Download Periodic Table Trends - Waterford Public Schools

Coincidence? I Think Not!
0 As you have realized, the Periodic Table provides
a great deal more information than just atomic
number and atomic mass!
0 Each period (row) corresponds to an energy
level
0 Each atomic orbital (s, p, d, f) is represented as a
“block”
0 This is because elements on the Periodic Table
are arranged by increasing atomic number
Blue = s block
0 As a result, a regular and repeating pattern of
chemical and physical properties emerges
0 Called periodic law
Red
Orange
Yellow
Green
Blue
Indigo
Violet
n=1
n=2
n=3
n=4
n=5
n=6
n=7
More on the Periodic Law
0 Atoms with similar properties appear in groups or families
(vertical columns) on the periodic table
0 They are similar because they all have the same number of
valence (outer shell) electrons, which governs their chemical
behavior
0 Remember, valence electrons are electrons in the highest-
numbered s- and p- orbitals!
0 Elements of the same period (horizontal row) have the
same number of energy levels
0 As you move across a period, the number of electrons and
protons increases, leading to increase in atomic number
0 Elements within the same period do not generally show
similarity in properties, except d-block and f-block
(lanthanides) elements
Trends in the Periodic Table
0 Due to this arrangement,
patterns of chemical and
physical properties
emerge
0 Called trends
0 These trends occur due
to the electronic
configurations of
elements!
So, What are These Trends?
Trend #1 - Atomic Radius
0 Defined as the distance between the nucleus and the
outer edge of the electron cloud
0 Since an electron cloud’s edge is difficult to define,
scientists use the covalent radius, or half the distance
between the nuclei of 2 bonded atoms
0 Atomic radii are usually measured in picometers (pm)
or angstroms (Å).
0 An angstrom is 1 x 10-10 m
Example – Atomic Radius
0 Two bromine atoms bonded together are 2.86 angstroms
apart. So, the radius of each atom is 1.43 Å.
2.86 Å
1.43 Å
1.43 Å
Atomic Radius Trend -Across a Period
0 Atomic radius DECREASES from left to right across a
period
0 Why?
0 As you move across a period, a proton AND electron are
added, as well as 1 or 2 neutrons
0 The result is a more positive nucleus and a more negative
electron cloud
0 The outer-shell electrons experience an increased
attraction to the nucleus and the electron cloud gets
pulled in, making atoms smaller
0 The amount of positive charge perceived by an electron
from the nucleus is called the nuclear effective charge
(Zeff)
0 Zeff increases across a period
Atomic Radius Trend
Down a Group
0 Atomic radius INCREASES down a group
0 Why?
0 Number of energy levels (n) increases so, orbitals are
larger
0 The bigger the distance over which the nucleus must
pull reduces the attraction for electrons
0 Furthermore, outer-shell electrons are shielded from the
full nuclear positive charge by the inner-shell electrons
Atomic Radius Animation
Increasing Atomic Radius
Decreasing Atomic Radius
Trend #2 - Ionic Radius
0 Defined as the distance from the nucleus to the other
edge of the electron cloud in a charged ion
0 Metal atoms can lose valence electrons to form positive
ions called cations
0 Nonmetal atoms can gain valence electrons to form
negative ions called anions
Ionic Radius Trend
Across a Period and Down a Group
0 Same radii trends apply once you divide the table into metal
(cation) and non-metal (anion) sections
0 Cation radii DECREASES from left to right with only minor
changes in the transition metals
0 Anion radii are larger than cations and DECREASES from left to
right
0 As electrons are added, the p+/e- ratio decreases and the electrons are
not as closely held
0 Increased electrons and electron repulsions also play a role in
expanding the electron cloud
0 Ionic radii increases down all groups because of the additional
energy levels (n)
Sodium Cation Formation
Effective nuclear charge on
remaining electrons
increases!
1 valence electron
11p
+
Remaining e- are pulled
in closer to the nucleus ionic size decreases!
Valence e- lost in
ion formation
Result - a smaller sodium
cation, Na+
Chlorine Anion Formation
A chloride ion is
produced - it is larger
than the original atom!
7 valence e-
17p
+
One e- is added to
the outer shell
Effective nuclear charge is reduced
and the e- cloud expands!
K+1
+
Cs
Br
is larger than
+
K
is bigger than Cl
When going down a group,
the ions get bigger
Cs +1
Cs+1
This is true for
cations & anions!
-
GOING ACROSS A PERIOD
CATION
S get
ANIONS get
smaller
smaller too
Ionic Radius Trend Summary
Trend #3 - Ionization Energy (IE)
0 Defined as the energy needed to remove an electron from a gaseous
atom or ion
0 In other words if an electron is given enough energy (in the form of a
photon) to overcome the effective nuclear charge in a gaseous atom or
ion, it can leave the atom completely and become ionized or charged
0 This is because the number of protons and electrons is no longer equal
0 IE is measured in kJ
0 The larger the I.E, the harder it is to remove the electron
0 This process is ALWAYS endothermic
0 This process is also a stepwise process
0 There are first, second, third, etc. ionization energies
0 First ionization energy is the energy required to remove the highest-
energy electron (one bound least tightly)
0 Second ionization energy is the energy required to remove an additional
electron from the mole of ions (+1 cations into +2 cations)
0 A large jump in ionization energy occurs when a CORE e- is
removed
Example
Ionization Energies of Magnesium
1st I.E.
736 kJ
2nd I.E.
1,445 kJ
3rd I.E.
7,730 kJ
Core e-
First Ionization Energy Trend
Across a Period
0 FIRST ionization energy INCREASES left to right across a period
0 First ionization energy and atomic radius are inversely proportional to each
other
0 Why?
0 Effective nuclear charge, Zeff, increases the attraction of the nucleus and therefore,
pulls the electron cloud closer to the nucleus
0 Holds electrons more tightly – harder to remove
Exceptions to the I.E Trend
0 Ionization energies do not increase smoothly across
the periods in the periodic table
0 In general:
- It is easier to remove
an electron if it
results in the
formation of a filled
or half-filled subshell
- It is harder to – It is harder to
remove an electron
from a filled or halffilled subshell
First Ionization Energy Trend
Down a Group
0 FIRST ionization energy DECREASES down a group
0 Why?
0 The larger the atom is, the easier its electrons are to
remove
0 Increased number of energy levels (n) increases the
distance over which the nucleus must pull and therefore,
reduces the attraction for electrons
0 Also, full energy levels provide some shielding between the
nucleus and valence electrons
Need More Clarification?
Ionization Energy Video
3-D Graph of First IE
Summary of First IE Trend
Trend #4 - Electronegativity
0 Defined as a measurement of the attraction of an atom
of the pair of outer shell electrons in a covalent bond
with another atom
0 In other words, electronegativity is a measure of an
atom’s attraction for another atom’s electrons
0 It is an arbitrary scale that ranges from 0 to 4
0 0 = no electronegativity
0 4 = high electronegativity
0 All electronegativity values are relative to fluorine, the
most electronegative element
Electronegativity Trend
Across a Period
0 Generally, metals are electron givers and have low
electronegativities
0 Nonmetals are electron takers and have high
electronegativities
0 What do you think about the electronegativity of noble
gases?
0 So, electronegativity INCREASES left to right across a period
0 Why?
0 Effective nuclear charge, Zeff, increases the attraction of the
nucleus and therefore, strengthens the attraction for electrons
0 Holds electrons more tightly
Electronegativity Trend Down a
Group
0 Electronegavity DECREASES down a group
0 Why?
0 Increased number of energy levels (n) increases the
distance over which the nucleus must pull and
therefore, reduces the attraction for electrons
0 Also, full energy levels provide some shielding between
the nucleus and valence electrons
Summary of Electronegativity Trend
0
A General Justification of the Trends
0 Across a period, electrons of the outermost shell experience
increased nuclear attraction due to increase in atomic number
0 Called nuclear effective charge (Zeff) – the amount of positive
charge perceived by an electron from the nucleus
0 Down a group, nuclear attraction decreases due to increased
distance from the nucleus and the “shielding effect”
0 Electrons intervening between the nucleus and an outer electron
are said to “shield” or screen the outer electron from the nucleus
0 As a result, the outer electron does not experience the full nuclear
charge
0 Only full energy levels, not full orbitals (s, p, d, f), are of concern
in a shielding argument
Summary of All Periodic
Trends