although newly added electrons that occupy
the d sublevels are not removed first; the first electrons
to be removed are those in the outermost s sublevels.
For example, iron, Fe, has the electron configuration
[Ar]3d64s2. It loses a 4s electron first to form Fe+
([Ar] 3d 64s 1). Fe+ can then lose the second 4s electron
to form Fe2+ ([Ar] 3d 6). Fe2+ can then lose a 3d electron
to form Fe3+ ([Ar] 3d 5).
Most d-block elements commonly form 2+ ions in
compounds. Some, such as iron and chromium, also
commonly form 3+ ions. The Group 3 elements form only
ions with a 3+ charge. Copper forms 1+ and 2+ ions, and
silver usually forms only 1+ ions.
As expected, the cations have smaller radii than the atoms
do. Comparing 2+ ions across the periods follows the
decrease in size that parallels the decrease in atomic radii.
Electronegativity - The d-block elements all have
electronegativities between 1.1 and 2.54. Only the active
metals of Groups 1 and 2 have lower electronegativities.
The d-block elements also follow the general trend for
electronegativity values to increase as radii decrease, and
vice versa.
The f-block elements all have similar electronegativities,
which range from 1.1 to 1.5
Electronegativity