Chapter 10 Notes
How are electrons arranged around the nucleus of an atom?
Misconception: electrons fly in circular patterns around the nucleus.
What really happens: Electrons form a cloud around the nucleus and spin on their axis.
Analogy: Think of the nucleus as the planet and the electrons make up the atmosphere.
They essentially make up a cloud around the nucleus.

This cloud is made up of energy levels.
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Each level is made up sublevels.
o
1st energy level has 1 sublevel (“s” sublevel)
o
2nd energy level has 2 sublevels (“s” and “p” sublevels)
o
3rd energy level has 3 sublevels (“s”, “p”, and “d” sublevels)
o
4th energy level has 4 sublevels (“s”, “p”, “d” and “f” sublevels)
http://www.youtube.com/watch?v=K-jNgq16jEY
Atomic orbitals – areas where electrons are likely to be found.
s orbital –only 1 s orbital per
sublevel
p orbital –3 p orbitals per
sublevel
d orbital – 5 d orbitals per
sublevel
f orbital – 7 f orbitals per
sublevel
The organization of the periodic table makes it EASY to determine the orbitals of
the electrons for each atom.
Rules for writing orbital notations
1. When writing orbital notations, use one arrow to represent each electron.
2. Electrons must enter the lowest energy sublevel possible before moving to a higher
energy sublevel (Aufbau principle)
3. an atomic orbital may hold at most two electrons
a. Electrons within the same orbital have opposite spins. (Pauli exclusion
principle)
4. one electron must be put in each orbital of a sublevel before any one orbital can
have two electrons in it. (Hund’s rule)
5. Even if you don’t have enough electrons to fill each orbital of a sublevel, you must
still show that those orbitals exist.
6. The total number of arrows (electrons) must be equal to the atomic # for each
element.
Write the orbital notation for oxygen.
Step 1: How many electrons does oxygen have?
Step 2: What orbital level range will the electrons occupy?
Step 3: Place one electron in each sublevel before adding a second arrow to each
sublevel.
Write the orbital notation for aluminum.
Step 1: How many electrons does aluminum have?
Step 2: What orbital level range will the electrons occupy?
Step 3: Place one electron in each sublevel before adding a second arrow to each
sublevel.
Electron Configurations – the way in which electrons are arranged around the
nucleus of an atom. Each configuration has 3 parts:
1s2
“1” represents the energy level
“s” represents the sublevel
“2” represents the number of electrons in that sublevel
* The total of superscripts is equal to the atomic number for the element.
Which element is represented by the following electron configuration:
1. 1s22s22p63s23p6
2.
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1
Write the electron configuration for the following elements:
1. Sulfur
2. Gallium
3. Thorium
4. Platinum
Noble gas configurations are used as a shorthand for long electron configurations.

Find the noble gas before the element you are writing the configuration for, put it
in brackets, and then start with the next s sublevel to fill out the rest of the
configuration.
Write the noble gas configuration for the following elements:
1. Sulfur
2. Iron
3. Gallium
4. Thorium
5. Platinum
What element is represented by the following noble gas configuration:
1. [Kr]5s24d105p2
2. [Ar]4s2
3. [Xe]6s24f145d6
Trends in the periodic table
Lithium –
Sodium –
Potassium –
Rubidium –
Fluorine –
Chlorine –
BromineIodine-
Columns are called ________or _________
Rows are called ___________

Each period shows the pattern of properties repeated in the next period
Atomic Radius Trend
Atomic radius – distance from the nucleus of an atom to its valence electrons. The
radius tells that size of the atom.

Moving from left to right across a period, atomic radius decreases.

Moving from top to bottom of a family, atomic radius increases

valence shell farther from nucleus
Why does the atomic radius decreases as you move left to right across period

adding electrons to same valence shell

valence shell held closer because more protons in nucleus
Why does the atomic radius increase as you move from top to bottom of a family?

The valence electrons get further and further from the nucleus because you are
adding more energy levels. Therefore the radius of the atom increases.
Example: Put the following elements in order of increasing atomic radius:
Zn, Sc, Se, K, Cs, O
Example: Put the following elements in order of decreasing atomic radius:
F, Cd, Ba, Ge, W, Cl
Ionization Energy Trend

Ionization energy – the energy required to remove an electron from an atom (1st
ionization energy).

Removing an electron creates a charge imbalance, so a cation (positive ion) is
formed.

2nd Ionization energy – the energy required to remove two electrons from an atom.

Moving from left to right across a period, ionization energy increases. Why?
Within the same energy level (period) electrons experience an increasing pull from
the nucleus, so it takes more energy to remove them.

Moving down a group, ionization energy decreases. Why?
The valence electrons feel less and less pull from the nucleus as they get further
from the nucleus.

As electrons are removed, ionization energy increases gradually until an energy level
is empty, then it makes a big jump.

Pulling an electron off of a alkali metal (Group 1 elements) is easy. Trying to pull an
electron off of a noble gas (Group 18 elements) takes much more energy.

Example: Put the following elements in order of increasing ionization energy:
Sr, Cr, As, S, Rb, Cu

Example: Put the following elements in order of decreasing ionization energy:
O, V, K, P, Ga, Fr

Which of the following elements ( Na, Al, Ne, Mg, Si) will have a very large second
ionization energy?

Third ionization energy?
Ionic Radius Trend
Ionic radius – similar to atomic radius but it is the radius for an ion instead of an
atom.

Positive ions are always smaller than their neutral atoms, and negative ions are
always larger than their neutral atoms.

As you go down a group, ionic radius increases.

As you go from left to right across a period, positive ions decrease in size.

Negative ions also decrease as you go across a period, but they start off being
much larger than positive ions.

Put the following ions in order of increasing ionic radius:

Hint: If all of the ions have the same number of electrons, than the one with the
highest number of protons has the smallest radius.
Na+1, Al+3, N-3, F-1, O-2, Mg+2
Electronegativity Trend
Electronegativity – how strongly the nucleus of an atom attracts the electrons of
other atoms in a bond.

Nonmetals tend to gain electrons when they form bonds, and have higher
electronegativities than metals, which tend to lose electrons, when they form
bonds.

Moving from left to right across a period, electronegativity increases.

Moving down a group, electronegativity decreases.
1. Put the following elements in order of increasing electronegativity:
Fe, Si, O, Ba, Ca, Cs
2. Put the following elements in order of decreasing electronegativity:
Se, F, Ag, Pt, Fr, Sb