Transcript Back to School Night

Bonding
Ionic bond (formula units)
 Between metal and a nonmetal
 Transfer electrons
 Covalent bond (molecules)
 Between 2 nonmetals
 Share valence electrons
 Alloy (metallic “bond”)
 Two metals just mix
 Don’t chemically bond or react
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Properties of Ionic Compounds
 Ionic compounds exist as crystalline solids.
 A crystal is a regular, repeating, three-dimensional arrangement
of positive and negative ions known as a crystal lattice.
 Held together by strong electrostatic forces (opposites attract).
 Identifiable properties:
• very high melting points
• hard but brittle (shatters when hammered)
• ions cannot move in the solid state
• when dissolved in water or melted to liquid state, ions
dissociate and form electrolytes
Characteristics of the Covalent Bond
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The forces of attraction are much weaker than
those of ionic bonds.
Molecules melt at low temperatures.
Can not conduct electricity in solution.
Diatomic molecules
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Elements with strong electronegativities that bond with
themselves
“Super 7”
H2 , N2 , O2 , F2 , Cl2 , Br2 , I2
“Super 7”
Ionic Formulas
A general rule is to “criss-cross” the absolute value of the charges to
balance out the transfer of electrons
example: aluminum bound to oxygen
1) Write the symbols for the ions.
Al3+ O2–
2) Cross over the numerical value of each ion’s charge (not the
charge itself) to the opposite element to form the subscripts
Al2 O3
3) Check the combined positive and negative charges to see if
they are equal. The overall net charge on an ionic compound is
equal to zero.
(2)(+3) + (3)(-2) = 0
Naming Ionic Compounds
(with representative elements)
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Binary Compound (only contains a total of two
types of elements)
a) Cation (metal) is identified simply by the element’s
name off the periodic table
b) Anion (non-metal) is named using the root name of the
element with an –ide ending.
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Naming Ionic Compounds
(with representative elements)
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Compound containing a polyatomic acid (more
than two types of elements)
a) Cation (metal) is identified simply by the element’s
name off the periodic table
(exception: NH4+1 = ammonium)
b) Anion is named using the name of the polyatomic ion
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Naming Ionic Compounds
(with transition metals)
a) Cation (metal) is identified simply by the element’s name
off the periodic table
*
if the metal can have multiple charges (those metals that
have “d” electrons available to give away), a Roman numeral is
added after the cation’s name to indicate the ion’s
charge
• examples:
Fe2+
iron(II)
Fe3+
iron(III)
b) Anion (non-metal) is named using the root name of the
element with an –ide ending or the name of the
polyatomic ion
Acids
Compounds that contain Hydrogen ion (H+) when
dissolved in water (aqueous solution = aq)
Two categories:
Binary acids: hydrogen bound to one other element
Named:
a) use the prefix “hydro-”
b) take the root name of the 2nd element and add –ic
ending
c) add the word “acid” to the end
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ex.
HCl Hydrochloric acid
H2S Hydrosulfuric acid
HF Hydrofluoric acid
Oxyacids: hydrogen bound to a polyatomic ion
NO “hydro”
a) name is based on the name of the polyatomic ion
• if it ends in –ate change to –ic ending
• if it ends in –ite change to –ous ending
b) add the word “acid” to the end
ex.
HNO2
H2CO3
HClO3
H2SO3
nitrous acid
carbonic acid
chloric acid
sulfurous acid
Covalent Compounds (molecules)
Bond between 2 non-metals that share their valence electrons
Naming binary covalent compounds (molecules):
a) Name the 1st element in the formula directly off the periodic
table
b) Name the 2nd element using the root name off the periodic
table with an –ide ending
c) Prefixes are used to indicate the number of atoms of each
element that are present in the compound *
1 – mono**
2 – di
3 – tri
4 – tetra
5 – penta
6 – hexa
7 – hepta
8 – octa
* mono- is never used for the 1st element
9 – nona
10 – deca
Covalent Bonds
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The forces of attraction are much weaker than those of
ionic bonds.
Molecules melt at low temperatures.
Can not conduct electricity in solution.
Some have the ability to share more than one pair of
electrons, forming multiple bonds.
 Double bond: two pair of electrons (4 total) are shared
between the two atoms (ex. O2)
 Triple bond: three pair of electrons (6 total) are shared
between the two atoms (ex. N2)
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The distance between the nuclei of two bonded atoms is
called the bond length.
Energy is released (exothermic) when a bond forms and is
absorbed (endothermic) when a bond breaks
The amount of energy required to break a covalent bond is
called the bond dissociation energy.
The stronger the bond, the greater the bond dissociation
energy and, therefore, the more difficult it is to break the
bond.
Shorter bonds have greater bond dissociation energies than
longer bonds.
Single bonds < double bonds < triple bonds
Lewis Structures: Uses electron dots to show the
arrangement of electrons in a molecule
Steps:
1. Predict the location of the atoms
a)
b)
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Hydrogen is always a terminal atom
Center atom is least electronegative element
Carbon (if present) is always a center atom
2. Count the total number of valence electrons in the elements to
be combined
3. Determine the number of “pairs” of electrons in the molecule
by dividing the total number of valence electrons by 2
4. Place a single line (“bonding pair”) between the center atom
and the terminal atoms
5. Subtract the number of pairs used from the total number of
pairs available
6. Starting with the terminal atoms, add unshared pairs so that
each atom is surrounded by eight electrons (remember
hydrogen only shares one pair)
7. If the center atom does not have an octet, one or two of the
lone pairs around the terminal atoms must be converted to
form multiple bonds
In general, carbon, nitrogen, oxygen and sulfur can form double or triple
bonds.
Resonance Structure:
• Occurs when more than one valid Lewis Structure can be
written for a molecule or ion.
• Only differ in the position of electron pairs, never in the
atoms position.
ex.
O3 (ozone)
NO2-1 (nitrate)
CO3-2 (carbonate)
Exceptions to the Octet Rule
1)
Fewer than eight electrons around the atom (hydrogen and boron
containing compounds such as BH3)
2)
Odd number of total valence electrons (These compounds usually
form polyatomic ions to “make-up” the difference)
ex. ClO2
2)
Expanded Octets: central atom contains more than 8 electrons
Usually occur with non-metals beyond period 3 when bound to
highly electronegative elements fluorine, oxygen, and chlorine.
ex. SF6
Molecular Geometry (Shape)
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VSEPR Theory: “valence-shell electron-pair repulsion.”
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The electron pairs are oriented as far away from each other as possible
to minimize the repulsion around the center atom.
The shape of a molecule refers to the positions of atoms only.
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5 main shapes (based on the octet rule)
1) Linear
2) Bent
3) Trigonal planar
4) Trigonal pyramid
5) Tetrahedral
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6) Trigonal bipyramid
7) Octahedral
Electronegativity and Polarity
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Electronegativity: a measure of the tendency of an atom to
attract electrons in a chemical bond
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Polarity: the uneven distribution of electrons (molecule is
asymmetrical around the center atom)
For polar covalent bonds, a dipole is established.
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The forces of attraction within a compound are known as
intramolecular forces. (holds together the atoms making
up a compound)
• Ionic
• Covalent
• Metallic
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The forces of attraction between molecules are known as
intermolecular forces.
1) London Dispersion Forces
 Very weak forces of attraction between non-polar
molecules
 Result from the temporary dipole occurring as
molecules approach one another
 The more electrons that are present, the stronger the
dispersion forces will be.
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Polar molecules have dipoles (partial positive and partial
negative regions.)
2) Dipole-dipole Forces
 Occur between polar molecules,
 The partial negative region in one molecule attracts the
partial positive region in a neighboring molecules.
There is an electrostatic attraction between the
molecules.
3) Hydrogen Bonding
 Hydrogen bound to an atom that has lone pairs of
electrons
 The hydrogen atom is attracted to an unshared pair of
electrons on a neighboring molecule.
 Are the strongest intermolecular force