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Periodicity
Trends in the Periodic Table
Electron Dot Diagrams
• Atoms ca be represented by electron dot
diagrams. The dots on the dot diagram
identify only the outside shell electrons.
Steps For Drawing Electron Dot
Diagrams
• Write the atomic symbol for the atom. This will
represent the nucleus and the inner energy levels.
• Use a dot to represent an outer shell electron.
• One dot is placed in each of the four sides before any
pairing occurs.
• Beginning with the fifth dot, pairing may occur to a
maximum of eight dots. This is the octet rule.
• Unpaired electrons are called bonding
electrons because they are involved in bond
formation.
• Paired electrons are called lone pairs and are
generally not involved in bond formation,
however some complex molecules and ions do
use the lone pairs to bond.
• So far you have learned that the elements are
arranged in the periodic table according to
their atomic number. You have also learned
that there is a correlation between the
arrangement of elements and their electron
configurations.
• There is also a correlation between how the
elements are arranged and their atomic radii,
ionization energy, electron affinity, ionic radii,
valence electrons, and electronegativity.
Explaining Reactivity
• Period Patterns:
– Each element’s period number is the same as the
number of shells it’s electrons occupy
– 2nd row has two shells (two energy levels)
• Group Patterns:
– All elements in the main group have the same
number of electrons in their outer shell. These are
called valence electrons.
– Elements in group 2 have two electrons in their
outer shell
Valence Electrons
• Chemical compounds form because electrons are
lost, gained, or shared between atoms. The electrons
that interact in this manner are those electrons in
the highest energy levels.
• The electrons available to be lost, gained, or shared
in the formation of chemical compounds are referred
to as valence electrons.
• Atoms with partially filled outer energy levels are
unstable.
• The close the element is to a noble gas, the more
reactive/unstable it is
Valence Electrons in Main-Group Elements
(any main energy level)
Number of Valence Electrons
Group Number
1
1
2
2
13
3
14
4
15
5
16
6
17
7
18
8
Graphing Activity
• Complete the graphing activity
‘Trends in the periodic table’
– Do questions 1-7
Atomic Radii
• The atomic radius is defined as half the distance
between the nuclei of identical atoms bonded
together. It is measure in picometers.
TREND: Atomic Radius
• Decrease across a period due to increasing positive
nuclear charge & increased # e- on shell (more
attraction)
• Increase down a group due to increasing number of
energy levels (outer electrons are farther from the
nucleus = less attraction)
Ionization Energies
• An electron can be removed from an atom if
enough energy is supplied.
• Li + energy  Li+ + e-
• Ionization energy is the energy required to
remove an electron from an atom.
TREND: Ionization Energy
Ion: An atom or group of atoms that has a
positive (cation) or negative (anion) charge
Ionization: The formation of an ion
Ionization Energy: The energy required to
remove one electron from a neutral atom
of an element, measured in kilojoules/mole
(kj/mol)
Ionization Energy
Think about where you would hold a basketball to
protect it… where would it take the most
energy to rip that ball away?
1. IE tends to increase across each period.
•
Atoms are getting smaller, electrons are closer to the
nucleus
2. IE tends to decrease down a group
•
•
3.
4.
5.
Atoms are getting larger, electrons are farther from the
nucleus
Outer electrons becomes increasingly more shielded from
the nucleus by inner electrons
Metals have characteristically low IE
Nonmentals have high IE
Noble gases have a very high IE
• Group 1 elements have the lowest ionization
energies. They lose electrons most easily.
• Group 17 elements have high ionization
energies because they do not lose electrons
easily. These elements would rather gain an
electron.
• Going across a period, from left to right, the
ionization energy increases.
• As you go down a group, the electrons are
further and further away from the nucleus,
and are therefore easier to remove.
• Going down a group, the ionization energies
decrease.
Pop Quiz!
1. Rank these
elements by
decreasing atomic
size:
a) Se, Br, Ca
b) N, Bi, As
2. Which has the
greater ionization
energy?
a) P or Bi?
b) P or Ar?
Complete Questions 8-10 form graphing sheet
Start working on your big trends review
Ionic Radii
• A positive ion is known as a cation.
– The formation of a cation by loss of one or more electrons
always leads to a decrease in radius. The remaining
electrons will be pulled in tighter to the nucleus.
• A negative ion is known as an anion.
– The formation of an anion by the addition of one or more
electrons always leads to an increase in radius.
– This is because the total positive charge of the nucleus
remains the same when an electron is added so the
electrons are not drawn to the nucleus as strongly as they
were before the addition.
– The electron cloud spreads out because of the greater
repulsion between the increased number of electrons.
• The metals on the left tend to form cations
while the non metals on the right will form
anions. Cationic radii (groups 1,2,13, and 14)
decrease going from left to right. Anionic radii
also decreases going from left to right.
• Just as there is an atomic radii increase as you
go down a group, there is an ionic radii
increase. The electrons are found in shells
farther away from the nucleus and are harder
to keep close.
TREND: Ionic Radius
• Cations have a smaller ionic radius than
corresponding atom
• Protons outnumber electrons
• Less shielding of electrons
• Anions have a larger ionic radius than
corresponding atom
• Electrons outnumber protons
• Greater electron-electron repulsion
• Ion size decreases down a group
Electron Affinity
• Neutral atoms can also acquire electrons. The
energy that occurs when an electron is
acquired by a neutral atom is called the
electron affinity.
• Most atoms release energy when they acquire
an electron.
• Cl + e-  Cl1- + energy
• Sodium on the other hand would rather give
up an electron than take one in. If an atom is
forced to take an electron, not much energy is
released due to the amount of energy used to
force the electron in.
• Na + e- + energy  Na+
• As you move across the table from left to
right, the electron affinity increases.
• As a general rule, electrons add with greater
difficulty as you move down a group,
therefore electron affinity decreases.
• This is because the atoms use up a lot of the energy to
hold electrons close to the nucleus. As shells get bigger,
the atom uses more energy to keep electrons close.
Adding another electron to try and keep close takes up
more energy than can be released.
TREND: Electron Affinity
1.
2.
3.
4.
EA tends to increase across each period.
EA tends to decrease down a group
Halogens have the highest electron affinities
Metals have low electron affinities
Electronegativity
• Valence electrons hold atoms together in a chemical
compound. In many compounds the negative charge
of the valence electrons is concentrated closer to
one atom than the other.
– This uneven concentration of charge has a significant
effect on the chemical properties of a compound.
– It is therefore useful to have a measure of how strongly
one atom attracts the electrons of another atom within a
compound.
• Electronegativity is a measure of the ability of
an atom in a chemical compound to attract
electrons.
• As you move from left to right on the periodic
table, the electronegativity increases.
• As you move from top to bottom in a group on
the periodic table, the electronegativity
decreases.
Electronegativity
The ability of an atom to attract electrons when bonded
1. EN tends to increase across each period atomic radius is
smaller, bonding pair can get closer to the nucleus and be
more attracted
2. EN tends to decrease down a group atomic radius is larger,
attracts bonding pairs less strongly
3. Nonmetals have high electronegativity
4. Metals have low electronegativity
Electronegativity
Electronegativity
Pop Quiz!
1. Which element has the greatest electron
affinity?
a) F, Br, Cl
b) I, Rb, Sn
2. Which element is most electronegative?
a) Po, O, Se
b) O, Li, C