Transcript Ionic Bonding

Chemical Bonds
To Bond or Not to Bond?

Why do atoms want to bond? (Hint: You
already know the answer)
To Bond or Not to Bond?
 Octet
rule: all elements want to
obtain 8 valence electrons and
become stable
Metallic Bonding

Bond between two metals

Sea of electrons

Produces an alloy
 Alloy:
solid solution of two or more metals
(ex: Steel, Brass, 14k Gold, Sterling silver,
etc.)
Ionic Bond

Electrostatic attraction: attraction between
positive and negative charge

Gains or loses valence electrons

Bond between a metal and nonmetal
Ionic Compound
Between a METAL and a NONMETAL
 Between a POSITIVE and a NEGATIVE
ion

 Ion-
a charged atom or group of atoms
 Charges between metal and nonmetal must
equal ZERO

Most are CRYSTALLINE solids
Naming Ionic Compounds
 The
metal is named just as you
see it off the periodic table
 The
nonmetal has the ending
dropped and replaced with –ide
Practice

MgCl2

KCl

Be3N

Cs2S
Writing Ionic Formulas

Label the Parts:
MgCl2

Steps:
 1.
Write the Element symbols
 2. Write each elements charge
 3. Determine subscripts by making overall
charge of the compound equal zero!
 4. Reduce if possible. All Subscripts (other
than 1) must be written
Li combines with S
Practice
K – Br
 Ca – P
 Strontium oxide
 Lithium nitride
 Sodium fluoride
 Beryllium iodide
 Magnesium oxide

Polyatomic Ions

Polyatomic ions: groups of atoms that
have a charge
 List

of Polyatomics given today
When using polyatomic ions you must use
parenthesis if there is more than 1 of
them.
Ca(OH)2
Naming Ionic Bonds with
Polyatomics

Name the metal as always. Name the
polyatomic just the way the name is on your
list.
 Ca(OH)2
 BaSO4
 Na3PO4
 (NH4)CO3
 Mg(ClO3)2
 (NH4)3N
Writing Formulas

Must protect polyatomic with parenthesis if
it is more than 1!
 Calcium
carbonate
 Lithium hydroxide
 Strontium acetate
 Potassium chromate
 Barium nitrate
 Aluminum sulfite
 Ammonium sulfide
Transition Metals

Most Transition metals and other special
metals have multiple oxidation numbers
 Exceptions:

Zn+2, Cd+2, Ag+1
Must write transition metals with there
oxidation number displayed in parenthesis
using roman numerals
 Example:
Iron (II) bromide
Roman Numerals

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
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Modern
Representation
One (I)
Two (II)
Three (III)
Four (IV)
Five (V)
Six (VI)
Old Representation
 Cupric
 Cuprous
 Ferric
 Ferrous
 Aurous
 Auric
Practice

ZnCl2

Fe3(PO4)2

CoN

Fe3N2

PbSO4

CuNO3
Practice

Cupric sulfate

Iron (II) hydroxide

Silver nitride

Lead (IV) phosphate

Iron (III) bromide

Copper (II) chlorate
Mixed

Strontium acetate

CoN

Copper (III) chromate

ZnCl2

Magnesium bromide

Na3PO4

Li2O
Covalent Bonds
Share electrons: all elements want 8
valence electrons (***except H and He want
2 valence electrons***)
 The sharing of covalent compounds gives
them different geometrical shapes
 Between a NONMETAL and a NONMETAL
 Two Types:

 Non-polar-
equal sharing of electrons
 Polar- unequal sharing of electrons
Naming Covalent Compounds
For covalent compounds both of the
nonmetals must have prefixes assigned
to represent the number of atoms
 The second element has a prefix and has
the ending changed to -ide
 One- mono
Six- hexa
 Two- di
Seven- hepta
 Three- tri
Eight- octa
 Four- tetra
Nine- nona
 Five- penta
Ten- deca

11 – undeca
 12 – dodeca
 13 – trideca
 14 – tetradeca
 15 – pentadeca
 16 – hexadeca
 17 – heptadeca
 18 – octadeca
 19 - nonadeca

Practice Naming Covalent
Compounds

CO

CO2

N2O5

NF3
Practice Writing Covalent
Compounds

Iodine pentafluoride

Nitrogen tribromide
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Diphosphorus pentoxide

Sulfur hexachloride
Diatomic Molecules

Molecule: another name for covalent
bonds

Diatomic molecule: covalent compounds
between two of the same atoms
 Seven
N2, H2
you need to know: I2, Br2, Cl2, F2, O2,
Lewis Dot for Covalent
Compounds

Lewis Dot Review: What is a Lewis dot
diagram?

Draw the Lewis dot for the following:
S
N
C
Xe
He
Lewis Dot for Covalent
Compounds
What does every element want?
 We must draw the Lewis dot diagram so every
element is SHARING 8 valence electrons
(EXCEPT: H only needs 2 valence e-)
 Covalent compounds can create:

 Single
bonds: one line drawn and represents 2 valence
electrons
 Double bonds: two lines drawn and represents 4 valence
electrons
 Triple bonds: three lines drawn and represents 6 valence
electrons
Examples:
VSEPR Theory

VSEPR: Valence Shell Electron Pair
Repulsion Theory

Definition: in covalent compounds the
geometric arrangement is determined solely
by the repulsions between electron pairs
present in the valence electron shell
 To
determine shape you must first draw the
Lewis dot diagram
Polar vs. Nonpolar
Polar
Definition: unequal
sharing of electrons


 In
a polar molecule
there will be at least
one free pair for the
central atom
Nonpolar
Definition: equal
sharing of electrons


 Nonpolar
molecule
has 0 free pairs from
the central atom
Ionic vs. Covalent



Ionic
Very strong bonds
Melt at very high
temperatures
Crystal structures
called salts and many
can dissolve in water


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Covalent
Weaker bonds
Melt at very low
temperatures
Most cannot dissolve
in water
Determine Type of Bond

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K – Br
S–O
Si – Cl
H–F
Se – S
H–O
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Na – Cl
Fe – S
H–N
Ca – I
Al – O
Mg - Br