Download Periodicity

Unit 13
Chemical Periodicity
A Little Bit of History
In the early 1800s, German chemist J.W.
Dobereiner observed that several elements
could be classified into sets of three which he
called triads.
• (Li, Na, K), (Ca, Sr, Ba), (Cl, Br, I)
• The elements within each triad had similar
properties.
• Several properties of the middle elements
were averages of the other two.
In 1865, English chemist J.A.R. Newlands
observed that when the then 62 known elements
were arranged in order of increasing atomic mass,
every eighth element had similar properties.
• This was known as the Law of Octives
• The first was light the eighth; Na was like Li
• The second was like the ninth, Mg was like
Be
• He knew nothing of the noble gases
In 1869, Russian chemist Dmitri Mendeleev and
German chemist Lothar Meyer published nearly
identical schemes for classifying elements.
• Mendeleev got the credit because he published
his first and he was more successful at
demonstrating it value.
• Mendeleev eventually produces the first working
periodic table of the elements.
• He arranged his table so that elements in the
same column had similar properties.
• He switched the order of three pairs of
elements so as to keep elements in columns
with similar properties
• He switched the order of three pairs of
elements so as to keep elements in columns
with similar properties.
• This was contrary to all other chemists who
arranged elements according to increasing
atomic mass.
• He boldly pronounced that perhaps the
calculated atomic masses for those mixed up
elements should be recalculated
• Mendeleev also had the insight to predict the
existence and some of the properties of three
new elements (these elements were missing
from his periodic table.
In 1913, English chemist Henry Moseley
determined the atomic number of the elements.
• He already know that protons are positively
charged particles that exist inside of nuclei.
• He used x-rays to determine that each element
had a certain number of positive particles in its
nucleus.
• He said these positive particles must be the
protons discovered by Goldstein.
• This discovery led to the term atomic number.
The modern periodic law states that when
elements are arranged in order of
increasing atomic number, their physical
and chemical properties show a periodic
pattern.
Mosley’s discovery proved Mendeleev to be
correct.
Different kinds of periodic tables may present different
kinds of information or present information in different
ways.
• The shape of the periodic table comes from the
periodic law.
• Elements with similar properties are aligned in
vertical columns called groups or families.
• The horizontal rows are called periods.
• Some of the groups have family names. (Alkali
metals, alkaline earth metals, halogens, and the
noble gases.)
• Elements can be classified as metals, nonmetals, and
semimetals.
• Metals are typically solids at room temperature.
• characteristic shine, conduct heat and electricity,
malleable (hammered into sheets), ductile (drawn
into wires), lose e-‘s to become cations
• Nonmetals can be solids, liquids, or gases
• do not have luster (dull), poor conductors of heat
and electricity, brittle.
• quite a variation in physical properties (some are
colored, others are colorless; some are soft solids
while others form hard solids); gain e-‘s to become
anions
• Semimetals have some properties of metals and some
of nonmetals, or they have properties that are
intermediate.
The outermost “s” and “p” electrons, which are largely
responsible for an atom’s chemical behavior, are called
valence electrons.
• The elements in a group have similar properties
because they have the same number of valence
electrons (which means similar electron
configurations).
• To save space in writing electron configurations and
to focus attention on valence electrons, chemists
often use abbreviated electron configurations.
• An atom’s inner electrons are represented by the
symbol for the nearest noble gas with a lower
atomic number. (This is called the noble gas inner
core)
Octet rule is a very important concept to understanding
bonding in chemistry.
• Octet refers to the eight outer electrons of an atom.
• An atom with all of its “s” and “p” electrons is
extremely stable.
• The noble gases have those eight outer electrons and
will not bond with another atom.
• Metals will lose electrons to back up to have the same
number of electrons as the previous energy level’s
noble gas.(they become cations)
• Nonmetals will gain electrons to have the same number
of electrons as the noble gas at the end of the row.
(they become anions)
• Sometimes nonmetals share electrons to gain the
octet
The key to understanding the shape of the periodic table is
to examine the elements’ electron configuration. (The four
main sections of the periodic table correspond to the four
sublevels, s p d f.)
• The s-block elements are groups 1 and 2.
• The p-block elements are groups 3-8.
• The d-block elements are the elements whose last
electrons fill the d sublevel.
• The f-block elements are the elements whose last
electrons fill the f sublevel.
The shape of the periodic table is a result of the way the
electrons fill the s, p, d, and f orbitals of the different
energy levels.
• The four blocks have other names as well.
• The elements in the s- and p- blocks are called the
representative elements.
• The elements in the d-block are called the transition
metals.
• Those in the f-block are known as the inner transition
metals.
Periodic Trends
Many properties of the elements change in a
predictable way as you move through the
periodic table. These systematic variations are
called periodic trends. The periodic trends that
you must understand are…
• Atomic radius
• Ionization energy
• Electronegativity
• Shielding Effect
Atomic Radius
The atomic radius is the distance from the center of
an atom’s nucleus to its outermost electron.
• Atoms get larger going down a group. (ie. size
increases with increasing atomic #)
• Atoms get smaller moving from left to right
across each period. Why?
– As a rule, atoms that have more positive charge in
their nuclei exert a stronger pull on the electrons in a
given principal quantum number.
– A stronger attractive force shrinks the electrons’
orbitals and makes the atom smaller.
• As you learned in Unit 5, an atom can gain or
lose electrons to form an ion.
• When an atom loses electrons it becomes
smaller.
• Loss of electrons not only vacates the atom’s
largest orbitals, it also reduces the repulsive
force between the remaining electrons,
allowing them to be pulled closer to the
nucleus.
• When an atom gains electrons it becomes
larger.
• A fluoride ion is larger than the fluorine atom
mainly because of the greater number of
electrons, which increases the electric
repulsive forces among them.
• The increased repulsions spread out the
electrons, thus making the ion larger than the
atom.
• The atoms of a particular element typically
form only certain ions.
• The elements on the left side of the table form
positive ions. (cations)
• The elements on the right side of the table
form negative ions. (anions)
• The elements in the eighth column do not
form ions. (noble gases)
• An atom’s ionization energy is the energy needed
to remove one of its electrons.
• You can think of ionization energy as a reflection
of how strongly an atom holds onto its outermost
electron.
• Atoms with high ionization energies hold onto
their electrons very tightly, whereas atoms with
low ionization energies are more likely to lose
one or more of their outermost electrons and
gain a positive charge.
• Ionization energies generally increase as you
move from left to right across a period.
• Metals have low ionization energies while the
noble gases have the highest ionization
energies.
• There will be peaks and valleys due to filled
sublevels or partially-filled sublevels.
• Ionization energies generally decrease as you
move down a group.
• The energy required to remove the first
electron from an isolated atom is called the
first ionization energy.
• The successive ionization energies are the
energies required to remove electrons beyond
the first electron.
• For each element you can find one very large
increase between a different pair of ionization
energies.
• This happens whenever the inner core of
electrons are attacked.
Electronegativity
• An atom’s electronegativity reflects its ability to
attract electrons in a chemical bond.
• Electronegativity increases as you go from left to
right across the period.
• Electronegativity decreases as you go down a
group.
– Metals have low electronegativity while nonmetals
have high electronegativity.
– Fluorine has the highest and francium has the lowest.
Shielding Effect
• Shielding effect is the result of full energy levels
separating the outermost electrons from the
nucleus.
• An element like lithium only has one energy level
separating its outer electron from the nucleus.
(therefore low shielding effect)
• But, an element like francium has many energy
levels separating its outer electrons from the
nucleus. (therefore high shielding effect)
• This shielding effect is the reason why
ionization energy and electronegativity
decreases as you go down a group.
• Shielding effect increases as you go down a
group.
• Shielding effect remains the same as you go
across a period.