Name _________________________________________________________ Date __________________________ Period ________
Periodic Table Trends
(Ionization Energy and Electronegativity)
Ionization Energy – The _______________ required to _______________ an electron from a gaseous atom or ion.
Period Trend:
 As the atomic number increases, ionization energy ____________________.
WHY?
 _______________ _______________ increases; Electrons are held more ________________ to the nucleus.
 Also, the outer energy level is closer to being an _______________.
Example 1: Circle the element with the higher 1st ionization energy.
Ca
or
Br
Group Trend:
 As the atomic number increases, ionization energy ____________________.
WHY?
 Outer electrons are in _______________ energy levels.
 Farther from the nucleus, which means more _______________; therefore, the electrons are more _______________
removed.
Example 2: Circle the element with the higher 1st ionization energy.
Na
Practice 1: Which has a higher ionization energy?
Mg
or
Ba
Which has a higher ionization energy?
Al
or
Si
Which has a higher ionization energy?
Mg
or
Cl
or
Cs
2nd Ionization Energy – The _______________ required to _______________ the ______________ electron. (There can also be 3 rd,
4th and so forth ionization energies.
Example 3: Aluminum
Al(g)  Al (g) + e
+
Al (g)  Al (g) + e
+
2+
I1 = 580 kJ/mol
-
Al (g)  Al (g) + e
2+
3+
Why is there an increase in successive ionization energies?
-
I2 = 1815 kJ/mol
-
I3 = 2740 kJ/mol
Al3+(g)  Al4+(g) + e- I4 = 11600 kJ/mol
Example 4: Consider atoms with the following electron configurations:
1s22s22p6
Which atom has the largest 1st ionization energy?
1s22s22p63s1
1s22s22p63s2
Which one has the smallest 2nd ionization energy?
Practice 2: The successive ionization energies for an unknown element are
I1 = 786 kJ/mol
I2 = 1,577 kJ/mol
To which family in the periodic table does the unknown element most likely belong?
Explain.
I3 = 3,232 kJ/mol
I4 = 4,355 kJ/mol
I5 = 16,091 kJ/mol
Electronegativity – The ability of an atom to _______________ electrons to itself in a chemical bond.
Period Trend:
 As the atomic number increases, electronegativity ____________________.
WHY?
 The atoms are getting _______________ in size with more protons in the nucleus, so they have more ability to
_______________other elements’ electrons and pull them toward themselves.
Example 5: Circle the element with the higher electronegativity.
Be
or
O
Group Trend:
 As the atomic number increases, electronegativity ____________________.
WHY?
 Elements at the ________________ of a group have electrons _______________ held by the nucleus.
 Elements at the bottom of a group are _______________ and have more _______________ (and _______________)
between the nucleus and outer/valence electrons.
Example 2: Circle the element with the higher electronegativity.
N
or
Sb
Name _________________________________________________________ Date __________________________ Period ________
CHEMICAL BONDING
_______________ _______________ describes the _______________ that hold adjacent atoms together in a compound.
3 General Types of Bonds
1.
IONIC BONDS – Form when one or more _______________ electrons are _______________ from one atom to another,
creating _______________ and _______________ ions.
a.
2.
COVALENT BONDS – Involves _______________ of valence electrons between atoms.
a.
3.
Properties of Ionic Bonds
i. High _______________ and _______________ points.
ii. Crystalline (__________) when dried and are _______________.
iii. Often ______________________________.
iv. Conduct electric current (electricity) in the _______________ form and _______________ forms.
Properties of covalent Bonds
i. _______________ melting and boiling points.
ii. _______________ electric conductor in any form.
iii. Exist as whole _______________, not ions.
iv. Most are _______________ in water.
METALLIC BONDS – Form in __________ atoms (positive metal ions with _______________ electrons). The force
holding the metal together is the electrostatic attraction among _____________________ and _______________.
a.
Properties of Metallic Bonds
i. Good _______________ ________________ - Electrons
flow _______________ in metals, conducting electrical
signals.
ii. Good _______________ ________________ - Free flowing
electrons transmit _______________.
iii. _______________ melting and boiling points.
iv. _______________,
_______________,
and
_______________.
b.
Alloys – A material composed of __________ or __________ metals.
Examples:
1. Brass – made of _______________ and ______________.
2. 14 Karat Gold – made of _______________ and ______________.
3. Bronze – made of _______________ and ______________.
Chemical reactions result in the _______________, _______________or ____________________ of valence electrons. So, only the
valence electrons are involved in _______________.
Lewis Electron Dot Symbols



A useful way to represent electrons in the valence shell of an atom.
The symbol of the element represents the atomic nucleus together with __________ electrons.
______________ electrons are represented by __________ and are placed one-by-one around the
element symbol.
Draw the Lewis electron dot symbols for each element in period 2.
How do you draw Lewis electron dot symbols for IONS?
To draw the Lewis electron dot symbol of cations:
1.
2.
3.
4.
_______________ the same number of electrons as the charge.
Draw the symbol for the element with no dots.
Place ______________ around the structure.
Write the _______________ of the ion _______________ the brackets.
Calcium Ion
Aluminum Ion
Sodium Ion
Sulfide Ion
Bromide Ion
Practice 1: Draw the Lewis electron dot structures for these
cations.
To draw the Lewis electron dot symbol of anions:
1.
2.
3.
4.
_______________ the same number of electrons as the charge.
Draw the new electron arrangement.
Place ______________ around the structure.
Write the _______________ of the ion _______________ the brackets.
Nitride Ion
Practice 2: Draw the Lewis electron dot structures for these
anions.
IONIC BONDING
Ionic Bonds form when one or more valence electrons are _______________ from one atom to another, creating _______________
and _______________ ions.
To draw Lewis structures for ionic compounds:
1. Write the correct formula for the compound.
2. Draw the Lewis electron dot symbol for the _______________ ion(s).
3. Draw the Lewis electron dot symbol for the _______________ ion(s) to the RIGHT of the positive ion.
Example 1: Draw the Lewis structure for lithium oxide.
Practice 3: Draw the Lewis structure for the following ionic compounds:
Sodium phosphide
Magnesium nitride
COVALENT BONDING
Covalent bonding involves the _______________ of valence electrons between atoms.



One pair of _______________ electrons is represented by a _______________ dash (
)
A pair of electrons __________ shared (__________-______________) are represented by a pair of __________ around their
atom ( )
Two atoms can share more than one pair of valence electrons – Double bonds ( ) and Triple bonds ( )
How to draw Lewis Dot Structures for Covalent Compounds:
1) Determine the arrangement of atoms within a molecule. The central atom is usually the _______________ electronegative
atom. Hydrogen is a _______________ atom because it typically bonds to only one other atom.
2) Determine the total number of _______________ electrons in a molecule or ion. In a neutral molecule, this number will be
the sum of the valence electrons for each atom.
a.
b.
For an anion, __________ the number of electrons equal to the negative charge.
For a cation, ___________ the number of electrons equal to the positive charge.
3) Place one pair of electrons between each pair of _______________ atoms to form a _______________ bond. Count the
number of valence electrons in the molecule. Subtract 2 electrons from the total valence electrons for every bond you drew.
4) Use any remaining pairs as __________ pairs around each ______________ atom (except hydrogen) so that each terminal
atom is surrounded by _____ electrons. If, after this is done, there are electrons left over, assign them to the
______________ atom. (If the central atom is an element in the third or higher period, it can have more than eight electrons.)
5) If the central atom has _______________ than 8 electrons at this point, change one or more of the lone pairs on the terminal
atoms into a bonding pair between the central atom and terminal atom to form a _______________ (_______________ or
_______________) bond.
a.
As a general rule, double or triple bonds are most often encountered when both atoms are from the following list:
Carbon, Nitrogen or Oxygen
Using the steps above, draw the Lewis structures for the following covalent compounds.
Example 2: Phosphorus trichloride
Formula:__________
Example 3: Carbon monoxide
# valence e-: __________
Example 4: Silicon dioxide
Formula:__________
# valence e-: __________
Practice 4: Arsenic tribromide
# valence e-: __________
Practice 5: Carbon tetrafluoride
Formula:__________
# valence e-: __________
Formula:__________
Formula:__________
# valence e-: __________
Practice 6: Water
Formula:__________
# valence e-: __________
EXCEPTIONS to the OCTECT RULE:
1) Incomplete Octets – A central atom with __________ than _____ electrons in its outer energy level.
a. Incomplete octets are pretty rare and generally are only found in some _______________, _______________ and
_______________ compounds.
b. Boron and aluminum form compounds in which they have _____ valence electrons, rather than the usual 8 as
predicted by the octet rule.
c. Beryllium will form compounds in which it only has _____ valence electrons.
d. DO NOT double bond to satisfy their octets.
2) Expanded Octets – A central atom with __________ than _____ electrons.
a. These structures are only possible when the principle quantum number is greater than or equal to n = _____ because
their _____ orbitals are available for bonding. Some can have up to 12 electrons surrounding the central atom!
Example 5: Xenon tetrafluoride
Formula:__________
Example 6: Aluminum chloride
# valence e-: __________
Example 7: Sulfur hexafluoride
Formula:__________
# valence e-: __________
Practice 7: Iodine pentachloride
# valence e-: __________
Practice 8: Bromine pentafluoride
Formula:__________
# valence e-: __________
Formula:__________
Formula:__________
# valence e-: __________
Practice 9: Boron trifluoride
Formula:__________
# valence e-: __________
Lewis Structures for POLYATOMIC IONS:
 Use the rules for drawing Lewis structures for covalent compounds.


When determining the total number of valence electrons:
o
For an anion, _______________ the number of electrons equal to the negative charge.
o
For a cation, _______________ the number of electrons equal to the positive charge.
Enclose the entire structure in _______________ and write the _______________ of the ion outside the brackets as a
superscript.
Example 8: Ammonium Ion
Formula:__________
Example 9: Sulfate Ion
# valence e-: __________
Practice 10: Carbonate Ion
Formula:__________
Formula:__________
# valence e-: __________
Practice 11: Perchlorate Ion
# valence e-: __________
Formula:__________
# valence e-: __________
Resonance Structures – The possible structures of a molecule for which more than one Lewis structure can be written.
Example 10: Ozone
Practice 12: Carbonate Ion
Formula: O3
# valence e-: __________
Formula: CO 32-
# valence e-: __________
Valence Shell Electron Pair Repulsion (VSEPR) Theory



Based on the idea that the bond and non-bond (lone) electron pairs in the valence shell of an element _______________each
other and seek to be as far apart as possible.
It is this repulsion that causes the molecule or ion to have a particular _______________.
VSEPR Vocabulary:
1) Electron Pair Geometry – The geometry of the _______________ _______________ on the central atom. (Count
the number of atoms and non-bonding pairs of electrons around the central atom to determine the electron pair
geometry.
2) Molecular Geometry – The __________ arrangement of the ____________________ and ____________________
electrons that represent the _______________ of the molecule.
VSEPR CHART:
Example
# of atoms
bonded to the
central atom
# of lone pairs of
electrons on the
central atom
Electron Pair Geometry
Molecular Geometry
(Shape)
Draw the following Lewis structures to determine the electron pair geometry and the molecular geometry (shape) of the molecule or
ion.
Example 11: PF3
# valence e-: __________
Electron Pair Geometry: _________________________
Molecular Geometry: _________________________
Example 12: NO21-
# valence e-: __________
Electron Pair Geometry: _________________________
Molecular Geometry: _________________________
Practice 13: H2O
# valence e-: __________
Electron Pair Geometry: _________________________
Molecular Geometry: _________________________
Practice 14: BCl3
# valence e-: __________
Electron Pair Geometry: _________________________
Molecular Geometry: _________________________