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Ionic- Stealing
Covalent- Sharing is caring
Metallic- Sea of Electrons
Ionic Bond – chemical bond resulting from the
electrostatic attraction between positive and
negative ions
How to determine:
*Look at electronegativity to determine the type of
bond.
*Difference in electronegativity of 3.3 to 1.7 is
considered an ionic bond.
Pure ionic bond – one atom has completely given
up one or more electrons; another atom has
gained them. (electrons are transferred)
**The bond is the opposite charge attraction that
occurs after each element has reached noble
gas stability **
Covalent bonds – chemical bonds resulting from
the sharing of electrons between two atoms
How to determine:
Difference in electronegativity
1.7 – 0.3 – 0.0
Polar covalent bond – united atoms have an
unequal attraction for shared electrons
Nonpolar covalent bond – bonding electrons are
shared equally by the bonded atoms, resulting in
a balanced distribution of electrical charge
Molecule – a group of two or more atoms held
together by covalent bonds and able to exist
independently
Diatomic molecule - molecule containing two
atoms (H2 , HCl)
Chemical formula – shorthand representation of
the composition of a substance using atomic
symbols and numerical subscripts.
A.K.A- Molecular formula
Ionic compound – composed of positive and
negative ions combined so that the positive and
negative charges are equal.
i.e. The positive and negative must be the same for
the compound to be stable
Other Ionic Compound Need-To-Know’s
Formula Unit – simplest unit indicated by the
formula of any ionic compound.
Formation of crystal structure involves many
ions. [orderly array of ions called a crystal
lattice].
Lattice energy – the energy released when one
mole of an ionic crystalline compound is formed
from gaseous ions.
Properties of Ionic compounds
+- attraction causes the macroscopic properties
-Relatively high melting points
- Brittle solids
-Conduct electricity in the molten state
(free charged ions)
Metallic Bonding
-Bonding that occurs within the atoms of a metal is different from ionic and covalent bonding.
-*It is a way that the atoms can bond with themselves*
-Metals can eject their two valence electrons becoming a positively charged ion but stay bonded through the localization of electrons
Metallic bond – a chemical bond resulting from the attraction between positive ions and mobile electrons.
Metallic bonding accounts for many properties of metals.
Examples:
Good conductors of electricity
Good conductors of heat
[free movement of electrons]
Luster-absorbs and re-emits light due to the many energies of electrons
Malleability (hammered into shape)
Ductility (drawn into a wire)
[metallic bonding is not directional]
What is an ionic bond?
What types of elements are involved in an ionic
bond?
Are Ionic bonds strong or weak?
(Hint: Melting Point)
What are the names of the following ionic
compounds
NaCl, MgCl2, Al2O3,Na3N
Covalent bonding- formation of a molecule that
requires the sharing of electrons between two or
more atoms.
Single bond – covalent bonds produced by the
sharing of one pair of electrons between two
atoms.
Double bond – covalent bond produced by the
sharing of two electron pairs
Triple bond – covalent bond produced by the
sharing of three electron pair
Unshared pair – a pair of electrons that is not involved in the bonding. These electrons belong exclusively to one atom.
Lewis structures:
Shows valence electrons
Structural formula:
Shows the bonds between the atoms
**Same as Ionic but the numbers matter now!
Step 1: Leave name of first element alone
Step 2: Change nane of second element to end
in –ide
Step 3: Add prefixes to signify the subscripts
Mono=1
Di=2
Tri=3
Step 1: Add up all of the valence electrons of EVERY atom in the chemical formula E.g. CH4 (C= 4, H=1, H=1, H=1, H=1;)
4+1+1+1+1= 8
Step 2: Determine a central atom (generally the first atom written in the compound)
**Never Hydrogen**
Step 3: Use 2 electrons or a – line to create single bonds between atoms.
Step 4: Add extra electrons as lone pairs to unstable atoms if needed**Remember the octet rule and that Hydrogen and Helium are weird**
Step 5: If atom is still unstable after lone pairs
are added. Look to move pairs of electrons to
form either double or triple bonds.
Step 6: Consider whether or not the compound
is a polyatomic ion
i.e. it cannot be stable and will have a charge
Covalent compounds
-Sharing electrons means that forces between the
molecules are weak
-Relatively low melting points
-Do not conduct electricity [No free ions]
Polyatomic ion – a charged group of covalently
bonded atoms
VSEPR Theory – electrostatic
repulsion between valence
level electron pairs
surrounding an atom causes
these pairs to be oriented as
far apart as possible
**i.e. the 3-d shape of the
molecule**
*Based on the fact that bonded
atoms and electrons want to
be as far away from each
other as possible*
Intermolecular Forces
A.K.A Van der Waals forces –attractive forces between molecules.
-Forces of attraction between molecules
i.e. the forces between all of the water molecules in a glass of water
** 2 types**
-Dipole-Dipole force- which occurs between the molecules of polar covalently bonded molecules
-London Dispersion Forces- which occur between the molecules of non-polar covalently bonded molecules
Dipole – equal but opposite charges separated by
a short distance.
Dipole – dipole force – forces of attraction between
polar molecules
Dipole-Dipole Forces are very common with
compounds containing Hydrogen
Hydrogen bonding – a dipole-dipole force;
London dispersion forces – weak intermolecular
forces responsible for the attraction between
nonpolar molecules.
London dispersion forces result from the constant
motion of electrons [creates instantaneous and
induced dipoles]