General Chemistry
Name:
Unit 4 Note Packet – Compound Formation, Nomenclature, and Bonding
_
Period: _______________
Compounds
Compounds are ___________ made up of ______________ elements in ______ proportions.

electrically neutral ( _____ numbers of
___________________________)

atoms combine by ________________________
electrons to _____ chemical bonds

atoms achieve greater __________ in ___________
with other atoms
3 Types of Bonding

________________
-
( _______________ +
_______________ )

________________
-
( _______________ +
_______________ )

________________
-
( _______________ +
_______________ )
METALLIC BONDING AND COMPOUNDS
Metallic Bonds
Characteristics of Metallic Compounds

electrons are _________


bond is an _____________ (positive-netagive)

conduct ___________ and heat

usually _______ at _______ temperatures
attraction between _______ and a “sea” of
_____________________________
+
(range of m.p./b.p.)

______________ in water

luster is _________

____________________ and
__________________
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Metallic Bonding and Compounds
Alloy: a ____________________ (usually a _____________) that contains
____________________________________ and has the __________________________ of a metal
Compositions of Selected Alloys
stainless steel: 74% ____, 18% ____,
brass: 67% ____, 33% ____
8%____0.18%___
18 carat gold: 75% ____, 10-20% ____, 5-15% ____
coinage silver: 90% ____, 10% ____
Nichrome: 60% ____, 40% ____
plumber’s solder: 67% ____, 33% ____
*note that the elements in an alloy are not present in specific ratios (the percentages may be
adjusted) and so alloys do not technically qualify as being true compounds
IONIC BONDING AND COMPOUNDS
Ionic Bonds

electrons are _______________
Characteristics of Ionic Compounds

bond is an ________________ attraction

__________ + _____________

called _________

_______ at _______ temperatures (high m.p.
between ________ and an _________ (
______ )
and b.p.)

atoms are often ____________ than ions

metals ______ electrons to form __________

_________________________
ions to achieve stability 

example:

______

___________

______________, NOT molecules

example:
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conduct ___________ when _________ (
_______ at high temperature)
non-metals ______ electrons to form
___________ ions stability 
may dissolve in _______ to form
they form ____________________ of
________________ anions and cations
2
COVALENT BONDING AND COMPOUNDS
Covalent Bonds
Characteristics of Compounds

electrons are ___________

__________ + ______________

_________________ attraction between

___________________ at ______
____________ and __________
temperatures ( _____________ m.p. and b.p.)

may dissolve in ________ but doesn’t form an
_______________

doesn’t ____________________ when
______ or molten ( _______ at high
temperature)

forms _____________
Describing Covalent Bonds
unshared pair - ______________________ that are not ____________ in bonds
single bond – only ______________ of _________ are _________ between two atoms (see examples above)
double bond – ___________ of ___________ are __________ between two atoms
examples:
1)
2)
triple bond – _______________ of ___________ are __________ between two atoms
examples:
1)
2)
Classify each compound as:
M- Metallic
I-Ionic
C- Covalent
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•
•
•
•
•
KCl
Brass (Cu + Zn + Sn)
CO2
NO2
Sterling silver (Ag + Cu)
3
•
•
•
•
•
SnF2
CH4
MgCl2
NH3
Li
CHEMICAL FORMULAS
Chemical formula: what _________________it contains and the __________________________ of those
elements
– Example:
Contains ___ sodium atom and ___ chloride atom
–
Example:
The formula is ______________________ of the symbols
_____________________________________
Contains ___ Hydrogen atoms and ___ Oxygen atom
Subscript means “___________________________” and is written ________________ the symbol. It tells
how many _______________ of that __________________ are in one unit of the compound. If the symbol
has _________________________, the unit contains _____________________ atom of that element.
Familiar Name
Chemical Name
Formula
Lye
Ammonia
Sand
Battery Acid
Octet Rule: atoms tend to ____________________________ electrons in order to acquire a
______________ of ______________ electrons

Think of ionic bond formation as a process:
– electrons are _______________ to achieve a ____________________ of electrons
– __________ form
– ions brought together by _________________________________________
Lewis Dot Diagrams:
 Recall that a way to show and emphasize an atom’s valence electrons is to draw the element’s dot
diagram
Li
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Be
N
O
4
Cl
Si
Ar

Dot diagrams can also be used to represent how electrons are rearranged during chemical reactions as
compounds are formed.
o Example: lithium oxide (Li2O) forming from lithium and oxygen
o Use dot diagrams to show how beryllium chloride (BeCl2) would form from beryllium and chlorine
How atoms combine
Fill out the following table for each element listed. Try to bond two or three different atoms by giving or
taking away electrons, thus forming a compound. Draw the Lewis dot diagram of each compound. Create a list of
formulas and the Lewis dot diagrams of the compounds that you are able to form below. Your list should contain
at least 10 compounds!!
Element
Element
# of
# of
# of valence
# of electrons to
Lewis dot
name
symbol
protons
neutrons
electrons
Sodium
Fluorine
Oxygen
Bromine
Chlorine
Magnesium
Lithium
Hydrogen
Beryllium
Silicon
Potassium
Phosphorus
Compounds formed:
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be gained or lost
diagram
Naming an Ionic Compound from a Formula
1. identify the compound as ionic
2. determine the name of the cation and the name of the anion
3. if the cation is a transition metal, determine the charge on the metal
(if a polyatomic ion is included, look up the name)
4. write the name with the cation first and then the anion with the “-ide” ending second
(include the charge of a transition metal after its symbol (Roman numeral in parentheses))
Practice:
K 2O
CaCl2
LiCO3
(hint: look up the name for CO3-)
FeCl 3
(hint: figure out the charge on the iron ion)
Writing an Ionic Formula from a Name
1. identify the compound as ionic
2. determine the symbol of the cation and then the symbol of the anion
3. determine the charge on each ion
4. balance the overall charge by combining the ions in the proper proportions (criss-cross)
5. write the cation symbol and then the anion symbol; use subscripts to denote more than one ion
6. if there is more than one polyatomic ion, put the ion in parentheses
Practice:
aluminum chloride

sodium iodide

potassium sulfate

titanium (IV) oxide

There are two types of ions:
1.
monatomic ion:

most monatomic cations have the same name as the element
Na+ is called
Mg2+ is called

if the cation is a transition metal with more than one oxidation state, a Roman numeral is added to
specify the charge
Fe2+ is called iron (II) [this is read as “
”]
Cu+ is called copper (I) [read as “

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”]
monatomic anions have the same name as the element but the suffix is replaced with “-ide”
Cl- is called
instead of chlorine ion
O2- is called
ion instead of oxygen ion
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2.
polyatomic ion:
(electrons
gained/lost are distributed throughout the group)
most polyatomic ions have names that end in “-ate” or “-ite”

SO42- is the sulfate ion
PO43- is the phosphate ion
see table
in your reference tables for more polyatomic ions
Hydrates:
the water molecules in the crystal structure are called “waters of hydration”
the number of water molecules in the substance is indicated with a prefix before “-hydrate”
for example: MgSO4 • 7H2O is



Naming and Formulas for Covalent Compounds
Covalent compounds are named using prefixes to indicate the
number of each type of atom.
1. identify the compound as covalent
2. determine the name of the first element in the compound
and indicate how many atoms of this element are present by
using the appropriate prefix
3. determine the number and name of the second element and
change the suffix to end with “-ide” (note that this is the
more electronegative element); use the most appropriate
prefix
Exceptions:
 the prefix mono- is not written with the first word of a
compound’s name (ex:
)
 prefixes are sometimes shortened to make the name easier
to say (ex:
)
 diatomic elements are called by the element name
(ex:
)
 common names are used for some common substances
(ex:
)
Practice:
NO2
BF3
nitrogen trifluoride
diphosphorous pentoxide
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Stability of Atoms and the Octet Rule in Covalent Compounds
 nonmetals in covalent compounds achieve stability by sharing electrons to obtain an octet
Dot structures can also be used to show how covalent bonding occurs. Example: silicon dioxide (SiO 2)
Use dot diagrams to show how ammonia (NH3) would form.
Structural formulas for compounds sometimes show pairs of electrons as lines instead of using two dots. Count
the total number of valence electrons in the compound and then distribute the pairs so that each atom has a
full set of valence electrons. Example: phosphorous trichloride (PCl 3)
Use a structural formula with lines to show a molecule of water.
Exceptions to the Octet Rule
There are several compounds that do not follow the octet rule.
 atoms with less than an octet of electrons
hydrogen is stable with only two valence electrons (H2)
boron is stable with only six valence electrons (BF3)

atoms with more than an octet - sometimes additional electrons fill the 3d orbitals of atoms beyond the
second period of the periodic table
sulfur sometimes has ten electrons involved in bonding (SF4)

molecules with an odd number of electrons
these tend to be short-lived because they are unstable (NO)
Empirical vs. Molecular Formulas
empirical formula
 shows the _____________________________________________ of _____________ in a
compound
 _____________ used for ______________ compounds
 can be useful for _____________________ describing _________________ compounds
 example:
molecular formula
 shows the ______________________ of ____________ in _______________ molecule
 ______________ be used for ____________ compounds
 examples:
Questions:
What is the empirical formula for sucrose?
What is the empirical formula for glucose?
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Empirical vs. Molecular Formulas
empirical formula

 always used for ionic compounds
 can be useful for partially describing covalent compounds
 example: Ca2+ and F- combine to form CaF2
molecular formula

 cannot be used for ionic compounds
 examples: sucrose- C12H22O11 and glucose- C6H12O6
Questions:
What is the empirical formula for sucrose?
What is the empirical formula for glucose?
Polarity of Covalent Bonds
 electrons are not always shared equally in a covalent bond
 different atoms share more or less equally depending only their electronegativity (def: an atom’s
attraction for electrons in a chemical bond)
nonpolar bond: a bond in which
(less than 0.4)
polar bond: a bond in which
(between 0.5 and 1.9)
ionic bond: note that an ionic bond can be thought of as a situation where electrons are shared so unequally
that an electron is essentially given/taken
Polarity in bonds is shown with the delta symbol and arrows that have a
cross at one end. Notice that polar bonds in a molecule can give
rise to an overall charge imbalance for the whole molecule.
Such an imbalance leads to a polar molecule.
Bond Type by Electronegativity
Electronegativity Difference
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Bond Type
nonpolar covalent
polar covalent
ionic
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MOLECULAR GEOMETRY
Electrons repel one another, so when atoms come together to form a molecule, the molecule will assume
the shape that keeps its different electron pairs as far apart as possible. When we predict the geometries of
molecules using this idea, we are using the
.
In a molecule with more than two atoms, the shape of the molecule is determined by the number of
electron domains on the central atom. The central atom forms hybrid orbitals, each of which has a standard
shape. Variations on the standard shape occur depending on the number of bonding pairs and lone pairs of
electrons on the central atom.
Here are some things you should remember about DOMAINS when dealing with the VSEPR modelA DOMAIN is a group of electrons, including both bonding pairs and lone pairs of electrons. In using
the VSEPR model, we look at the DOMAINS around the central atom in a covalent compound.
 Single, double and triple bonds are all DOMAINS. Double and triple bonds are treated in the same way
as single bonds in predicting overall geometry for a molecule, but multiple bonds have slight more
repulsive strength and therefore will occupy a little more space then single bonds.
 A lone pair of electrons is a DOMAIN as well. Lone electron pairs have a little more repulsive strength
then the bonding pairs, so lone pairs will occupy a little more space than the bonding pairs.
The tables below show the different geometries:

If the central atom has 2
electron domains, then its
basic shape is linear.
If the central atom has 3
electron domains, then its
basic shape is trigonal planar.
If the central atom has 4
electron domains, then its
basic shape is tetrahedral.
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POLARITY (of molecules)
 recall that polar bonds are bonds in which electrons are not shared equally
 polar bonds can give rise to polar molecules, or dipoles that have positive and negative ends
 polarity of molecules gives a molecule a variety of different properties
Determining polarity
nonpolar bonds  nonpolar molecule
polar bonds  polarity of the molecule depends on the polarity of its bonds AND the shape of the molecule (in
other words, a molecule that contains polar bonds is not necessarily polar)
Examples:
oxygen (O2)
formaldehyde (H2CO)
carbon dioxide (CO2)
INTRAMOLECULAR FORCES
-forces within a molecule
-tend to be strong
vs.
INTERMOLECULAR FORCES
-forces between molecules
-tend to be weaker than
Intermolecular forces
Matter has entropy (the tendency to be disordered) so a force must be present to keep individual atoms,
molecules, or ions in a solid or liquid in place. These forces are… INTERMOLECULAR FORCES.
Van der Waals forces:
1. London/dispersion
2. Dipole-dipole forces
3. Hydrogen bonds
4. Molecule-ion attractions
1. Dispersion Forces
 Due to attractive forces between e-s of one atom and the nucleus of another
 occurs btw molecules that are
 moving e-s create a temporary dipole
 temporary dipole can induce dipoles in nearby molecules
 stronger for atoms/molecules with more electrons
2. Dipole-dipole forces
 due to attractive forces between the positive end of one molecule and the negative end of another
 occurs between
molecules
3. Hydrogen Bonding
 a special dipole-dipole attraction between
very high electronegativities and small radii (
 explains high boiling point of water
and three other elements with
)
4. Molecule-ion attractions
 ionic compounds dissolve in water and other polar liquids because of attraction between the dipoles
and the ions
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