Chemical Bonding
Chemical Bonds:

Elements form bonds to be in a ___________ energy state
1. Ionic Bonds – _____________ of electrons, between ________ and _____________
2. Covalent Bonds – ___________ of electrons, between two nonmetals
3. Metallic Bonds- _______________ atoms in ___________ metals form bonds
Octet rule: atoms tend to ________, _______, or __________ electrons until they are surrounded by
___________ valence electrons to achieve a stable octet (noble gas configuration)
Electron Dot Symbols

______________ electrons: reside in the ____________ occupied energy level, reside in the outer s & p
orbitals and are the electrons involved in chemical ________________.

Electron-dot symbols are convenient way of showing the s & p electrons & tracking them in bond formation

They consist of the chemical _____________ for the element plus a dot for ________ valence electron

Consider sulfur whose electron configuration is [Ne]3s23p4, thus there are
valence electrons.
S
Ionic Bonds –
•
atoms ________________ electrons from a ______________ (positive ion) to an
______________ (negative ion) to achieve an octet.
•
Ionic compounds are ___________ due to the ______________ ____________ between unlike
charges organizing the ions of ionic substances into a rigid, organized three-dimensional
arrangement:
•
The ions are ____________together
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Energy is ________________
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Ions form solid _________________ structure
Ionic Rule 1: Metals with a Single Oxidation Number Bound to Non-Metals
 The metal will take its _________________ oxidation number and the non-metal will have to take its
__________________ oxidation number.
 Only one compound can be formed
Basic Formula: ___________________________________________

Example 1: Sodium reacts with oxygen to produce Na2O, what is the name of this compound
 Since there is only one possible compound, we do not have to indicate the number of elements
 Answer: _________________________________

Example 2: What is the chemical formula for aluminum oxide
 First write the ________________ of the elements
 Next write the oxidation number of each element _____________ that element
 Switch the oxidation numbers and _________________
 Answer: _____________________________________________
Comprehension Check:

What is the name of Mg3N2?
What is the formula for indium chloride?

What is the name of Li2Se?
What is the formula for hafnium phosphide?
Ionic Rule 2: Metals with Multiple Oxidation Numbers Bound to Non-Metals
The metal will take one of its positive oxidation numbers and the non-metal will have to take its negative oxidation
number.
o Since the metal has more than one possible oxidation number, ______________ compounds can be formed
o We need a ____________________name for each
Formula: _________________________________________________________________________________
Example 1: What is the name of IrBr6?
o First we need to determine how many _______________ that iridium needs to lose in order to satisfy 6
bromine atoms.
 Each bromine needs one ____________________
 There is only one iridium in this compound
 Therefore, the iridium atom will have to supply all ____ electrons, giving it a +6 oxidation number.
Answer: _________________________________________________________
Example 2: What is the formula for mercury(II) nitride?
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


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First write the symbols of the elements
Next write the oxidation number of each element _______________ that element
__________________ the oxidation numbers and reduce
Answer: ________________________________________________________
Comprehension Check:
What is the name of RuN?
What is the formula for paladium(IV) bromide?
What is the name of MnO3
What is the formula for molybdenum(V) sulfide?
Ionic Rule 3: Metals with a single Oxidation Number Bound to Polyatomic Ions

o
o
o
Polyatomic Ions – strongly bound group of atoms that have either lost or gained electrons and become charged.
List of common Polyatomic Ions are on the back of your Periodic Table
Polyatomic ions act as a ________________ atom, with a single name
Subscripts within the ion ____________________ be changed
o
Since there is only one oxidation number for the metals and Polyatomic Ion, only _____ compound can be
produced.
 Naming these compounds is just like rule 1, except we do not add –ide to the end of the polyatomic ion
Basic Formula: _________________________________________________________________
Example 1: What is the name of Mg(NO3)2
Recognize that there are more than _____ elements involved, which means that a Polyatomic Ion is involved
Next, look up the ______________ in the periodic table and confirm that it has a single oxidation number
Look up the name of the ____________________ __________
Answer: ______________________________________________________________
Example 2: What is the formula for calcium iodite?
 First, since the second name does not end in __________, a polyatomic ion is involved.
 Write the symbol for calcium and formula for iodite.
 Write the ______________ _______________ above the metal and the polyatomic ion
 Switch the numbers, and use ____________________ around the polyatomic ion if necessary
Answer: ____________________________________________________________
Comprehension Check:
What is the name of KHSO4?
What is the formula of strontium bromate?
What is the name of In2(C2O4)3?
What is the formula for germanium phosphate?
Ionic Rule 4: Metals with Multiple Oxidation Numbers Bound to Polyatomic Ions
When the metal has ___________ than one possible oxidation number, more than one compound can be formed
 We must use _________ ______________ to indicate which oxidation number the metal is using
Basic Formula: ______________________________________
Example 1 : What is the name of RhSO4?
 First, there are more than __________ elements involved
 Look up the oxidation and name of SO4
 Finally, figure out which oxidation number the metal is using.
There is only one rhodium, so it must account for ____________ of the electrons & would have to take
a ________ oxidation number
Answer: __________________________________
Example 2: What is the formula for nickel (II) ferrocyanide?
 First, since the second name _________ end in -ide, is a polyatomic ion is involved? It is an _____________
 Write the ________________________ for nickel and formula for ferrocyanide
 Write the oxidation numbers above the metal and the polyatomic ion
 Switch the _______________, and use parenthesis around the polyatomic ion if necessary & ___________
Answer: _________________________________________________________
Comprehension Check:
What is the name of Cr(IO)3?
What is the formula for palladium(IV) ferricyanide?
What is the name of CuMnO4?
What is the formula for molybdenum(VI) dichormate?
Properties of Ionic Compounds:
•
•
Arranged in repeating three-dimensional ________________
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Large ________________ forces result in very _________________ structures
•
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Ionic compounds can ____________ an electric current when melted or dissolved in water
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Bonding in Metals
 Metallic bonds: _______________________________________________________
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Bonding is due to ___________ _____________ which are _________________
throughout the entire solid
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The metal is held ____________________ by the strong forces of ________________
between the positive _______________ and the delocalized ___________________.
Metals:
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Metals are good _________________ of ________ and ________________because the
_______________ electrons are able to flow _____________
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Valence electrons of metals can be thought of as a ______ of ___________ very ___________
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Metallic Bonding - “Electron Sea”
Metallic Properties:
•
•
Have _________________________________________________________________
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Luster = ___________
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Ductile = ______________________________________________________
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Malleable = ____________________________________________________
Properties can be explained by the ___________________ of electrons in metals
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When subjected to pressure , the _____________ easily slide past each other like a ball
bearing immersed in oil.
Covalent Bonding
What is a covalent bond?

A chemical bond that results from the ________________ of electrons, to form a stable ______
or __________ (Hydrogen only needs 2 to be stable)

Molecule = ______ or more atoms that are held together by _______________ bonds

Majority of covalent bonds form between _______________ (CLOSE together on periodic table)
Covalent Rule 1: Nonmetals Bound to Nonmetals
Since nonmetals have more than one oxidation number, there will always be more than one compound produced
 Therefore we have to have a distinct name for each compound
 To do this we use a prefix to indicate how many atoms of each element are present
One :______________
Two: ______________
Three: _____________
Four: _______________
Five: _______________
Six: _________________
Seven: ______________
Eight: ________________
Using prefixes
 The prefix mono- is only used on the second element
Ex: PF3 is named ____________________________
 If two vowels are adjacent, leave them
Ex: NI3 is named ______________________________
 In the case of monoxide only, drop one “o”
Nine: ______________
Ten:_______________
Ex 1: What is the name of P2S3?
________________________________
Ex 2: What is the name of As7I3?
________________________________
Ex 3: What is the chemical formula of dihydrogen monoxide?
________________________________
Ex 4: What is the chemical formula of dinitrogen pentaoxide?
________________________________
Diatomic molecule
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molecule containing the ______ _____ atoms
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Some elements _________ exist this way because they are _____ stable than the individual atoms

The Seven Diatomic Elements :___________________________________
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Bonds in _____ the ________________ ions and diatomics are all covalent bonds
Single Covalent Bonds
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Two atoms held together by a sharing of _____ pair of
electrons are joined together by a single covalent bond.
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An electron dot structure represents the __________
pair of electrons of the covalent bond by ______ dots.
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A ______________ formula represents the covalent bonds by _________ and shows the arrangement
of covalently bonded atoms.
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A pair of valence electrons that is ____ ___________ between atoms is called an unshared pair, also
known as a ______ pair of a _________________ pair.
Double & Triple Covalent Bonds
o
o
o
Atoms form _______ or ____________ covalent bonds if they can attain a noble gas structure by
sharing two or three pairs of electrons.
A double bond involves _________ ________ pairs of electrons.
A __________ bond involves sharing ___________ pairs of electrons
Lewis Diagrams
1. Arrange atoms
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___________ atom is usually in the center (often Carbon)
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If not Carbon, _______________ electronegative atom is in center
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Hydrogen is always _____________ (on the side, not a central atom)
2. Find ___________ # of e- available to bond (_____________ e-! only )
3. Place a __________ of electrons between _____________ atom and each terminal atom
4. Place remaining electrons in pairs around _____________ atoms (except H) to satisfy octet rule
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Any ______________ pairs are assigned to central atom
5. Determine whether or not central atom satisfies _____________
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If not, __________ one or more lone pairs from terminal atoms to double or triple bonds
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Certain atoms can be ___________ to octet rule – H, Be, B, S, P, Xe
Ex 1 : CF4
Ex 2: CO2
Polyatomic Ions:
 To find ___________ # of valence e-:
 _________ 1e- for each negative charge
 ____________ 1e- for each positive charge
 Place ______________ around the ion and label the charge
Ex 3: ClO4-1
Octet Rule Exceptions:

______________  2 valence e-

Boron & Beryllium get ____ & ____ valence e- respectively

_____________ octet  more than 8 valence e- (e.g. S, P, Xe)
Resonance Structures

Molecules that ____ be correctly represented by a _________ Lewis diagram

Actual structure is an _____________ of all the possibilities

Show all possible structures separated by double-headed ___________
Ex: SO3
Bond Polarity
-Most bonds are a blend of ________ & __________ characteristics.

____________ in electronegativity determines bond type.

Electronegativity
o ______________ an atom has for a shared pair of electrons.
o
___________ e-neg atom  -
o
___________ e-neg atom 
+
Electronegativity Difference

If the difference in electronegativities is between:

1.7 to 4.0: ___________

Greater than 0.3 & less
than 1.7: _____ Covalent

0 to 0.3: __________
________________
The type of bond can usually be calculated by finding the difference in electronegativity of the two atoms
that are bonded.
o
o
o
Compound:
Difference:
Type of Bond:
F2 or F-F
_______
_______
o
o
o
Nonpolar Covalent Bond
e- are shared _____________
________________ e- ______________
usually _____________ atoms
 Ex: H2 or Cl2
CF4
_______
_______
o
o
o
LiF or Li - F
_______
_______
Polar Covalent Bond
e- are shared ________________
_____________ e- density
results in ________ ___________ (dipole)
 Ex: H2O
VSEPR THEORY
o
o
________________________________________________________
Electron pairs ____________ themselves in order to _______________ repulsive forces
Types of e- Pairs
o __________________ – form bonds
o Lone pairs – _______________ eo Total e- pairs– _________ + _______ pairs
 Lone pairs ________ more ____________ than bonding pairs!!!
 Lone pairs __________ the bond angle between atoms
Determining Molecular Shape
o
o
o
o
Draw the ____________ Diagram
Tally up e- pairs on central atom (bonds + lone pairs)
double/triple bonds = ONE pair
Shape is determined by the _________ of bonding pairs and lone pairs
Molecular Polarity:
 ___________ molecule = one end slightly __________ and one end slightly __________
 Molecule with 2 poles = ___________ molecule or ___________
 _____________, symmetry and bond polarity determines molecular polarity
 H – O bond is __________ and water is ______________, so H2O is polar
 C – Cl bond is polar, but CCl4 is _________________, so molecule is _____________
Identify each molecule as polar or nonpolar
o
o
o
o
O2 : Nonpolar Bonds __________
CS2 : Linear  ______________
CF4 : Tetrahedral  __________
H2O : Bent  ____________
Intermolecular Forces
Intramolecular and Intermolecular Forces:
 Covalent & Ionic bonds - ______________________________________________________
•
These are ___________________ forces.
 There are also forces that cause molecules to ___________ each other. These are called
______________________ forces.
Definition of Intermolecular Forces:
 Attractive forces _________________ molecules.
 Much ____________than chemical bonds
_______________ molecules.
a.k.a. __________________________________________________
Types of IMF: