Transcript Review of Bonding and Bond Types

Review of Bonding and Bond
Types
Valence Electrons
• The electrons in the highest occupied level of an
element’s atoms.
• The number of valence electrons largely determines the
chemical properties of an element.
• Can be determined by looking at the electron
configuration or dot diagram for an element.
• Alkali metals have 1, Alkaline Earth metals have 2,
Transition metals have 2 or 3, Boron group has 3,
Carbon group has 4, Nitrogen group has 5, Chalcogens
have 6, Halogens have 7, and the Noble Gases have 8
(an octet).
• Chemist Gilbert Lewis explained that atoms form certain
kinds of bonds due to atoms tendency to achieve the
electron configuration of a noble gas (octet rule).
Electronegativity
• The tendency
for the atoms
of the element
to attract
electrons
when they are
chemically
combined with
atoms of other
elements.
Ionic Bonding
• Two elements with very different electronegativities
(difference between 4.0 and 2.0).
• Elements with high electronegativity pull electron(s)
more strongly and therefore steal the electron(s)
from the element with a low electronegativity.
• The stronger elements therefore becomes
negatively charged (anion) since it has gained
electrons and the weaker elements becomes
positively charged (cation) since it has lost
electrons.
• Since they have opposite charges, they attract one
another by electrostatic forces which creates the
bond.
• Usually between one non-metal and one metal.
Drawing Lewis Structures for Ionic
Bonds
•
The overall charge on the compound must equal zero, that is, the number of
electrons lost by one atom must equal the number of electrons gained by
the other atom.
• The Lewis Structure (electron dot diagram) of each ion is used to construct
the Lewis Structure (electron dot diagram) for the ionic compound.
Examples:
• Lithium fluoride, LiF
– Lithium atom loses one electron to form the cation Li+
– Fluorine atom gains one electron to form the anion F– Lithium fluoride compound can be represented as
•
[Li+]
Lithium oxide, Li2O
– Each lithium atom loses one electron to form 2 cations Li+ (2 electrons in total are
lost)
– Oxygen atom gains two electrons to form the anion O2– Lithium oxide compound can be represented as
[Li+]
[Li+]
Ionic Bond Characteristics
• Compounds are called formula units
• Most are crystalline solids at room temperature (pg 422 has some great
pictures)
• In solids NaCl, each sodium cation is surrounded by six chloride anions
and each chlorine anion is surrounded by six sodium cations. Giving
them both a coordination number (the number of ions of opposite charge
that surrounds the ion in a crystal) of 6.
• In this arrangement, each ion is attracted strongly to each of its
neighbors and repulsions are minimized, which results in a very stable
structure.
• High melting point (usually above 300 °C)
• High solubility (dissolving ability) in water since water is polar.
• Good conductor of electricity either when melted (the orderly crystal
structure breaks down allowing cations and anions to move freely
through the liquid) or when mixed in water (ions also become free to
move around allowing electrons to flow).
• Brittle, not very malleable because when struck with a hammer, the blow
tends to push ions of like charge into contact and they repel, shattering
the crystal.
Metallic Bonding
• Metals are made up of closely packed cations
rather than neutral atoms. The cations are
surrounded by mobile valence electrons
(delocalized electrons), which drift freely from
one part of the metal to another.
• The attractions of the free-floating valence
electrons for the positively charged metal cations
creates the bond that holds the metal together.
Metallic Bond Characteristics
• Good conductors of electrical current because electrons can flow
freely in them. As electrons enter one end of a bar of metal, an equal
number leave the other end.
• Ductile and malleable due to the metals ability to be forced through
a narrow opening to produce a wire or shaped with a hammer
without shattering. A sea of drifting valence electrons insulates the
metal cations from one another. When a metal is subjected to
pressure, the metal cations easily slide past one another like ball
bearing immersed in oil.
• Also crystalline solids at room temperature (except for mercury).
Three arrangements of the atoms
– Body-centered cubic
– Face-centered cubic
– Hexagonal close-packed
Covalent Bonding
• Two elements with similar
electronegativities (difference is less than
2.0).
• Since they have similar pull on the
electron(s) they end up both holding onto it
and the electron(s) and sharing them,
which creates the bond.
• Usually both non-metals.
Covalent Bonding Characteristics
•
•
•
•
Compounds are called molecules.
Can be solid, liquid or gas at room temperature.
Low melting point (usually below )
High to low solubility in water (depending on the
molecules’ polarity)
• Poor to nonconducting when dissolved in water
(aqueous).
• A greater range of physical properties is found due to the
widely varying intermolecular attractions (polar and
nonpolar bonding)
• Can have single, double or triple covalent bonds
between two atoms.
Polarity
• Polar bonds are when one atoms has a slightly
stronger pull on the electrons that are being
shared (difference between 2.0 and 0.4).
Therefore, one end of the bond is slightly more
negative (the one that is slightly stronger) and
one is slightly more positive (the one that is
slightly weaker).
• All the diatomic elements (Nitrogen, Oxygen,
Fluorine, Chlorine, Bromine, Iodine and
Hydrogen) have non-polar covalent bonds since
they share the electrons equally.
Polar Bonds vs. Polar Molecules
• The presence of a polar bond in a molecule
often makes the entire molecule polar, where
one end of the molecule has a slightly more
negative charge and the other end is slightly
more positive (called a dipole).
(Ex: H20 and NH3)
• However, if the polar bonds are arranged
symmetrically around the molecule the entire
molecule is actually non-polar.
(Ex: CO2 and CH4)
Other Intermolecular Forces
• Molecules are often attracted to each other by a variety of
other forces, although weaker than ionic or covalent bonds.
– Van der Waals forces:
• London Dispersion forces-the weakest of all molecular interactions
caused by the motion of electrons.
• Dipole-induced dipole interactions-occurs when polar molecules cause
electrons in another molecule to move to one side of the molecule as
they are attracted to the dipole of the other molecule.
• Dipole-dipole interactions- occur when polar molecules are attracted to
eachother since they have different charges on each side.
– Hydrogen bonds:
• Attractive forces in which a hydrogen covalently bonds to a very
electronegative atom is also weakly bonded to an unshared pair of
electrons of another electronegative atom. Since hydrogen only has
one electron when it bonds with a highly electronegative atom, the
nucleus of the hydrogen atom becomes very electron deficient.
Therefore, the sharing of a nonbonding electron pair on a nearby
electronegative atom compensates for the deficiency.