Chemistry Review
Electron Configuration
For an electron in an atom, specifying that n=2, ℓ=1 is equivalent to saying that it is located in the 2p subshell.
Similarly, n=3, ℓ=0 specifies the 3s subshell, and so forth. Thus, one can construct the following table:
Quantum numbers Subshell
n=2, ℓ=1
2p
n=4, ℓ=0
4s
n=6, ℓ=1
6p
n=6, ℓ=2
6d
n=4, ℓ=3
4f
*Orbitals from smallest to largest energy: 1s,2s,2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
Atomic Size

Atomic size tends to decrease left to right across the rows and increase down the columns.
*In the top right corner of the periodic table are the smallest atoms
Example: Rb>Sr>Sb>As>Se>F
Atomic Radius

Atomic radii increase going down a group, because successively larger valence-shell orbitals are
occupied by electrons. Moving left to right across a period, atomic radius decreases because the
effective nuclear charge increases.
Example: Atomic Radius increasing down a group- Rb>K>Na>Li
Atomic radius decreasing across a period- Na>Mg>Al>Si
Shielding of Electrons
Example: The shielding of electrons gives rise to an effective nuclear charge, Zeff, which explains why boron is
larger than oxygen. Estimate the approximate Zeff felt by a valence electron of boron and oxygen, respectively?
Boron has five protons and two inner (nonvalence) electrons, so 5-2=+3; Oxygen has eight protons and two
inner (nonvalence) electrons, so 8-2=+6. The valence electrons in an oxygen atom are attracted to the nucleus
by a positive charge nearly double that of boron. Therefore, the electrons in oxygen are held closer to the
nucleus, giving it a smaller radius.
*The ordering of electrons in orbitals according to their shield ability would look like 4s>4f>4d…
*Zeff increases across a period
Cations and Anions

An atom is made up of a positively-charged nucleus surrounded by negatively-charged electrons. The
atom is neutral when there are equal numbers of protons and electrons. Ions are formed when there are
unequal numbers of protons and electrons in an atom. A positive ion is called a cation whereas a
negative ion is called an anion.
Example: Using the fact that the number of electrons in the Cl− ion is equal to the number of electrons present
in the neutral chlorine atom plus the charge on the Cl− ion, eighteen electrons are in this ion.
Therefore, the new electron configuration is [Ne]3s23p6
Ionic Radii

In a group of ions with the same number of electrons, the most negative ion is the largest and the most
positive ion is the smallest

Ionization energies tend to decrease going down a column of the periodic table but increase going across
a row.

If an atom has a relatively high first ionization energy, this indicates that the electrostatic interactions
between the positively charged nucleus and the negatively charged electron are relatively high. In other
words, the ionization energy of such an atom is high because the nucleus is attracting that electron more
strongly, making it more difficult to pull that electron away. This property also affects the size of an
atom. Thus, nuclei that attract their electrons more strongly tend to pull the electrons into closer
proximity to the nucleus, shrinking the atom's overall size.
Element
Radius
(pm)
X
119
Y
191
Z
261
*Highest to lowest first ionization energies: X>Y>Z, so smaller the radius, larger ionization energies
Electron Affinity


The electron affinity, Eea, is the energy associated with the gain of an electron. The more stable the
electron configuration obtained, the more energy is released. Recall that half-filled and completely filled
orbitals are more stable. Begin by looking at the electron configuration of the anions derived from the
given electron configurations.
The tendency to gain an electron is quantitatively measured by the electron affinity, the amount of
energy involved in the addition of an electron to a neutral gaseous atom. Ordering these elements by the
electron affinity provides an identical order: Cl>S>Si>Na>Mg
Naming Compounds
Cations
Naming convention for cations
Example
If the element forms only one ion,
simply name the element.
K+ (potassium)
Cations that make multiple ions are
distinguished by Roman numerals.
Fe3+ [iron(III)]
Polyatomic ions have special names.
NH4+ (ammonium)
Anions
Naming convention for anions
Example
If the element forms only one ion,
add ide to the element name.
Cl− (chloride)
When there are only two oxoanions in
a series, use ite for the anion with
fewer oxygen atoms and ate for the
anion with more oxygen atoms.
SO 2−4 (sulfate)
SO 2−3 (sulfite)
ClO− (hypochlorite)
Halogens and transition metals that can
ClO2− (chlorite)
form up to four oxoanions require
ClO3− (chlorate)
distinguishing prefixes.
ClO4− (perchlorate)
Lattice Energy

As you move down a group, lattice energy decreases
For bond length, just add the two atomic radii together; (Units in pm)
Electron Negativity

Electronegativity is another periodic property. An atom's electronegativity describes its ability to attract
electrons to itself when it is part of a chemical compound. Electronegativity increases diagonally from
the lower left to the upper right of the periodic table. Highly electronegative elements (with a chi value
of χ≥2.2) are insulating nonmetals, whereas elements with low electronegativity (with χ≤1.8) are
conducting metals.
Bond type depends on the atoms' relative attraction for electrons. Therefore
- a metal–nonmetal bond tends to be ionic,
- a nonmetal–nonmetal bond tends to be covalent, and
- a covalent bond between atoms of the same element is nonpolar.
Use the periodic table to classify each element in the bonds given. Recall that nonmetals are mostly in the upper right
corner of the periodic table.
The difference in electronegativity between the two elements in a bond can be used to predict bond type. As a
general rule,
-
if the difference is zero, the bond is nonpolar covalent,
if the difference is greater than two, the bond is substantially ionic, and
if the difference is greater than zero but less than two, the bond is polar covalent.
Example: Classify the bond in each compound as ionic or covalent. You can use either method (metal versus
nonmetal classifications or electronegativity differences) to do this. The electronegativity values are listed here:
Element Electronegativity
Na
0.9
H
2.1
Cl
3.0
F
4.0
-
Once you know the type of bond (and therefore the type of compound), you can relate this to the boiling point.
Recall that more charge separation leads to greater attraction within a substance as a whole. The greater the
attraction, the harder it is for atoms or molecules to be released into the gas phase.
-
Hint 2. Relate the type of compound to boiling point
-
Which class of compound tends to have high melting and boiling points?
ionic compounds
covalent compounds
*Bigger difference, the more polar it is
Lewis Dot Structure and Formal Charges

To determine the formal charges of the atoms in a molecule, we must assign each electron to a particular atom.
Then, we compare the number of electrons assigned to each atom to the expected valence of that element.
All lone pairs of electrons should be assigned to the atom on which they are found; bonding pairs of electrons should be
divided equally between the atoms in the bond. The formal charge is calculated by subtracting the number of electrons
assigned to the atom from the number of valence electrons in the free atom.
Formal charge = (Number of valence e− in free atom) - 12(Number of bonding e−) - (Number of nonbonding e−)
For example, in ∙∙C≡O∙∙, each atom is assigned five electrons (two nonbonding from the atom's lone pair and three
bonding from half of the triple bond). Since the expected valence of carbon is four, its formal charge
is (4 valence)−(5 assigned)=−1. Since the expected valence of oxygen is six, its formal charge
is (6 valence)−(5 assigned)=+1.
For a neutral molecule such as CO, the sum of the formal charges is zero. For a polyatomic ion, the sum of the formal
charges is equal to the overall charge on the ion.
1.
2.
3.
4.
5.
Octet Rule:
To determine whether the octet rule is followed, the number of electrons around each atom in the formulas must be
identified. This can best be done by drawing the electron-dot structure for each atom. Here is a set of guidelines to follow
in drawing electron-dot structures:
Determine the total number of valence electrons using the number of valence electrons for each atom in the formula
If the species is a cation, subtract a number of electrons equal to the ionic charge. If the species is an anion, add a
number of electrons equal to the ionic charge.
The central atom is usually the more electropositive element. This excludes H, which is never a central atom.
Bond all atoms in the appropriate skeletal arrangement so that all atoms (except H) have an octet (or their normal number
of bonds). If there are not enough electrons to accomplish this, use double or triple bonds to achieve the octet.
If there are electrons left over, place the electrons as lone pairs on the central atom if the atom is in period 3 or higher.
Molecular Shapes and Bond Angles
The shape and bond angles of a molecule can be predicted using VSEPR (valence-shell electron-pair repulsion) theory.
According to VSEPR theory, bonds and nonbonding electron pairs exist as localized charge clouds around a central atom.
They repel one another, forming the shapes listed here.
O uter Lone Charge
atoms pairs clouds
Shape
Angle
(degrees)
2
0
2
linear
180
3
0
3
trigonal planar
120
2
1
3
bent
120
4
0
4
tetrahedral
109.5
3
1
4
trigonal pyramidal
109.5
2
2
4
bent
109.5
5
0
5
4
1
5
seesaw
90, 120, 180
3
2
5
T-shaped
90, 180
2
3
5
linear
180
6
0
6
octahedral
90
5
1
6
square pyramidal
90
4
2
6
square planar
90
trigonal bipyramidal 90, 120, 180
If BP stands for a bonded pair of electrons and LP stands for a lone pair of electrons, then BP-BP represents the repulsion
between two bonded pairs, LP-LP represents the repulsion between two lone pairs, and BP-LP represents the repulsion
between a bonded pair and a lone pair of electrons.
The molecular geometry of a molecule describes the three-dimensional shape of just the atoms. This is in
contrast to the electronic geometry, which describes the shape of all electron regions.
To determine the molecular geometry of a molecule, one must first determine the electronic geometry by
drawing the Lewis structure.