Transcript Binnie periodic_trends

Periodic Trends
Elemental Properties and Patterns
The Periodic Law
• Dimitri Mendeleev was the first scientist to
publish an organized periodic table of the
known elements.
• He was perpetually in trouble with the
Russian government and the Russian
Orthodox Church, but he was brilliant
never-the-less.
The Periodic Law
• Mendeleev even went out on a limb and
predicted the properties of 2 at the time
undiscovered elements.
• He was very accurate in his predictions,
which led the world to accept his ideas
about periodicity and a logical periodic
table.
The Periodic Law
• Mendeleev understood the ‘Periodic Law’
which states:
• When arranged by increasing atomic
number, the chemical elements display a
regular and repeating pattern of chemical
and physical properties.
The Periodic Law
• Atoms with similar properties appear in
groups or families (vertical columns) on the
periodic table.
• They are similar because they all have the
same number of valence (outer shell)
electrons, which governs their chemical
behavior.
A Different Type of Grouping
• Besides the 4 blocks of the table, there is
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another way of classifying element:
Metals
Nonmetals
Metalloids or Semi-metals.
The following slide shows where each
group is found.
Metals, Nonmetals, Metalloids
Metals, Nonmetals, Metalloids
• There is a zig-zag or
staircase line that
divides the table.
• Metals are on the left
of the line, in blue.
• Nonmetals are on the
right of the line, in
orange.
Metals, Nonmetals, Metalloids
• Elements that border
the stair case, shown
in purple are the
metalloids or semimetals.
• There is one important
exception.
• Aluminum is more
metallic than not.
Metallic Character
• This is simple a relative measure of how
easily atoms lose or give up electrons.
Metals, Nonmetals, Metalloids
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How can you identify a metal?
What are its properties?
What about the less common nonmetals?
What are their properties?
And what the heck is a metalloid?
Metals
• Metals are lustrous
(shiny), malleable,
ductile, and are good
conductors of heat and
electricity.
• They are mostly solids
at room temp.
• What is one
exception?
Nonmetals
• Nonmetals are the
opposite.
• They are dull, brittle,
nonconductors
(insulators).
• Some are solid, but
many are gases, and
Bromine is a liquid.
Metalloids
• Metalloids, aka semi-metals
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are just that.
They have characteristics of
both metals and nonmetals.
They are shiny but brittle.
And they are
semiconductors.
What is our most important
semiconductor?
Periodic Trends
• There are several important atomic
characteristics that show predictable trends
that you should know.
• The first and most important is atomic
radius.
• Radius is the distance from the center of the
nucleus to the “edge” of the electron cloud.
Atomic Radius
• Since a cloud’s edge is difficult to define,
scientists use define covalent radius, or half
the distance between the nuclei of 2 bonded
atoms.
• Atomic radii are usually measured in
picometers (pm) or angstroms (Å). An
angstrom is 1 x 10-10 m.
Covalent Radius
• Two Br 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
• The trend for atomic radius in a vertical
column is to go from smaller at the top to
larger at the bottom of the family.
• Why?
• With each step down the family, we add an
entirely new PEL to the electron cloud,
making the atoms larger with each step.
Atomic Radius
• The trend across a horizontal period is less
obvious.
• What happens to atomic structure as we step
from left to right?
• Each step adds a proton and an electron
(and 1 or 2 neutrons).
• Electrons are added to existing PELs or
sublevels.
Atomic Radius
• The effect is that the more positive nucleus
has a greater pull on the electron cloud.
• The nucleus is more positive and the
electron cloud is more negative.
• The increased attraction pulls the cloud
in, making atoms smaller as we move from
left to right across a period.
Effective Nuclear Charge
• What keeps electrons from simply flying off
into space?
• Effective nuclear charge is the pull that an
electron “feels” from the nucleus.
• The closer an electron is to the nucleus, the
more pull it feels.
• As effective nuclear charge increases, the
electron cloud is pulled in tighter.
• The
overall
trend in
atomic
radius
looks like
this.
Atomic Radius
Cation Formation
Effective nuclear
charge on remaining
electrons increases.
Na atom
1 valence electron
11p+
Valence elost in ion
formation
Result: a smaller
sodium cation, Na+
Remaining e- are
pulled in closer to
the nucleus. Ionic
size decreases.
Anion Formation
Chlorine
atom with 7
valence e17p+
One e- is added
to the outer
shell.
Effective nuclear charge is
reduced and the e- cloud
expands.
A chloride ion is
produced. It is
larger than the
original atom.
Sizes of Ions
• Cations are smaller
than their parent
atoms.
• The outermost
electron is removed
and repulsions
between electrons
are reduced.
© 2009, Prentice-Hall, Inc.
Sizes of Ions
• Anions are larger
than their parent
atoms.
• Electrons are added
and repulsions
between electrons
are increased.
© 2009, Prentice-Hall, Inc.
Sizes of Ions
• In an isoelectronic series, ions have the same
number of electrons.
• Ionic size decreases with an increasing nuclear
charge.
© 2009, Prentice-Hall, Inc.
Shielding
• As more PELs are added to atoms, the inner
layers of electrons shield the outer electrons
from the nucleus.
• The effective nuclear charge (enc) on those
outer electrons is less, and so the outer
electrons are less tightly held.
Ionization Energy
• This is the second important periodic trend.
• If an electron is given enough energy (in the
form of a photon) to overcome the effective
nuclear charge holding the electron in the
cloud, it can leave the atom completely.
• The atom has been “ionized” or charged.
• The number of protons and electrons is no
longer equal.
Ionization Energy
• The energy required to remove an electron
from a gaseous atom is ionization energy.
(measured in kilojoules, kJ)
• The larger the atom is, the easier its electrons
are to remove.
• Ionization energy and atomic radius are
inversely proportional.
• Ionization energy is always endothermic, that
is energy is added to the atom to remove the
electron.
Ionization Energy
• Energy required to remove an electron from a gaseous atom or
ion.
• X(g) → X+(g) + e–
Mg → Mg+ + e–
Mg+ → Mg2+ + e–
Mg2+ → Mg3+ + e–
I1 = 735 kJ/mol(1st IE)
I2 = 1445 kJ/mol
(2nd IE)
I3 = 7730 kJ/mol
*(3rd IE)
*Core electrons are bound much more tightly than valence
electrons.
Ionization Energy
• In general, as we go down a group from top to
bottom, the first ionization energy decreases.
• Why?
• The electrons being removed are, on average,
farther from the nucleus.
Ionization Energy
• In general, as we go across a period from left to
right, the first ionization energy increases.
• Why?
• Electrons added in the same principal quantum
level do not completely shield the increasing
nuclear charge caused by the added protons.
• Electrons in the same principal quantum level
are generally more strongly bound from left to
right on the periodic table.
Ionization Energy
The Values of First Ionization Energy for the Elements in the
First Six Periods
Trends in First Ionization Energies
However, there
are two apparent
discontinuities
in this trend.
© 2009, Prentice-Hall, Inc.
Trends in First Ionization Energies
• The first occurs between
Groups 2 & 3
• In this case the electron
is removed from a porbital rather than an sorbital.
• The electron removed is
farther from nucleus.
• There is also a small
amount of repulsion by
the s electrons.
© 2009, Prentice-Hall, Inc.
Trends in First Ionization Energies
• The second occurs
between Groups 15 and
16
• The electron removed
comes from doubly
occupied orbital.
• Repulsion from the other
electron in the orbital aids
in its removal.
© 2009, Prentice-Hall, Inc.
CONCEPT CHECK!
Which atom would require more energy to
remove an electron? Why?
Na
Cl
CONCEPT CHECK!
Which atom would require more energy to
remove an electron? Why?
Li
Cs
CONCEPT CHECK!
Which has the larger second ionization
energy? Why?
Lithium or Beryllium
Successive Ionization Energies (KJ per Mole) for the Elements
in Period 3
PES Photoelctron Spectroscopy
• Method that provides information on all the
occupied energy levels of an atom
• We can compare the energy required by a
photon to knock an electron out of an atom
(ionization energy), to the signal output to
deduce how many electrons are at each
energy level.
Presented By, Mark Langella, APSI
Chemistry 2013 , PWISTA.com
Methodology
• Very high energy photons, such as very-shortwavelength ultraviolet radiation, or even x-rays, are
used in this experiment.
• The gas phase atoms are irradiated with photons of a
particular energy.
• If the energy of the photon is greater than the energy
necessary to remove an electron from the atom, an
electron is ejected with the excess energy appearing as
kinetic energy, ½ mv2, where v is the velocity of the
ejected electron.
• The speed of the ejected electron depends on how much
excess energy it has received.
Presented By, Mark Langella, APSI
Chemistry 2013 , PWISTA.com
Ionization Energy
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IE = the ionization energy of the electron
KE = the kinetic energy with which it leaves the atom
Ephoton = IE + KE or
IE = Ephoton – KE
kinetic energy of the electrons is measured in a
photoelectron spectrometer
• photons of sufficient energy are used, an electron may
be ejected from any of the energy levels of an atom
• every electron in each atom has an (approximately)
equal chance of being ejected
Presented By, Mark Langella, APSI
Chemistry 2013 , PWISTA.com
Spectrum
• for a large group of identical atoms, the
electrons ejected will come from all possible
energy levels of the atom
• because the photons used all have the same
energy, electrons ejected from a given energy
level will all have the same energy.
• Only a few different energies of ejected
electrons will be obtained, corresponding to
the number of energy levels in the atom.
Presented By, Mark Langella, APSI
Chemistry 2013 , PWISTA.com
Photoelectron spectrum
• The photoelectron spectrum is a plot of the number of
ejected electrons (along the vertical axis) vs. the
corresponding ionization energy for the ejected
electrons (along the horizontal axis)
• The “X” axis is labeled high to low energies so that you
think about the XY intersect as being the nucleus.
http://www.chem.arizona.edu/chemt/Flash/photoelectron.
html
Presented By, Mark Langella, APSI
Chemistry 2013 , PWISTA.com
Presented By, Mark Langella, APSI
Chemistry 2013 , PWISTA.com
Interpretations from the data:
1. There are no values on the y axis in the tables above. Using the
Periodic Table and Table 1, put numbers on the y axis.
2. Label each peak on the graphs above with s, p, d, or f to indicate
the suborbital they represent..
3. What is the total number of electrons in a neutral potassium atom?
PES Question
• If a certain element being studied by a PES
displays an emission spectrum with 5
distinct kinetic energies. What are all the
possible elements that could produce this
spectrum?
Presented By, Mark Langella, APSI
Chemistry 2013 , PWISTA.com
Electron Affinity
• What does the word ‘affinity’ mean?
• Electron affinity is the energy change that
occurs when an atom gains an electron
(also measured in kJ).
• Where ionization energy is always
endothermic, electron affinity is usually
exothermic, but not always.
Electron Affinity
• Electron affinity is exothermic if there is an
empty or partially empty orbital for an
electron to occupy.
• If there are no empty spaces, a new orbital
or PEL must be created, making the process
endothermic.
• This is true for the alkaline earth metals and
the noble gases.
Electron Affinity
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Electronegativity
• Electronegativity is a measure of an atom’s
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attraction for another atom’s electrons.
It is an arbitrary scale that ranges from 0 to 4.
The units of electronegativity are Paulings.
Generally, metals are electron givers and have
low electronegativities.
Nonmetals are are electron takers and have
high electronegativities.
What about the noble gases?
Electronegativity
0
Overall Reactivity
• This ties all the previous trends together in
one package.
• However, we must treat metals and
nonmetals separately.
• The most reactive metals are the largest
since they are the best electron givers.
• The most reactive nonmetals are the
smallest ones, the best electron takers.
Overall Reactivity
0