Chemical Bonding
I.
II.
Groups - columns going down the periodic table. Elements within a group have
similar properties because they have the same number of valence electrons.
a. Valence Electrons are in the highest occupied energy level (outer shell). They
determine the physical and chemical properties of the element.
i. The group number for the representative element tells the number of
valence electrons.
Group 1A has one valence electron.
Group 2A has two valence electrons.
Group 3A has three valence electrons.
Group 4A has four valence electrons.
Group 5A has five valence electrons.
Group 6A has six valence electrons.
Group 7A has seven valence electrons.
Group 8A has eight valence electrons (except Helium with only 2).
b. Lewis Electron Dots show the valence electrons of the atom.
i. Inner electrons and the nuclei are represented by the element’s symbol.
ii. Bots representing electrons are arranged symmetrically around symbol.
iii. Bonds are formed between atoms using unpaired valence elecctrons
1. Write the electron dot structures for the following compounds
Water H2O
Ammonia NH3
Methane CH4
Carbon Dioxide CO2
Noble Gases are inert – they are extremely stable and do not react under normal
laboratory conditions. They have eight electrons in their highest energy level – eight
valence electrons is a stable electron configuration.
a. Octet Rule - atoms in compounds tend to create the electron configuration of a
noble gas – eight electrons in their highest occupied energy level.
i. Metals lose electrons to obey this rule and become stable like a noble gas.
ii. Nonmetals gain or share electrons to obey the octet rule, become stable.
b. Cations - neutral atom lost valence electrons & became a positively charged ion.
i. Metals with up to three valence electrons that can be easily removed.
(Group 1A, 2A, 3A)
ii. Unstable neutral atoms lose enough electrons to fulfill octet rule in the
next energy level down to become more stable as positively charged ions.
iii. Proton number never changes. A neutral atom has equal numbers of
protons and electrons. If electrons are lost then there are less electrons
than protons and the atom is now positively charged and called a cation.
iv. Examples: K loses 1e- to attain the stable noble gas configuration of
Argon. Al loses 3e- to attain the stable noble gas configuration of Neon.
(Board Game Analogy)
c. Anions - neutral atom gained valence electrons & became a negatively charge
ion.
III.
IV.
i. Nonmetals with up to three valence electrons that can be easily added.
(Group 5A, 6A, 7A)
ii. Unstable neutral atoms gain enough electrons to fulfill octet rule in their
highest energy level and become more stable as negatively charged ions.
iii. Proton number never changes. A neutral atom has equal numbers of
protons and electrons. If electrons are gained then there are more electrons
than protons and the atom is now negatively charged and called an anion.
iv. Examples: Cl gains 1e- to attain the stable noble gas configuration of
Argon. N gains 3e- to attain the stable noble gas configuration of Neon.
(Board Game Analogy)
d. Polyatomic Ions – a group of atoms that acts as a unit with a single charge
i. Begin memorizing polyatomic ions…get the list from the website and
make flashcards.
1. Know the formula, the charge, and the correct spelling of the name
of the seventeen polyatomic ions listed on the website.
Ionic Compound – between metal (cation) and a nonmetal (anion) = (M)(NM)
a. Ionic Bond - Completely transfer electrons.
i. Positive charge – cation – lost electrons to the anion.
ii. Negative charge – anion – gained electrons from the cation.
b. Positive charge must equal and, therefore, cancel the negative charge. Example:
Sodium Chloride – sodium wants to lose one electron to become stable and
chlorine wants to gain one electron to become stable. Will fulfill the octet rule
once they combine. (+)(-) = 0
c. Formula unit – a chemical formula of the smallest sample of an ionic compound.
d. Properties of Ionic Compounds
i. Crystalline solids at room temperature. Arranged in repeating threedimensional patterns. Very stable. Structures determined by X-ray
diffraction crystallography. Example: In solid NaCl, each Na is
surrounded by six Cl and each Cl is surrounded by six Na.
ii. Have very high melting points. Separates each ion from one another. Hard
to break the attraction between the ions. Very stable. Example: NaCl melts
at 800 ˚Celsius.
iii. Conduct electric currents when molten (liquid) or dissolved in water
(aqueous). The cations and anions migrate freely.
iv. Ionic compounds – are electrically neutral salts. (Many appear as minerals
in the Earth’s crust.)
e. Ionic Character
i. Ionic compounds have the greatest ionic character with full on charged
ions. The further the ions are apart in electronegativity, the more the ionic
character.
ii. Molecular compounds have very low electronegativity. The closer the ions
are in electronegativity, the less the ionic character.
Molecular Compounds - (NM)(NM) – are often are multiples of the lowest wholenumber ratios of nonmetals. Examples: C3H6 and C4H10 Note: Do not reduce
molecular compounds.
a. Covalent Bonds - the sharing of electrons between two nonmetals – creates a
molecular compound (or molecule). The goal is to attain eight valence electrons –
stability – similar to a noble gas electron configuration.
i. Do not forget that Hydrogen is a nonmetal.
ii. Do not forget your diatomic molecules in Group 7A: N2, O2, F2, Cl2, Br2,
I2, and H2
iii. Lewis Structures
1. Shared Pairs – both atoms can claim the electrons to achieve the
octet rule and become like noble gas configurations.
2. Unshared Pairs - pairs of valence electrons that are not shared
between atoms – also called lone pairs. Example: F2
iv. Single Covalent Bond - formed when one pair of electrons is shared
between two atoms. Example: Hydrogen – diatomic molecule…H2
v. Double Bond - involve two shared pairs of electrons. (Oxygen will form
the double bond but be an exception to the octet rule.) Example: Oxygen diatomic molecule…O2
vi. Triple Bond - involve three shared pairs of electrons. Example: Nitrogen
– diatomic molecule…N2
b. Structural Formulas - chemical formulas that show the arrangement of atoms in
molecules. A dash represents a pair of shared electrons (never used to show ionic
bonds because ions do not share electrons).
i. No unshared (lone) pairs visible.
ii. Shared pairs: represented with a line (dash) instead of two dots.
iii. Resonance - when two or more electron dot structures can be written for a
molecular compound. Example: NO2
iv. Exceptions to the Octet Rule
1. It is impossible to fulfill the octet rule whenever the total number
of valence electrons in the compound is an odd number. Oxygen
and Boron tend to be satisfied with less than eight valence
electrons. (6-7) Example: BF3
1. Phosphorus and Sulfur tend to accept more than eight valence
electrons because they have the d sublevel to expand into. (10-12)
Example: SF6 and PCl5
c. Properties of Molecular Compounds
i. Do not conduct electricity.
ii. Solids, Liquids, and gases.
iii. Low melting and boiling points that only separate one molecule from
another as opposed to separating each atom from another.
d. VSEPR Theory - valence shell electron pair repulsion. The electron dot
structures are not flat 2D structures, but are 3D in real life.
i. Shapes - Electron pairs repel. Molecules adjust their shapes so that the
valence electron pairs are as far apart as possible.
1. Linear – angles are 180 degrees – definitely will be linear if only
have two atoms in the molecule. No lone pairs and two covalent
bonds or three lone pairs and one covalent bond around central
atom. Example: CO2
2. Bent – again, unshared pair(s) strongly repels the covalent bonding
pairs. Two lone pairs and two shared pairs around central atom. All
angles are 105 degrees. Example: H2O
3. Trigonal-Planar – three shared pairs (covalent bonds ) separate as
much as possible, but are unaffected by a lone pair (no lone pairs)
of electrons like the pyramidal structure. Example: BF3
4. Trigonal-Pyramidal – one unshared pair strongly repels the three
shared pairs (covalent bonding), pushing them closer together. All
angles are 107 degrees. Example: NH3
5. Tetrahedral – four faced – four shared pairs and no lone pairs, all
angles are 109.5 degrees. Example: CH4
6. Trigonal Bipyramidal – five shared pairs separate as much as
possible, but are unaffected by a lone pair of electrons (no lone
pairs). Example: PCl5
7. Octahedral – six shared pairs separate as much as possible, but are
unaffected by a lone pair of electrons (no lone pairs). Example: SF6
ii. Can predict the shape of the molecule (as a general rule) using the group
number (valence electrons)
1A Linear
Example:
Li
2A Linear
Example:
Be
3A Bent
Example:
B
4A Tetrahedral
Example:
C
5A Pyramidal
Example:
N
6A Bent
Example:
O
7A Bent
Example:
F
iii. Hybridization - two atoms combine, their atomic orbitals overlap to
produce molecular orbitals. One electron from each atomic orbital
combines to create a shared pair in a molecular orbital.
1. sp3 hybridization – has electrons in 4 orbitals
2. sp2 hybridization – has electrons in 3 orbitals
3. sp hybridization – has electrons in 2 orbitals
e. Polarity – nonmetals of unequal strength (electronegativity) do not share electrons
equally. Note: This only applies to Molecules.
i. Polar Molecules - has a polar bonds, one end of the molecule has a
slightly negative charge while the other end has a slightly positive charge.
Dipole – a molecule with oppositely charged ends.
1. Polar Covalent Bonds - when the atoms are of different types, the
bonding electrons are shared unequally. The atom with the stronger
electronegativity acquires a slightly negative charge as it draws the
electrons toward itself. The atom with the lower electronegativity
acquires a slightly positive charge as the electrons are drawn away
from it. (Reminder: Electronegativity is the ability of the atom to
attract electrons to itself.)
1. Examples: HCl has 2 different types of atoms. Chlorine is
more electronegative than Hydrogen. Chlorine pulls the
shared pair closer to its own nucleus creating a partial
negatively charge pole. Hydrogen allows the shared pair to
be pulled farther from its own nucleus creating a partially
positively charge pole. H2O has 2 different types of atoms.
Oxygen is more electronegative than Hydrogen. The bent
shape due to the lone pairs creates oppositely charged ends.
2. The greater the difference in electronegativity the greater the
polarity.
1. Example: SO2 is less polar than H2O. Sulfur and Oxygen
are located close together on PT small difference in
electronegativity slightly polar molecule. Hydrogen and
Oxygen are located far apart on PT large difference in
electronegativity highly polar molecule.
2. Polar compounds are high in ionic character due to the
partially charged poles on the molecule.
1. The greater the difference in electronegativity, the
greater the polarity on the molecule and the more
ionic character the molecule has.
ii. Nonpolar Molecules – Either the molecule has no oppositely charged
ends or the ends cancel each other out.
1. Polar Covalent Bonds cancel each other out. Nonpolar molecules
may have polar bonds but the overall molecule is nonpolar because
polar ends cancel. Note: It will be a nonpolar molecule if the
molecule is symmetrical in 3D and all bonds are exactly the same.
Example: CO2
2. Nonpolar Covalent Bonds - when atoms are the same type, they
share the bonding electrons equally. Because have same
electronegativity. This is the case with all the diatomic molecules.
H2 O2 N2 etc.
3. Nonpolar compounds have very little ionic character as they
sometimes exhibit dispersion forces when their ions vibrate to
create a momentary dipole.
f. van der Waals Forces – intermolecular attractions - attractions between
molecules – weak compared to ionic or covalent bonds but still substantial in
strength
i. Dipole Interactions - when polar molecules are attracted to one another:
opposite charged regions of polar molecules are attracted.
1. Hydrogen Bonds – a particularly strong dipole interaction
specifically involving hydrogen at the partially positive pole.
Hydrogen is covalently bonded to a very electronegative atom
AND to an unshared pair of another atom.
1. Hydrogen is able to bond with the unshared pair of
electrons from another molecule because its’ valence
electrons are not shielded from the nucleus by another layer
of electrons (hydrogen’s valence electrons are directly up
against the nucleus). Example: H2O
2. The more electronegative the element that hydrogen is
bonded to the stronger the intermolecular attractions.
Example: Which element has stronger intermolecular
interactions: H2S or H2O?
ii. Dispersion Forces - weakest of all molecular interactions – caused by the
motion of electrons.
1. Vibrating electrons may end up moving randomly closer to one
atom or another creating a momentary dipole.
2. The more electrons the greater the interaction between nonpolar
molecules.
Examples: Halogens
F2 – molecules not touching - gas
Br2 – molecules touching and sliding - liquid
I2 – touching and vibrating in place - solid
V.
Metallic Bonding groups of closely packed cations in a “sea” of free moving
valence electrons
a. Valence Electrons create an attraction between the free moving valence electrons
in the positively charged metal cations
b. Properties of Metals
i. Good conductors of electricity – electrons enter one end of the metal bar
and leave the other.
ii. Ductile – can be stretched into wires.
Malleable – can be pounded into shapes. Metals ions slide passed one another in a sea of drifting