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The Periodic Table
Chemistry I Honors
What Element is...?
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Half of a Dime Nickel
The Lone Ranger’s horse Silver
Not fat
Tin
Watered down gin Hydrogen
A police officer
Copper
What I do when I’m hungry Iodine
What torpedoed ships do Zinc
Male member of the Ganese tribe Manganese
What he did with a bucking horse Rhodium
What should be done with an ailing man Helium/Curium/Barium
History of the Periodic Table
Mendeleev and Chemical Periodicity
• Mendeleev noticed that when elements were arranged in order of increasing
atomic mass, certain similarities in their chemical properties appeared at
regular intervals.
• Repeating patterns are referred to as periodic.
• Mendeleev created a table in which elements with similar properties were
grouped together - a periodic table of elements.
History of the Periodic Table
Mendeleev and Chemical Periodicity
• After Mendeleev placed all the known elements in his periodic table, several
empty spaces were left.
• In 1871, Mendeleev predicted the existence and properties of elements that
would fill three of the spaces.
• By 1886, all three elements had been discovered.
History of the Periodic Table
Moseley and the Periodic Law
• In 1911, the English scientist Henry Moseley discovered that the elements fit
into patterns better when they were arranged according to atomic number,
rather than atomic weight.
• The Periodic Law states that the physical and chemical properties of the
elements are periodic functions of their atomic number.
History of the Periodic Table
The Modern Periodic Table
• The Periodic Table is an arrangement of the elements in order of their atomic
number so that elements with similar properties fall in the same column, or
group.
History of the Periodic Table
Periods and Blocks of the Periodic Table
• Elements are arranged vertically in the periodic table in groups that share
similar chemical properties.
• Elements are also organized horizontally in rows, or periods.
• The periodic table is divided into four blocks, the s, p, d, and f blocks. The
name of each block is determined by the electron sublevel being filled in
that block.
History of the Periodic Table
Periods and Blocks of the Periodic Table
History of the Periodic Table
Periods and Blocks of the Periodic Table
• The elements of Group 1 of the periodic table are known as the alkali
metals.
• lithium, sodium potassium, rubidium, cesium, and francium
• In their pure state, all of the alkali metals
• have a silvery appearance
• soft enough to cut with a knife
• cannot be found in nature as free elements
• most reactive metals
• 1 valence electron
History of the Periodic Table
Periods and Blocks of the Periodic Table
History of the Periodic Table
Periods and Blocks of the Periodic Table
• The elements of Group 2 of the periodic table are known as the alkaline
earth metals.
• beryllium, magnesium, calcium, strontium, barium, radium
• Group 2 metals are
• less reactive than alkali metals
• too reactive to be found in pure form
• 2 valence electrons
History of the Periodic Table
Periods and Blocks of the Periodic Table
• The s-block elements consist of all the elements of Groups 1 and 2 and
helium.
• The p-block elements consist of all the elements in Groups 13-18 except
helium.
• Together, the s-block and p-block elements are known as the main group
elements.
History of the Periodic Table
Periods and Blocks of the Periodic Table
• The properties of the elements of the p-block vary greatly.
• At its right-hand end, the p-block includes all of the nonmetals expect
hydrogen and helium.
• All six of the metalloids are also in the p-block.
• At the left-hand end and bottom of the block, there are eight p-block metals.
History of the Periodic Table
Periods and Blocks of the Periodic Table
History of the Periodic Table
Periods and Blocks of the Periodic Table
• The elements of Group 17 of the periodic table are known as the halogens.
• fluorine, chlorine, bromine, iodine, astatine
• Group 17 elements are
• most reactive nonmetals
• 7 valence electrons
• elements are diatomic (except astatine)
• elements have distinctive color
• react vigorously with most metals to form examples of the type of
compounds known as salts.
History of the Periodic Table
Periods and Blocks of the Periodic Table
History of the Periodic Table
Periods and Blocks of the Periodic Table
• The metalloids, or semiconducting elements, are located between the
nonmetals and metals in the p-block.
• boron, silicon, germanium, arsenic, antimony, tellurium
• The metals of the p-block are generally harder and denser than the s-block
alkaline earth metals, but softer and less dense than the d-block.
History of the Periodic Table
Periods and Blocks of the Periodic Table
History of the Periodic Table
Periods and Blocks of the Periodic Table
• The elements of Group 18 are known as the nobel gases.
• helium, neon, argon, krypton, xenon, radon
• Group 18 elements are
• largely unreactive
• full valence shell (2 for helium, 8 for all other elements)
History of the Periodic Table
Periods and Blocks of the Periodic Table
History of the Periodic Table
Periods and Blocks of the Periodic Table
• The elements in Groups 3-12 are known as the transition elements.
• Properties of Transition Elements
• good conductors
• multiple oxidation states
• occurs as uncombined elements in nature
• less reactive than Group 1 & 2
History of the Periodic Table
Periods and Blocks of the Periodic Table
History of the Periodic Table
Periods and Blocks of the Periodic Table
• In the periodic table, the f-block elements are wedged between Groups 3 &
4 in the sixth and seventh periods.
• Their position reflects the fact that they are involved in filling the 4f and 5f
sublevels.
• The first row of the f-block, the lanthanides, are shiny metals similar in
reactivity to the Group 2 metals and they are magnetic.
• The second row of the f-block, the actinides, are between actinium and
rutherfordium. The actinides are all radioactive.
Periodic Trends
Ionization Energy
• Ionization energy is the energy required to remove an electron from an atom.
• increases for successive electrons
• increase across a period
• Electrons in the same quantum level do not shield as effectively as
electrons in inner levels
• Irregularities at half filled sublevels due to extra repulsion of electrons
paired in orbitals, making them easier to remove
• decreases with increasing atomic number within a group
• Electrons further from the nucleus are easier to remove due to a decrease
in nuclear attraction
Periodic Trends
Ionization Energy
Sample Problem A
• Referring to a periodic table, arrange the following atoms in order of
increasing first ionization energy: Ne, Na, P, Ar, K
Periodic Trends
Ionization Energy
Sample Problem B
• Which has the lowest first ionization energy, B, Al, C, or Si? Which has the
highest?
Periodic Trends
Electron Affinity
• Electron affinity is the energy change associated with the addition of an
electron.
• increase across a period
• decreases down a group
• Electrons farther from the nucleus experience less nuclear attraction
• Some irregularities due to repulsive forces in the relatively small p-orbital.
Periodic Trends
Electron Affinity
Sample Problem C
• The electron affinity of lithium is a negative value, where as the electron
affinity of beryllium is a positive value. Use electron configuration to
account for this observation.
Periodic Trends
Atomic Radius
• Atomic radius is half the distance between radii in a covalently bonded
diatomic molecule.
• decreases across a period
• increase in effective nuclear charge (Zeff) due to
decreased shielding
• more protons pulling on the electrons
• increases down a group
• addition of principal quantum levels makes the electron
farther from the nucleus
Periodic Trends
Atomic Radius
Sample Problem D
• Referring to a periodic table, arrange the following atoms in order of
increasing size: P, S, As, Se.
Periodic Trends
Ionic Radius
• Anions (negative ions) are larger than their parent atom.
• due to a greater repulsion of fewer protons and more electrons.
• Cations (positive ions) are smaller than their parent atom.
• due to a greater attraction of more protons and fewer electrons.
• Ionic radius follows the same trend as atomic radius
• Isoelectronic Ions
• ions with the same number of electrons
• size decreases as the nuclear charge increases
Periodic Trends
Ionic Radius
Sample Problem E
+
2+
2• Arrange the ions, K , Cl , Ca , and S , in order of decreasing size.
Periodic Trends
Photoelectron Spectroscopy (PES)
• Photoelectron Spectroscopy (PES) is a technique that is used to gather
information about the electrons in an atom.
• An atom is bombarded with photons. Some of the photons are absorbed
and electrons are emitted. The energy is analyzed. Since we can know
the energy of the photons, and we know that energy is conserved we know
that the difference in energy between the photons sent into the atom and
the energy of the electrons emitted will be the potential energy of the
electrons when they are attached to the atom. Remember that the
potential energy of the electron in the atom is the work needed to remove
the electron from the atom.
• Energy of emitted electron = energy of photon - work needed to remove
electron from atom
Periodic Trends
Photoelectron Spectroscopy (PES)
• How do scientist find ionization energy?
• PES can tell what about an atom?
• PES provides evidence for ___.
• How does PES work?
• The size of a peak on a graph means ___.
• Electrons are held most tightly ___.
Periodic Trends
Photoelectron Spectroscopy (PES)
Hydrogen
• The whole number is the integration of the number of electrons in the
spectrum. The decimal number is the work needed to remove the electron
from the atom in MJ/mol.
Periodic Trends
Photoelectron Spectroscopy (PES)
Helium
• Helium has an integration of 2 and the energy is higher than hydrogen. This
is because He has 2 electrons and H has 1. Also, H and He have the same
level of shielding (none) while He has 2 protons and H has 1. That makes it
harder to remove the electrons from He than from H.
Periodic Trends
Photoelectron Spectroscopy (PES)
Lithium
• Lithium has two sets of electrons. One electron has lower energy, the 2
electrons have higher energy.
• On the PES spectra, a high energy number means the electron is closer to
the nucleus. A low number means it is further from the nucleus.
Periodic Trends
Photoelectron Spectroscopy (PES)
Lithium
• Lithium has two sets of electrons. One electron has lower energy, the 2
electrons have higher energy.
• Note that the peak for an electron farther from the nucleus is closer to the
beginning of the graph while the peak for an electron closer to the nucleus
is farther from the beginning of the graph.
• Think of the energy number on the graph as how hard it is to take the
electron away from the atom.
Periodic Trends
Photoelectron Spectroscopy (PES)
Boron
• Boron has 2 electrons at 19.3 in the 1s. There are 2 electrons at 1.36 and 1
electron at 0.80. The 2 electrons are 2s and the one electron is 2p. Why are
they different? They are both in the same energy level (shell), so why do
they have diffent potential energies?
Periodic Trends
Photoelectron Spectroscopy (PES)
Boron
• The reason the 2s and 2p orbitals have different energy is that s penetrates
better than p. An s orbital can overcome the effects of shielding better than
p. Remember that penetration ability is: s > p > d > f
Periodic Trends
Photoelectron Spectroscopy (PES)
Boron
• So it is easier to remove an electron from the 2p orbital in B than from the
2s. That’s why the 2s are at 1.36 and the 2p is at 0.80.
Periodic Trends
Photoelectron Spectroscopy (PES)
Scandium
Periodic Trends
Photoelectron Spectroscopy (PES)
Scandium
• The filling order has 3d after 4s
because 3d does not penetrate as
well as 4s.
• However, in the spectrum for Sc, we
can see that there are 2 electrons at
0.62 and one at 0.77
• This suggests that it is easier to remove the 4s than the 3d.
• This is because n=3 shields n=4. This raises the energy of 4s once 3d
starts to fill.
Periodic Trends
Photoelectron Spectroscopy
Sample Problem F
• Identify the element shown above.
Periodic Trends
Photoelectron Spectroscopy
Sample Problem G
• What happens to the 1s peak?
• What is different about the
graph than the order you fill the
electron orbital diagrams?
• Why is the 4s lower in energy
than 3d?
Periodic Trends
Periodic Trends
Sample Problem H
Account for each of the following in terms of principles of atomic structure,
including number, properties, and arrangements of subatomic particles.
a.The second ionization energy of sodium is about three times greater than
the second ionization energy of magnesium.
b.The difference between atomic radii of Na and K is relatively large
compared to the difference between atomic radii of Rb and Cs.
c.A sample of nickel chloride is attracted into a magnetic field, whereas a
sample of solid zinc chloride is not.
d.Phosphorus forms the fluorides PF3 and PF5, whereas nitrogen forms only
NF3.
Periodic Trends
Periodic Trends
Sample Problem I
Explain each of the following observations using principles of atomic
structure and/or bonding.
a.Potassium has a lower first ionization energy than lithium.
b.The ionic radius of
3N
is larger than that of
2O .
c.A calcium atom is larger than a zinc atom.
d.Boron has a lower first ionization energy than beryllium.