ch. 3: periodic properties of the...
TRANSCRIPT
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Ch. 3: Periodic Properties of
the Elements
Dr. Namphol Sinkaset
Chem 200: General Chemistry I
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I. Chapter Outline
I. Introduction
II. The Periodic Table
III. Electrons in the Atom
IV. Electron Spin
V. Sublevel Energy Splitting
VI. Using the Periodic Table
VII. Periodic Properties and Trends
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I. Organizing Chemical Info
• When information of the elements was
organized, chemistry began to advance
quickly.
• Element “triads” and “octaves”
• Mendeleev’s periodic table in 1869
• Quantum mechanics explains why the
periodic table appears as it does.
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II. Periodic Law
• Initially, Mendeleev ordered elements
by increasing atomic mass.
• Later work by Moseley showed that they
should be ordered by atomic number.
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II. The Modern Periodic Table
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II. Major Divisions of the Table
• Main-group elements have properties
that are largely predictable based on
their location.
• Transition and inner-transition elements
have properties that are less predictable
based on their location.
• Each column within the main group
region is known as a family or group.
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III. Electrons Occupying Orbitals
• From Chapter 3, we know how orbitals
are ordered for the hydrogen atom
• Since hydrogen has only one e-, the
ground state can be written as an
electron configuration:
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III. Many e- Atoms
• The Schrödinger equation can’t solve
multi-e- atoms; we only get approximate
solutions.
• We use quantum #’s from H atom
solution to describe orbitals of other
atoms.
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III. New Considerations
• An atom with more than 1 e- is more
complicated.
• Two more concepts are needed to
understand these larger atoms:
1) Electron spin
2) Sublevel energy splitting
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IV. H Atoms in a Magnetic Field
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IV. e- Spin
• e- generate a small magnetic field as if they were spinning.
• There are two possible directions e- can spin, so there are two possible states.
• spin quantum number (ms) can be either +1/2 or –1/2.
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IV. Representing e- Spin
• Orbital diagrams are used to show
electron occupation and spin.
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IV. Pauli Exclusion Principle
• No two e- in the same atom can have the same 4 quantum #’s!!
• H: n=1, l=0, ml=0, ms=1/2
• He has two p+, so it needs two e-:
1st e-: n=1, l=0, ml=0, ms=1/2
2nd e-: n=1, l=0, ml=0, ms=-1/2
• The orbital is filled and the e- have paired spins.
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IV. Electrons in Helium
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V. H vs. He Energy Levels
• One additional e- complicates the He
spectrum greater than expected. Why?
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V. Removal of Degeneracy
• In H atom, energy of an orbital depends
only on n.
e.g. Energies of 3s, 3p, 3d are degenerate.
• In every other atom, this is not true.
E (s orbital) < E (p orbital) < E (d orbital) <
E (f orbital), etc.
• What removes the degeneracy?
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V. Sublevel Energy Splitting
• Three factors contribute to differing
sublevel energies:
1) Coulomb’s Law (Z)
2) shielding
3) penetration
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V. Coulomb’s Law
• The PE of like charges is positive (unstable),
but decreases as they move apart.
• The PE of unlike charges is negative (stable)
and increases as they get closer.
• The magnitude of the interaction increases as
charges on particles increases.
r
qqE 21
04
1
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V. Nuclear Charge
• p+ in nucleus
constantly pull all e-.
• Higher charges
attract more strongly.
• More p+ lowers
orbital E by
increasing e-/nucleus
attraction.
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V. Shielding
• Electrons shield each
other from the full
charge of the nucleus.
• The effective nuclear
charge, Zeff, is the
actual positive charge
an e- feels.
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V. Penetration
• The movement of an
outer e- into the
region occupied by
inner e- is called
penetration.
• Penetrating e-
experience higher
nuclear charge,
lowering its PE.
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V. 2s and 2p Radial Distribution
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V. 3s, 3p, 3d Penetration
• This is the reason why energetically, s < p < d.
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V. Order of Sublevels
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V. The Aufbau Principle
• Since e- are “lazy,” they want to
“occupy” the lowest energy level
possible.
• Thus, if we know the energy order of
sublevels, then we can “build up” the e-
configurations of each atom.
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V. Writing e- “in” Orbitals
• Two ways to represent how e- are
situated in atoms:
1) e- configuration, nl#
2) orbital diagram, which uses arrows
indicating e-’s and their spin
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V. Hund’s Rule
• In the orbital diagram of C, there was a
choice as to where to place the 2nd p
orbital.
• We follow Hund’s rule.
When filling degenerate orbitals, electrons
fill singly first with parallel spins.
• Hund’s rule leads to lower energy.
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V. Examples
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VI. The Periodic Table
• As you go left to right on the periodic table,
you are using the Aufbau principle.
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VI. The Periodic Table
• Each region of the periodic table indicates
what orbitals are being “filled.”
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VI. Using the Periodic Table
• You can use an element’s location to
write its full or condensed electron
configuration/orbital diagram.
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VI. Using the Periodic Table
• Therefore, Cl is: [Ne] 3s2 3p5.
• From the orbital diagram, we can write
specific quantum numbers for each e-.
• Which e-’s are identified with the
following quantum #’s {n, l, ml, ms}?
{3, 0, 0, -1/2}
{3, 1, 1, 1/2}
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VI. Some Caveats
• Because energy differences between s
and d are small, some exceptions to
how e-’s fill exist.
Same for d and f.
• Remember that d principal quantum #
lags by one.
• Remember that f principal quantum #
lags by two.
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VI. Sample Problem 3.1
• Write condensed electron configurations
and orbital diagrams for the following
elements.
Mn
Sb
Nd
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VI. The Periodic Table
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VI. Important Parts of the Periodic Table
1) Each element placed in box w/ atomic #,
atomic mass, and atomic symbol.
2) Atomic # increases as go L to R.
3) Each horizontal row is period.
4) Each vertical column is a group or family.
5) Main group elements are in groups 1,2
and 13-18 (s and p blocks).
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VI. Important Parts of the Periodic Table
6) Transition elements are in groups 3-12 (d
block).
7) Inner-transition elements at the bottom
(lanthanides and actinides, f block).
8) Staircase line separates metals on L
from nonmetals on R. Metalloids or
semimetals lie adjacent to the line.
9) Some groups have special names: alkali
metals, alkali earth metals, halogens,
noble gases.
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VI. Types of Elements
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VI. Core vs. Valence e-’s
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VI. Valence Electrons
• valence electrons: the outermost e- in an atom
• Valence e- determine an atom’s chemistry; thus, atoms in the same vertical column have similar chemical properties.
• Valence e- can be determined from the Group number.
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VI. Formation of Ions
• Metals tend to lose e-’s and nonmetals
tend to gain e-’s.
• Main-group ions can be predicted.
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VI. Transition Metal Cations
• When forming transition metal cations,
remove e-’s from highest n-value orbital
first!
V: [Ar] 4s2 3d3
V2+: [Ar] 4s0 3d3
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VI. Magnetic Properties
• Some metals exhibit magnetism
paramagnetic: atom or ion that has
unpaired e-’s
diamagnetic: atom or ion in which all e-’s
are paired
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VI. Sample Problem 3.2
• Draw condensed orbital diagrams for
the following and determine whether
they are diamagnetic or paramagnetic.
Sc3+
Ir2+
Mn4+
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VII. Atomic Radii
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VII. Trend in Atomic Radii
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Trend in Atomic Radii
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VII. Trend in Ion Size
• Why?
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VII. Trend in Ionization Energy
• ionization energy:
energy in kJ needed to
remove an e- from
gaseous atoms/ions
• Why?
• What about 1st, 2nd, 3rd,
ionization energies?
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VII. Successive IE’s
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VII. Electron Affinity
• electron affinity:
energy change in kJ
when e- added to a
gaseous atom/ion
(generally negative)
• Why?
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VII. Trend in Metallic Character