Study Guide
Chemical Bonding
Electronegativity – the tendency for an atom to attract electrons to itself when it is
chemically combined with another element – the measure of the ability of an atom or
molecule to attract electrons in forming a chemical bond
Coulomb’s law – the electrical force of one body exerted on a second body is equal to
the force exerted by the second body on the first
Vector – a object with a magnitude and a direction
Valence Electrons
– A valence electron is an electron in the highest occupied energy level of an atom
– All of the elements in each group of the periodic table have similar behaviors
because they have the same number of valence electrons. Group IA electrons, for
example, all have 1 valence electron in their highest energy levels.
– The number of valence electrons determines the chemical properties of elements.
– Look at an element’s electron configuration to determine the number of valence
electrons. Also look at the periodic table to determine the number of valence
electrons for Group A elements.
– Transition elements (metals)have varying numbers of valence electrons. For
example, copper may have 1, 2, or 3 valence electrons.
Periodic Table
Number of Valence
Electrons
Group 1 (I)
Alkali Metals
1
Group 2 (II)
Alkaline Earth Metals
2
Groups 3-12
Transition Metals
1 or 2
Group 13 (III)
Boron Group
3
Group 14 (IV))
Carbon Group
4
Group 15 (V)
Nitrogen Group
5
Group 16 (VI)
Oxygen Group
6
Group 17 (VII)
Halogens
7
+
Group 18 (VIII or 0)
Noble Gases
8
+
Helium is an exception to Group 18. It has only 2 valence electrons.
Family or Group Number
Family or Group Name
Electronegativity increases going up the periodic table and going left to right across the
periodic table. This is opposite to atomic radius size. The atomic radius is the distance
from the nucleus to the outermost stable electronic orbital. As atomic radius decreases,
the electronegativity increases
Chemical Bonds
- Chemical bonds are the attraction of atoms to each other through the sharing
(covalent bond), the exchanging (ionic bond) of electrons, or through the
interactions of electrostatic forces
- Intramolecular – bonds occur within a molecule – can be ionic, covalent, or
metallic
Ionic bonds
o Ionic bonds are bonds that form between a metal and a nonmetal element
o Metals have low electronegativity and nonmetals have high
electronegativity.
o Ionic bonds occur because of this difference in electronegativity
o Metals give up electrons to nonmetals
o When metals give up electrons, they form cations. Cations are positively
charged particles.
o When nonmetals take on electrons, they form anions. Anions are
negatively charged particles.
o Ionic compounds are good conductors of electricity when in solution or in
a molten (melted) state. They have a high melting point and are generally
soluble in water.
o Ionic bonds have electronegativity values greater than 1.7.
Covalent bonds
o Covalent Bonds occur when valence electrons between nonmetal elements
are shared with one another.
o Nonmetals have similar electronegativity values
o The atoms do not become anions or cations – they simply share electrons
o Covalent bonds are directional in nature – strength of the bonds are
influenced by the bond angles
o Bonds can be polar or non-polar. Polar bonds have an electronegativity
difference greater than zero but less than 1.7. Polar covalent molecules
share both the properties of ionic and non-polar covalent molecules.
o A non-polar covalent bond has an electronegativity difference of zero.
They are insoluble in water and are not good conductors of electricity.
They usually exist as individual molecules.
o Polymerization occurs when molecules react and chain together to form
sheets or structurally repeating units of materials – plastics, nylon, vinyl
etc.
Metallic Bonds
o A metallic bond is the force of attraction that holds metals together; it
consists of the attraction of free-floating valence electrons for positively
charged metal ions – bonding within metals
o The metals in metallic bonds have low ionization energy and atoms are
easily stripped of their electrons; they tend to flow around very easily;
crystal structures form and the positively charged metal ions are held
infixed positions
o Metallic bonding is non-polar – little or no difference in electronegativity
o Bonds are strong due to high boiling and melting points
Characteristics of Metals
o Metals are usually hard due the crystal structure of the bonds between the
charged ions
o Metals conduct heat.
o Metals are malleable – they can be hammered or shaped into other forms
o Metals are ductile – they can be formed into wire
o Metals are lustrous – valence electrons absorb and emit light and the light
interacts with the surface for the metal or the crystalline structure.
Intermolecular Bonds
o Intermolecular bonds form as a result of the attractive forces that hold
molecules together
o Intermolecular forces are weak.
Three types of intermolecular forces
o Hydrogen Bonds
 The strongest of the intermolecular forces, but not a true bond
 exist between two polar molecules that contain Hydrogen and one
other highly electronegative atom. (Fluorine, Oxygen, or Nitrogen)
 Polar molecules have a slight positive charge on one end of the
molecule (usually where the Hydrogen is located) and a slight
negative charge on the other end of the molecule. (The Fluorine,
Oxygen, or Nitrogen) Water (H2O) is a perfect example.
 Polarity depends on the difference in electronegativity between the
atoms
 Hydrogen has low electronegativity; when bonded to atoms with
high electronegativity, like fluorine, nitrogen, or oxygen (H2O
again), the bond becomes highly polarized
o Van der Waals Forces
 exist between non-polar molecules
 Non-polar molecules have very little, if any, difference in
electronegativity
 Examples of non-polar molecules include H2, CO2, N2, and the
noble gases (Group VIIIA)
 At STP (standard temperature and pressure), there are virtually no
Van der Waals forces but, as pressure increases and temperature
decreases, these forces become active and gases like hydrogen,
carbon dioxide, nitrogen, and the noble gases will turn into liquids.
 Large molecules have more electrons and tend to have stronger
Van der Waals forces
o Molecue-Ion Attractions
 Molecule – ion attractions occur when ionic solids (like salt
NaCl) dissolve in water
 Water is made up of polar molecules and these polar molecules
strongly attract the ions in the metal forcing anions to be removed
from the crystal lattice that makes up the metal
Electron Configuration
- Electron configuration is the arrangement of electrons in an atom or molecule.
- The behavior of electrons is determined by their various configurations in atoms
and molecules.
- Orbitals – certain areas of atoms or molecules where electrons are likely to be
found
- Atomic orbitals – the energy states of individual electrons in an electron cloud –
orbitals are identified by a set of quantum numbers: n, m, l, and s. n is the
principal quantum number which determines the distance from the electron to the
nucleus – always a positive integer – all electrons with the same value of n are the
same distance from the nucleus and make up a shell.
- Molecular orbital – a linear expansion or combination of atomic orbitals
- Electrons have the ability to move from one energy level to another as their
energy states change. As they move from a higher energy level to a lower energy
level, they emit particle of light energy called photons.
- Schrodinger equation – describes the states of electrons in the atomic orbitals
- Subshells – regions of an electron that have two or more orbitals. The orbitals are
named s,p,d, and f. the s orbital is closest to the nucleus of an atom
- azimuthal quantum number (l) – the orbital’s angular momentum – a nonnegative integer
- subshells –orbitals with the same value of n and the same value of l
- magnetic quantum number ml – determines the energy shift of an atomic orbital
due to external magnetic fields
- spin quantum number(s) – a property of the electron – independent of the other
numbers
Writing Electron Configurations
Several rules govern the placement of electrons in orbitals:
1. No two electrons in an atom may have the same set of values for all four quantum
numbers. This is known as the Pauli exclusion principle.
2. Only two electrons may occupy a single orbital as long as they have different
electron spin values (quantum number s).
3. An electron always tries to occupy the lowest possible energy state before filling
higher states. This is called the Aufbau principle.
4. When there are multiple orbitals at a particular energy level, such as 3 p orbitals
or 5 d orbitals, only 1 electron will fill each orbital until each has an electron.
After this, pairing will occur starting with the lowest energy states and working up
to the higher states. This is called Hund’s Rule of Maximum Multiplicity.
5. An electron can occupy any orbital as long as it does not violate the Pauli
exclusion principle. If lower-energy orbitals are available, the electron will
eventually lose energy and drop into the lower orbital.
6. The maximum number of electrons in orbitals at a particular sublevel is as
follows:
A. s=2, or 1 orbital
B. p=6, or 3 orbitals
C. d=10, or 5 orbitals
D. f=14, or 7 orbitals
Remember that the maximum number of electrons in each individual orbital is 2.
7. The order of sublevel filling occurs as follows:
1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d10, 4p6, 5s2, 4d10, 5p6, 6s2, 4f14, 5d10, 6p6, 7s2
Now that we know this information, we can write orbital names. Orbital names are
written in the following manner:
n shapey
Here’s another way to look at all this:
The regions where electrons are likely to be found in an atom are called atomic orbitals.
Letters denote the atomic orbitals – s,p,d, and f. The numbers and kinds of atomic
orbitals depend on the energy sublevel. The lowest principal energy level has only one
sublevel – 1s. The second energy level has two sublevels – 2s and 2p. The third energy
level has three sublevels – 3s, 3p, and 3d. The fourth energy level has four sublevels –
4s, 4p, 4d, and 4f. Notice that the principal quantum number always equals the number
of sublevels. In the first energy level, a maximum number of two electrons can be held.
The second level can hold a maximum number of 8 electrons. The third energy level can
hold up to 18 and the fourth can hold 32 electrons. In order to understand the electron
arrangement in atoms we must keep several rules in mind:
1. Electrons enter orbitals of lowest energy first – kind of like filling up seats in a
theater. Start filling up seats closest to the stage first (in this case the nucleus) and
work outward. Energy is lowest closest to the nucleus.
2. An atomic orbital may contain at most two electrons.
3. When electrons occupy orbitals of equal energy, like a p orbital, one electron
enters each orbital until all the orbitals contain one electron.
An oxygen atom contains eight electrons. The orbital of lowest energy, 1s, gets one
electron and then a second with an opposite spin. That’s the max that can be held in this
first energy level. The next orbital to fill is 2s. It also gets two electrons. One electron
then goes into each of the three 2p orbitals. The remaining electron goes into the first of
the 2p orbitals. The other two p orbitals have only one electron since the total of eight
has been reached.
Take a look at these common configurations:
1s1
1s2
1s22s1
1s22s22p2
1s22s22p3
1s22s22p4
H
He
Li
C
N
O
Notice that the total number of electrons (written in superscript) adds up to each atomic
number (the number of electrons).
Here’s how to write the electron configuration for phosphorus:
Look at the periodic table and see that phosphorus has 15 electrons. Start placing the
electrons in the orbitals with the lowest energy (1s). Remember there is a maximum of
two electrons per orbital (example - the p sublevel has 3 orbitals with a max of 6
electrons). Electrons do not pair up within an energy sublevel (orbitals of equal energy)
until each orbital already has one electron. I cannot show the diagrams with the arrows
but I can show you the configuration:
P 1s22s22p63s23p3
Here is nickel:
Ni 1s22s22p63s23p63d84s2
And a few more:
C 1s22s22p2
N 1s22s22p3
Na 1s22s22p63s1
B
1s22s22p1