Periodic Trends
Elemental Properties and Patterns
• I CAN use the PT to predict a number
of properties among the elements.
The Periodic Law
• Though he didn’t realize it at the time,
Mendeleev made an important discovery
about the elements.
• Mendeleev’s repeating patterns on his
Periodic Table led to the formulation of 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.
What Does the Periodic Table Tell Us?
• Recall that the blocks of the periodic table are
arranged in to LEFT TO RIGHT ROWS called a
PERIOD/SERIES and UP AND DOWN COLUMNS called
a FAMILY/GROUP.
• Each square on the periodic table is known as an
ELEMENT KEY and provides the basic information
about each element:
• CHEMICAL SYMBOL – a one or two letter
abbreviation for the name of an element.
– Some symbols are based on the Greek or Latin name of
the element.
• Example Lead = Pb from Plumbus (where we get plumber from).
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Gold = Au from Arium (to shine like the sun).
• ELEMENT NAME – the name of the element.
– Some names are more than two thousand years
old!
• ATOMIC NUMBER – Number of protons
(nucleus) or electrons (electron cloud).
• ATOMIC MASS NUMBER – combined number
of protons and neutrons in the nucleus.
VALENCE ELECTRONS
• One of the most important thing about an atom
is the number of VALENCE ELECTRONS it has.
• Valence Electrons are those electrons found on
the outside edge of the atom, farthest away
from the nucleus.
Valence Electrons are those on the
outer edge of the atom.
• An atom will have from 1 to 8 valence
electrons, abbreviated as Ve- .
• Valence electrons are used to bond to other
atoms to form COMPOUNDS.
Valence Electrons
• Determining the number of Valence Electrons an
element has is easy:
• Look at the GROUP NUMBER the element is in.
• Remember the GROUP NUMBER is the numbers
1 to 18 across the tops of the vertical columns.
• IGNORE the TRANSITION ELEMENTS (groups
3-12) for now.
• ONLY LOOK AT GROUPS 1, 2, 13, 14, 15, 16,
17 and 18.
• The number of Valence Electrons is from the group
number (Groups 13-18 just use the SECOND
number):
• Group 1 = 1 Ve• Group 2 = 2 Ve• Group 13 = 3 Ve• Group 14 = 4 Ve• Group 15 = 5 Ve• Group 16 = 6 Ve• Group 17 = 7 Ve• Group 18 = 8 Ve• We will learn the Transition Elements later!
PRACTICE PROBLEMS
• Look at the questions on the back of this packet.
• Use the periodic table on the NEXT
problems.
SLIDE, to complete these
• I. FIND THE ELEMENT – use the period/series and family/group
information given in each statement to find the element referred to on the
Periodic Table.
• II. HOW MANY VALENCE ELECTRONS – use the position of each element
on the Periodic Table to determine how many Valence Electrons each
element has.
If you can’t read this table, follow this link to display it online:
http://www.ptable.com/Images/periodic%20table.png
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I. FIND THE ELEMENT – use the period/series and family/group information given in each statement to find the
element referred to on the Periodic Table.
a. Period 3 Group 15 ______________________
b. Period 6 Group 12 ________________________
c. Period 2 Group 17 ______________________
d. Period 3 Group 2 _________________________
e. Period 7 Group 11 ______________________
f. Period 4 Group 16 ________________________
g. Period 1 Group 18 _______________________
h. Period 3 Group 13 ________________________
i. Period 5 Group 4 ________________________
j. Period 4 Group 2 __________________________
k. Period 6 Group 14 ______________________
l. Period 4 Group 16 _________________________
m. Period 3 Group 14 ______________________
n. Period 6 Group 11 ________________________
II. HOW MANY VALENCE ELECTRONS – use the position of each element on the Periodic Table to determine how
many Ve- each element has.
o. Hydrogen ______
p. Iodine ________
q. Calcium ___________
r. Oxygen _____________
s. Boron _________
t. Magnesium _____
u. Francium _________
v. Tin ___________
w. Carbon ________
x. Xenon _________
y. Sulfur _____________
z. Potassium __________
•STOP HERE
Grouping the Elements
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Elements are often classified by their type:
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.
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 is a metalloid?
Metals
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Metals are:
Lustrous
Have a shiny surface.
Malleable
Can be hammered thin.
Ductile
Can be stretched (wire)
Good Conductors
Allow heat and electricity to flow
through easily.
• They are mostly solids at room temp.
• What is one exception?
• Mercury (Hg) is a LIQUID metal!
Nonmetals
• Nonmetal properties are the basically
opposite of metals:
• Dull
• Not shiny/can’t be polished.
• Brittle
• Break easily…cannot be hammered thin.
• Poor/non-conductors
• Limited or no flow of heat or electricity
through them.
– (also called insulators).
• Some are solid, several are gases, and
BROMINE is a liquid.
Metalloids
• Metalloids, aka semi-metals are just
that.
• They have characteristics of both
metals and nonmetals.
• They are shiny but brittle.
• They are semiconductors: elements
that have conduction capacities
between non metals and metals.
• What are our most important
semiconductors?
• Silicon, Germanium, and Arsenic.
Periodic Trends
• ATOMIC RADIUS
• 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 a defined covalent radius, or half
the distance between the nuclei of 2 bonded
atoms.
• Atomic radii are usually measured in
picometers (pm = 1 x 10-9 m) or angstroms (Å).
An angstrom is 1 x 10-10 m.
Example of Covalent Radius
• Two Br atoms bonded together are 2.86
angstroms apart.
2.86 Å
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Arad = 2.86 Å = 1.43 Å
2
Atomic Radius
• The trend for atomic radius in a
FAMILY/GROUP (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
• 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).
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
closer to the nucleus, making atoms smaller as
we move from left to right across a period.
INCREASES TOWARD BOTTOM LEFT
Overall Trend of Atomic Radius
DECREASES TOWARD TOP RIGHT
e- Shielding
• As more PELs are added to atoms, the inner
levels of electrons shield the outer electrons
from the nucleus’ attraction.
• The effective nuclear attraction for these
outer electrons is less, and so the outer
electrons are less tightly held.
• This makes it easier to remove some of these
e- when bonding to other atoms.
Ionization Energy
• This is the second important periodic trend.
• If an electron is pulled toward another atom with
enough energy to overcome the effective nuclear
attraction of the parent atom holding the electron
in the cloud, it can leave the atom completely.
• Once an electron has been removed from one atom
and added to another, both atoms are changed.
• These atoms have been “ionized”.
• The number of protons and electrons is no longer
equal.
• The atom that lost an electron becomes POSITIVELY
charged (more P+ than e-) while the atom that
gained an electron becomes NEGATIVELY charged
(more e- than P+ ).
Ionization Energy
• The energy required to remove an electron
from an atom is ionization energy.
• 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
• Draw arrows on your help sheet like this:
INCREASES LEFT TO RIGHT
DECREASES
DOWNWARD
TOP TO
BOTTOM
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
• Your help sheet should look like this:
+
+
Metallic Character
• This is simple a relative measure of how easily
atoms lose or give up electrons.
• Your help sheet should look like this:
Electronegativity
• Electronegativity is a measure of an atom’s
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
• Your help sheet should look like this:
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
• Your help sheet will look like this:
0
The Octet Rule
• The “goal” of most atoms (except H, Li and
Be) is to have an octet or group of 8
electrons in their valence energy level.
• They may accomplish this by either giving
electrons away or taking them.
• Metals generally give electrons, nonmetals
take them from other atoms.
• Atoms that have gained or lost electrons
are called ions.
Ions
• When an atom gains an electron, it becomes
negatively charged (more electrons than
protons ) and is called an anion.
• In the same way that nonmetal atoms can gain
electrons, metal atoms can lose electrons.
• They become positively charged cations.
Ions
• Here is a simple way to remember which is the
cation and which the anion:
+
This is Ann Ion.
She’s unhappy and
negative.
+
This is a cat-ion.
He’s a “plussy” cat!
Ionic Radius
• Cations are always smaller than the original
atom.
• The entire outer PEL is removed during
ionization.
• Conversely, anions are always larger than the
original atom.
• Electrons are added to the outer PEL.
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.