Unit 4
Chemical Bonding Notes
1. All Elements (except the 6 noble gases) form bonds to achieve stability.
2. Instability comes from unfilled energy levels.
3. A stable compound is formed when the total energy of the combined atoms is lower
than the separate atoms.
4. Atoms other than the noble gases react so that they end up with a full outer shell of
electrons; i.e. their electron configurations mimic or copy that of the noble gases.
Metals Lose Electrons
Nonmetals Gain Electrons
Nonmetals Share Electrons
Ionic Bond Example:
Na (metal) 1s22s22p63s1 loses 3s1 electron
Cl (nonmetal) 1s22s22p63s23p5 gains the 1e- from Na
After reacting:
Na is now 1s22s22p6 (Neon’s config.) Na+1
Cl is now 1s22s22p63s23p6 (Argon’s config.) m Cl-1
Na+1 + Cl-1
NaCl
Covalent Bond Example:
Electrostatic Forces:
Opposite Charges Attract >> + and -<<
Like Charges Repel <<- and ->> <<+ and +>>
Important Pages:
p. 414 Common Group A Element e- dot structures
p. 429 Common Alloys
p. 445 Common Molecular Cpds.
2 Kinds of Elements Involved in Bonding:
1. Metals (alkali, alkaline earth, transition, metalloid)
2. Nonmetals (H & above stairstep)
Sugar and Salt Conductivity Demonstration
Ionic Compounds
Metal &
Nonmetal
Ionic Bonding
Ions
Electrolytes
Salts
Molecular Compounds
All Nonmetals
Covalent Bonding
Molecules
Nonelectrolytes
Sugars, Fats, Alcohols, Oils, Carbohydrates,
etc.
Compound (cpd.) – a chemical combination of 2 or more elements
Ionic Compound - compound composed of 1 metal and 1 or more nonmetals, a positive ion and
a negative ion, ex. NaCl, Fe2O3, NaHCO3 (salt, rust, baking soda)
Molecular Compound - compound composed of 2 or more nonmetals, ex. H2O, C6H12O6, NH3
(water, glucose, ammonia)
Ions – are atoms with a + or – charge, due to the loss or gain of electrons
Positive ions have more protons than electrons (have lost electrons)
ex. Na+1 (lost 1 e-) Fe+2 (lost 2 e-), etc.
Negative ions have less protons than electrons (have gained electrons)
ex. Cl-1(gained 1 e-) O-2(gained 2 e-), etc.
Electrolyte – a compound that conducts electricity in aqueous solution
Nonelectrolyte – a compound that does not conduct electricity in aqueous solution
Octet Rule – atoms react by gaining, losing (ionic) or sharing (covalent) electrons so as to
acquire the stable e- configuration of a noble gas, usually 8 electrons
For Group A Elements: the group number tells you the number of valence electrons
Valence electrons – the outermost (highest energy level) electrons involved in bonding
Chemical Bonding – an attraction between atoms brought about by sharing of electrons
(covalent) or a complete transfer of electrons (ionic) that results in formation of a
compound
Ionic Bond – attraction of positively charged metal ion and negatively charged nonmetal
ion that forms salts (crystalline solids)
Metallic Bond – the force of attraction that holds metals together; it consists of the
attraction of free-floating (mobile) valence electrons for positively charged metal ions, p.
427
Electrical Conductivity, Malleability and Ductility of metals are explained by the mobility
of valence electrons, the “sea” of drifting electrons insulates the positive metal ions from
each other. When metals are hammered the positive ions easily slide past each other
without repelling one another like ball bearings immersed in oil
Covalent Bond – a bond in which 2 atoms share a pair of electrons
Single covalent bond – a single pair of shared electrons
Double covalent bond – two shared pairs of electrons
Triple covalent bond – three shared pairs of electrons
Electronegativity – the tendency for atoms of an element to attract electrons to itself when
chemically bonded to another element
Bond Type (Ionic, Covalent, Polar Covalent, etc.) can be predicted by looking at the
difference in electronegativities of the two elements involved.
p. 462 19 & 20 sample problem 16-4
Types of Bonds:
1. Nonpolar covalent – electrons are shared equally
2. Polar covalent – bonding electrons are shared unequally, the most electronegative
element pulls on the shared electrons more giving it a negative charge
3. Ionic – bond – complete transfer of electrons
polar molecule – a molecule like water, in which one or more atoms is slightly negative and one
or more is slightly positive
VSEPR theory – valence shell electron-pair repulsion theory; the shape of molecules
adjusts so that e- pairs are as far apart as possible
electron dot (Lewis) structure – a notation that depicts valence electrons as dots around the
atomic symbol of the element; the symbol represents the inner electrons and atomic nucleus
Model e- Dot or Lewis Dot Structures for the following:
1. Diatomic Molecules – H2, N2, O2, F2, Cl2, Br2, I2
2. H2O (bent)
3. CO2 (linear)
4. NH3 (trigonal pyramid)
5. CH4 methane (tetrahedron)
6. HCl, HF, HI, HBr
7. H2O2
8. C2H6
9. C3H8
10. C4H10
11. COCl2 (trigonal planar)
Electron Configurations
The “Secret Code” for describing where a particular electron can be found.
The quantum mechanical or electron cloud model of the model is based on the mathematical
probability of finding an electron in a certain position. This probability can be portrayed as a
blurry cloud of negative charge. If a line is drawn around the model of the electron cloud so that
the electrons is in the cloud 90% of the time. This gives the electron cloud a distinct shape.
Think of being able to see where the blades of a fan are most of the time while it is moving. You
could draw a line around where the blades are most of the time.
electron configuration – the arrangement of electrons around the nucleus of an
atom in its ground state
energy level – a region around the nucleus of an atom where an electron is likely
to be moving
principal energy levels – correspond to a particular period or row on the periodic
table
quantum numbers – numbers that describe various properties of electron orbitals
atomic (electron) orbitals – mathematical calculations used to express electron cloud shape; a
region in space where there is a high probability of finding an electron
1s2
Kinds of orbitals (clouds):
p. 365 AW or p. 313 Zum.
s – sphere shaped (1 suborbital)
p – dumbbell shaped (3 suborbitals)
d – four clover leafs and a pacifier (5 suborbitals)
f – complex (7 suborbital)
Each suborbital can hold a maximum of up to 2 electrons, so each cloud shape can hold the
following maximum number of electrons:
s2
p6
d10
f14
Rules for governing the filling of atomic orbitals:
1. Aufbau principle – electrons enter orbitals of lowest energy first
2. Pauli exclusion principle – an atomic orbital may describe at most two electrons
with opposite spins
3. Hund’s rule – when electrons occupy orbitals of equal energy, one electron enters
each orbital until all the orbitals have one electron with parallel spins
Aufbau Periodic Table p. 395 AW
Na configuration 1s22S22p63s1
Na Shorthand [Ne]3s1
There are exceptional electron configurations, for example:
Cr expected: [Ar]4s23d4
Cr actual:
[Ar]4s13d5
Cu expected: [Ar]4s23d9
Cu actual:
[Ar]4s13d10
An Unexcited atom has a Normal electron configuration,
An Excited atom may have electrons that skipped to higher levels (not normal); for example:
1s22s22p83s1