Download Periodic Law

The History of
the Modern
Periodic Table
See separate slide show for Periodic Table History
Periodic Law
• When elements are arranged in order of
increasing atomic #, elements with similar
properties appear at regular intervals.
Atomic Radius (pm)
250
200
150
100
50
0
0
5
10
Atomic Number
15
20
Chemical Reactivity
Families  Similar valence e- within a group
result in similar chemical properties
1
2
3
4
5
6
7
•Alkali Metals
•Alkaline Earth Metals
•Transition Metals
•Halogens
•Noble Gases
Periodic Table Reveals Periodic Trends
• Effective Nuclear charge
• Reactivity
• atomic size or radius
• bonding characteristics
• ionization energy
• crystal configurations
• electron affinity
• acidic properties
• electronegativity
• densities
• metallic character
• Melting/Boiling points
Electron screening or shielding
• Electrons are attracted to the nucleus
• Electrons are repulsed by other electrons
• Electrons would be bound more tightly if
other electrons weren’t present.
• The net nuclear charge felt by an electron is
called the effective nuclear charge ( Zeff ).
Quantum Mechanical Model
Zeff is lower than actual
nuclear charge.
Zeff increases
toward nucleus
ns > np > nd > nf
This explains certain periodic
changes observed.
Effective Nuclear Charge ( Zeff)
• The effective nuclear charge acting on an
electron equals the number of protons in
the nucleus, Z, minus the average number
of electrons, S that are between the nucleus
and the electron in question.
Zeff = Z  S
Zeff = attractive forces  repulsive forces
Zeff = # protons  # shielding electrons
For Example, Lithium vs. Carbon
Li Zeff = 3  2 = 1
C
Zeff = 6  2 = 4
When moving across a row:
The greater the Zeff value,
the smaller the atom’s radius.
So, carbon has a much smaller atomic radius compared to lithium: Rcarbon =77
pm Rlithium = 152 pm
Trend #1 Atomic Radii
Increases to Left and Down
1
2
3
4
5
6
7
•Why larger going down?
•Higher energy levels have larger orbitals
•Shielding - core e- block the attraction between the nucleus and the valence e•Why smaller to the right?
• Increased nuclear charge without additional shielding pulls e- in tighter
Practice…
• Referring to a periodic table, arrange the
following atoms in order of increasing size:
– Phosphorus
– Sulfur
– Arsenic
– Selenium
• S < P < Se < As
Atomic radii
The Periodic Table & Radii
Periodic Trend is Due to
Effective Nuclear Charge
Atomic Radii vs. Zeff:
Trends in Ionic Radii
• Using your knowledge of Zeff, how would
the size of a cation compare to neutral
atom? Anion?
Trends in Ionic Radii
• The cation of an atom decreases in size.
• The more positive an ion is, the smaller it is because
Zeff increases
• The anion of an atom increases in size.
• The more negative an ion, the larger it is because
Zeff decreases.
Cations  lose electrons, become smaller
Anions  gain electrons, become bigger
Ion Radii
Increases down
1
2
3
4
5
6
7
Increases moving across, but depends
if cation OR anion
+3 +4 -3 -2 -1
Ions and Ionic Radii
Practice…
• Arrange the following atoms and ions in order
of decreasing size:
– Mg2+
– Ca2+
– Ca
• Which of the following ions is the largest:
– S2–S
– O2-
Practice…
• Arrange the following ions in order of decreasing
size:
– S2– Cl– K+
– Ca2+
• Which of the following ions is the largest?
– Rb+
– Sr2+
– Y3+
Trend in Ionization Energy
• Ionization NRG is the NRG required to
remove an electron from an atom
Successive Ionization NRG
• Ionization energy increases for successive
electrons from the same atom.
Why do you think there is such a big jump for Mg3+?
*Notice the large jump in ionization energy
when a core e is removed.
• The smaller the atom, the higher the
ionization energy due to Zeff
• Bigger atoms have lower ionization NRG
due to the fact that the electrons are
further away from the nucleus and
therefore easier to remove.
Decreases
Increases
Practice…
• Which of the following elements would
have the highest second ionization
energy? Justify your answer.
–Sodium, Sulfur, or Calcium
• Which will have the greater third
ionization energy, Ca or S? Justify your
answer.
Practice…
• Referring to a periodic table, arrange the
following atoms in order of increasing first
ionization energy (Ne, Na, P, Ar, K) Justify your
answer.
• Based on the trends discussed in this section,
predict which of the following atoms (B, Al, C
or Si) has the lowest first ionization energy
and which has the highest first ionization
energy.
Electron Affinity
• The energy change associated with the addition of
an electron
• Tends to increase across a period
• Tends to decrease as you go down a group
• Abbreviation is Eea, it has units of kJ/mol. Values are
generally negative because energy is released.
• Value of Eea results from interplay of nucleus
electron attraction, and electron–electron
repulsion.
Ionization NRG vs. Electron Affinity
• Ionization energy measures the ease with
which an atom loses an electron
• Electron affinity measures the ease with
which an atom gains an electron
Electron Affinity
Trends in Electronegativity
• tendency for an atom to attract
electrons when it is chemically combined
with another atom.
• decreases as you move down a group
• increases as you go across a period from
left to right.
Trend #5 Metallic Character
• The metallic character of atoms can be related
to the desire to lose electrons.
• The lower an atom’s ionizatoin energy, the
greater its metallic character will be.
• On the periodic table, the metallic character of
the atoms increase down a family and decreases
from left to right across a period.
Metals
Nonmetals
• Shiny Luster
• Various colors (most
silvery)
• Solids are malleable and
ductile
• Good conductors of heat
and electricity
• Most metal oxides are
ionic solids that are basic
• Tend to form cations in
aqueous solution
•
•
•
•
No luster
Various colors
Brittle solids
Poor conductors of heat
and electricity
• Most nonmetal oxides
are molecular
substances that form
acidic solutions
• Tend to form anions or
oxyanions in aqueous
solution
Metallic Character
Increases moving down and across to the left
1
2
3
4
5
6
7
Rb
Cs Ba
Fr Ra
Lower left corner -- elements most
likely to lose their valence electrons
Metals and Nonmetals
• Low ionization energies of metals means they
tend to form cations (positive ions) relatively
easily
• Due to their electron affinities, nonmetals
tend to gain electrons when they react with
metals.
# 6 Melting/Boiling Points
• Highest in the middle of a period (generally).
1
2
3
4
5
6
7
Some Important Properties of Alkali Metals
• Soft metallic solids
• Easily lose valence electrons (Reducing
Agents)
– React with halogens to form salts
– React violently with water
• Large Hydration NRG
– Positive ionic charge makes ions attractive to
polar water molecules
Alkaline Earth Metals…
• Harder and more dense than Alkali Metals
• Less reactive than alkali metals (lower first
ionization energies)
• Reactivity increases as you move down the
periodic table.
The Halogens…
• “Salt Formers”
• Melting and Boiling Points increase with
atomic number.
• Highly negative electron affinities
• Tendency to gain electrons and form halide
ions
Noble Gases …
•
•
•
•
Monoatomic ions
Gases at room temperature
Large 1st ionization energies
“Exceptionally” unreactive
Practice…
• Look at Sample Integrative Exercise 7 on page
264