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Unit 13 Chemical Periodicity
A Little Bit of History
C In the early 1800s, German chemist J.W. Dobereiner observed that several elements could be
classified into sets of three which he called triads.
C (Li, Na, K), (Ca, Sr, Ba), (Cl, Br, I)
C The elements within each triad had similar properties.
C Several physical properties of the middle elements were averages of the other two.
C 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.
C The first element was like the eighth; the second element was like the ninth; etc
C In 1869, Russian chemist Dmitri Mendeleev and German chemist Lothar Meyer published nearly
identical schemes for classifying elements.
C Mendeleev got the credit because he published his first and he was more successful at
demonstrating its value.
C Mendeleev eventually produced the first periodic table of the elements.
C He arranged his table so that elements in the same column have similar properties.
C He switched the order of three pairs of elements so as to keep elements in columns with
similar properties.
C He boldly pronounced that perhaps the calculated atomic masses for those mixed up
elements should be recalculated.
C Mendeleev also had the insight to predict the existence and some of the properties of three
new elements.
C In 1913, English chemist Henry Moseley determined the atomic numbers of the elements. He did
this with X-ray analysis. We now know about protons.
C The periodic law states that when elements are arranged in order of increasing atomic
number, their physical and chemical properties show a periodic pattern.
Classification of the Elements
C Different kinds of periodic tables may present different kinds of information or present information
in different ways.
C The shape of the periodic table comes from the periodic law.
C Elements with similar properties are aligned in vertical columns called groups or families.
C The horizontal rows are called periods.
C Some of the groups have family names. (Alkali metals, alkaline earth metals, halogens, and the
noble gases.)
C Elements can be classified as metals, nonmetals, and semimetals.
C Metals are typically solids at room temperature.
C characteristic shine, conduct heat and electricity, malleable (hammered into sheets),
ductile (drawn into wires), lose e-‘s to become cations
C Nonmetals can be solids, liquids, or gases
C do not have luster (dull), poor conductors of heat and electricity, brittle.
C 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
C Semimetals have some properties of metals and some of nonmetals, or they have properties
that are intermediate.
C The outermost “s” and “p” electrons, which are largely responsible for an atom’s chemical
behavior, are called valence electrons.
C The elements in a group have similar properties because they have the same number of
valence electrons (which means similar electron configurations).
C To save space in writing electron configurations and to focus attention on valence electrons,
C
C
chemists often use abbreviated electron configurations.
C An atom’s inner electrons are represented by the symbol for the nearest noble gas with a
lower atomic number. (Noble gas inner core)
Octet rule is a very important concept to understanding bonding in chemistry.
C Octet refers to the eight outer electrons of an atom.
C An atom with all of its “s” and “p” electrons is extremely stable.
The key to understanding the shape of the periodic table is to examine the elements’ electron
configuration. (The fours main sections of the periodic table correspond to the four sublevels, s p
d f.)
C The s-block elements are groups 1 and 2.
C The p-block elements are groups 3-8.
C The d-block elements are the elements whose last electrons fill the d sublevel.
C The f-block elements are the elements whose last electrons fill the f sublevel.
C 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.
C The four blocks have other names as well.
C The elements in the s- and p- blocks are called the representative elements.
C The elements in the d-block are called the transition metals.
C Those in the f-block are known as the inner transition metals.
Periodic Trends
C Many properties of the elements change in a predictable way as you move through the periodic
table. These systematic variations are called periodic trends.
Î The atomic radius is the distance from the center of an atom’s nucleus to its outermost electron.
C Atoms get larger going down a group. (ie. size increases with increasing atomic #)
C Atoms get smaller moving from left to right across each period. Why?
C 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.
C A stronger attractive force shrinks the electrons’ orbitals and makes the atom smaller.
C As you learned in Unit 5, an atom can gain or lose electrons to form an ion.
C When an atom loses electrons it becomes smaller.
C 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 closed to the nucleus.
C When an atom gains electrons it becomes larger.
C 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.
C The increased repulsions spread out the electrons, thus making the ion larger than the
atom.
C The atoms of a particular element typically form only certain ions.
C The elements on the left side of the table form positive ions.
C The elements on the right side of the table form negative ions.
C The elements in the eighth column do not form ions.
Ï An atom’s ionization energy is the energy needed to remove one of its electrons.
C You can think of ionization energy as a reflection of how strongly an atom holds onto its
outermost electron.
C 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.
C Ionization energies generally increase as you move from left to right across a period.
C
(Metals have low ionization energies while the noble gases have the highest ionization
energies.)
C There will be peaks and valleys due to filled sublevels or partially-filled sublevels.
C 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.
C The successive ionization energies are the energies required to remove electrons beyond the
first electron.
C For each element you can find one very large increase between a different pair of ionization
energies.
C This happens whenever the inner core of electrons are attacked.
Ð An atom’s electronegativity reflects its ability to attract electrons in a chemical bond.
C Electronegativity increases as you go left to right across the period.
C Electronegativity decreases as you go down a group.
C Metals have low electronegativity while nonmetals have high electronegativity.
C Fluorine has the highest and francium has the lowest.
Ñ Another important trend is the idea of shielding effect.
C Shielding effect is the result of full energy levels separating the outermost electrons from the
nucleus.
C An element like lithium only has one energy level separating its outer electron from the
nucleus.
C But, an element like francium has many energy levels separating its outer electrons from
the nucleus.
C This shielding effect is the reason why ionization energy, electron affinity, and electronegativity
decreases as you go down a group.
C Shielding effect increases as you go down a group.
C Shielding effect remains the same as you go across a period.
Summary of Periodic Trends
Remains the same 6 Shielding Effect
Increases 6 Ionization Energy, Electronegativity
Decreases 6 Atomic size, Ion size
1A
Decreases 9
Ionization Energy
Electronegativity
Increases 9 Nuclear Charge
Atomic size
Shielding Effect
Ion size
8A
H
2A
3A
4A
5A
6A
7A
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Ga
Ge
As
Se
Br
Kr
Rb
Sr
In
Sn
Sb
Te
I
Xe
Cs
Ba
Tl
Pb
Bi
Po
At
Rn