Chemistry Unit 7 Review
Alec Jotte
Honors Chemistry
Test Date: 3-21-13
Excited State vs. Ground State

Excited State
o

Atom with excess energy.
Ground State
o
Atom in the lowest possible state of energy
Electromagnetic Radiation

The dual nature of light
o
Wave and a Photon (packet of energy)
o

Wavelength ()
o


The distance between two peaks or troughs in a wave
Frequency (v)
o

Alec Jotte
Number of waves (cycles) per second that pass a given point in space
Speed (c)
o
Speed of light (2.9979 x 108 m/s)
o
c = v
Visible Spectrum
o
In order from highest to lowest energy

Quantized Emission
Purple, blue, green, yellow, red

Atoms give off light but must first receive energy and become excited.

The energy is released in the form of a photon that corresponds exactly to the energy change
experienced by emitting the photon.

The energy levels of all atoms are quantized (analogy of the stairs)
Hydrogen Orbitals

Bohr’s Model of the atom
o
Different orbits around the nucleus.
o
Only certain types of photons are produced when H atoms release energy. Why?

Since only certain energy changes occur, the H atom must contain discrete energy
levels.
o

Bohr’s model does no apply to any atoms other than hydrogen therefore it is ditched.
Orbitals: the sphere that contains 90% of the total electron probability.
o
This model gives no information about when an electron occupies a certain point in space or
how it moves.
o

Orbitals do not have sharp boundaries.
Hydrogen has discrete energy levels called principle energy levels.
o
They are labeled with whole numbers.
o
Each principal energy level is divided into sublevels labeled with numbers and letter which
indicate the shape of the orbital.

The letter s means a spherical orbital.

The letter p means a two-lobed orbital. The x, y, and z subscript tells along which of
the coordinate axes the two lobes lie.

Orbitals are potential spaces for atoms, and so all atoms contain all of the orbitals.
Electron Configurations

Longhand and Shorthand (Noble gas configuration)
o
Y, ZR, Fe, Po, S, Mf
Orbital Diagrams

Orbital is a box grouped by sublevel containing arrow(s) to represent electrons
o
Li, Sc, Si, F, Ar, Se
Periodic Trends

In a principal energy level that has d orbitals, the s orbital from the next level fills before the d orbitals
in the current level.

After lanthanum, a group of fourteen elements called the lanthanide series, or the lanthanides, occurs.
This series of elements corresponds to the filling of the seven 4f orbitals.

After actinum, a group of fourteen elements called the actinide series, or actinides, occurs. This series
corresponds to the filling of the seven 5f orbitals.

Except for helium, the group numbers indicate the sum of electrons in the ns and np orbitals in the
highest principal energy level that contains electrons (where n is the number that indicates a particular
principal energy level). These electrons are the valence electrons.
Bonding


Forces that hold groups of atoms together and make them function as a unit.
o
No simple, and yet complete, way to define this.
o
A bond will form if the energy of the aggregate is lower than that of the separated atoms.
Kinds of Bonding
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Ionic bonding: Ionic compounds result when a metal and a non-metal react and electrons are
shared.
o
Covelant Bonding: Electrons are shared by nuclei of two or more atoms.
o
Polar Covelant bonding: Unequal sharing of electrons between atoms in a molecule.

One atom is more electronegative than the other (a difference of at least .4) and
results in a charge separation in the bond, one side being partly negative and the
other slightly positive.
Ionization Energy

The energy required to remove an electron from a gaseous atom or ion.
o
*Core electrons are bound much more tightly than valence electrons.
o
Ionization energy increases from left to right and decreses from top to bottom. Why?

Electrons added to the same principal quantum level do not completely shield the
increasing nuclear charge caused by the added protons.

Electrons in the same principal quantum level are generally more strongly bound
from left to right on the periodic table.

The electrons being removed are, on average, farther from the nucleus as you go
down on the periodic table.
Electronegativity

The ability of an atom in a molecule to attract shared electrons to itself.

On the periodic tale, electronegativity generally increases across a period and decreases down a group.

The range of electronegativity is 7.0 for fluorine to 0.7 for cesium and francium (the least
electronegative).

The polarity of a bond depends on the difference between the electronegativity values of the atoms
forming the bond.
Lewis Dot Structures

Show how valence electrons are arranged among atoms in a molecule.
o
Most important rule: Octect/Duet rules must be follow (exceptions: boron, beryllium,
phosphorus).

Steps for writing Lewis Dot Structures
o
Sum the valence electrons from all the atoms in a bond.
o
Use a pair of electrons to form a bond between each pair of bound atoms.
o
Atoms usually have noble gas configurations. Arrange the remaining electrons to satisfy the
octet (duet) rule.

Single Bond: covelant bond in which 1 pair of electrons is shared by 2 atoms.

Double Bond: covelant bond in which 2 pairs of electrons are shared by 2 atoms.

Triple Bond: covelant bond in which 3 pairs of electrons are shared by 2 atoms.
o
Resonance

A molecule shows resonance when more than one Lewis structure can be drawn for
the molecule.
Dipole Moments

Property of a molecule whose charge distribution can be represented by a center of positive charge.
o
An arrow is used to represent a dipole moment. The arrow points to the negative charge center
with the tail of the arrow indicating the positive center of charge.
o
Represented using + and -, each referring to the charge of the atoms involved in the polar
covelant bond.
Molecular Structures

The three dimensional arrangement of the atoms in a molecule.
3/20/2013 3:20:00 AM
3/20/2013 3:20:00 AM