Chapter 3 Notes for Chemistry Exam

3.1 – The Periodic Law and Table:
Scientists were trying to find order or some sort of pattern to explain the chemistry of
the elements.
Periodic Law – When arranged in increasing atomic number, elements with similar
chemical properties occur at regular (periodic) intervals.
Dmitri Mendeleev was the first to observe this periodic behavior and he organized
the elements into a table with vertical columns or groups and horizontal rows or
periods.
Group or Family – A vertical column of elements in the periodic table having
similar chemical properties. Scientists are still fighting over how to label the groups,
but we will number them from left to right.
Period – A horizontal row of elements across the periodic table. Periods are
numbered from top to bottom.
Each element in the periodic table belongs to both a group and a period.

3.2 – Electronic Arrangements in Atoms:
So why do elements have similar properties at regular periodic intervals?
The answer lies in their electrons.
Model A Model B
Ernest Rutherford proposed a planetary model for the atom where e-s orbit the
nucleus of an atom like planets orbit the sun.
Niels Bohr proposed that e-s can only orbit the nucleus at fixed distances from the
nucleus, and therefore e-s can only have specific energies.
Bohr also proposed that e-s can only jump to orbits of different distances from the
nucleus when energy is absorbed or released.
Absorbing energy causes an e- to jump farther away from nucleus to higher energy
and less stable orbits.
Releasing energy causes the e- to fall back
down closer to the nucleus to lower energy
and more stable orbits.
Erwin Schrodinger further refined this
model stating that the precise paths of e- can’t
be determined accurately (Like Bohr
thought). According to Schrodinger, the
location and energy of electrons around the
nucleus can be specified using 3 terms: shell,
subshell, and orbital. This model is called
the quantum mechanical model.
Shell – A location and energy of e-s around a
nucleus that is designated by a value for n,
where n = 1, 2, 3.
The lowest energy shell is assigned n = 1, and the next lowest is n = 2, etc.
The higher the value of n, the farther away from the nucleus the e-s are and the more
energy they have.
Subshell – A sub-compartment of a shell designated by the letters: s, p, d, f.
Subshells are identified using shell # (value for n) and the subshell letter (s, p, d, f).
Example: for n = 3, there are 3 subshells designated 3s, 3p, and 3d
All e-s within a specific subshell have the same energy (they are the same distance
from the nucleus).
Subshells are further divided into 1 or more atomic orbitals.
Orbital – A specific volume of space around the nucleus of an atom where e-s of the
same energy move. All the orbitals of a subshell have the same value of n.
The volume of space around the nucleus of each subshell is different.
All e-s in the same orbital have the same energy, regardless of the orbital they are in.
Example: An e- in a 4d orbital has the same energy as any other e- in another one of
the five 4d orbitals.
All s subshells consist of 1 singular orbital.
All p subshells consist of 3 orbitals
All d subshells consist of 5 orbitals
All f subshells consist of 7 orbitals
Each orbital within a subshell can contain a
maximum of 2 e-.
Energy Diagram of Atomic Shells, Subshell, and Orbitals
Practice Problems:
What is the e- configuration for Nitrogen?
What is the e- configuration for Calcium?
What is the e- configuration for Nickel?
What is the e- configuration for Chlorine?

3.3 The Shell Model and Chemical Properties:
The arrangement of e-s into orbitals, subshells, and shells provides and explanation
for the similarities in chemical properties of various elements.
All the elements in a specific group (column) have the same number of electrons in
their outer shell.
This shell is called the valence shell.
Valence Shell – The outermost shell of an element that contains the highest energy
electrons in that atom.
Elements with the same number of e-s in their valence shell display similar properties.
The number of e-s in a valence shell is identical the Roman numeral above the group
on the periodic table.
The n value for the valence shell increases by one as you go down the periodic table.
Practice Problem:
How many valence e-s does oxygen have?
How many valence e-s does magnesium have?
How many valence e-s does Yttrium need to lose or gain to obtain a full valence?
How many valence e-s does Sulfur need to lose or gain to obtain a full valence?
How many valence e-s does Fluorine need to lose or gain to obtain a full valence?

3.4 Electronic Configurations:
Electronic Configurations – The detailed arrangement of e-s indicated by a specific
notation. (1s2, 2s2, 2p6, 3s2, 3p6, 3d10)
Hund’s Rule – Electrons will always fill any available empty orbitals of the same
energy before pairing up to share an orbital. Think about people sitting on a bus.
Pauli Exclusion Principle – Only e- spinning in opposite directions can
simultaneously occupy the same orbital.
Order for filling atomic orbitals with e-s.
Helpful hint for determining e- configuration:
Noble Gas Configuration – An electronic configuration consisting of completely
filled s and p outermost subshells. 8 valence e-s, very stable, essentially inert. Far
Right Group of elements are called the “Noble gases”, they do not form bonds with
any other atoms. They are completely inert. Given the choice, atoms always want to
gain or lose e-s to obtain a full octet of 8 valence e-s.

3.5 Another Look at the Periodic Table:
Elements in the same family have the same # of valence e-s. (F, Cl, Br, I, & At each
have 7). As a result they all have similar chemical properties.
Distinguishing Electron – the last and highest-energy electron in an element.
Distinguishing e-s are categorized by the type of subshell (s, p, d, or f) they are found
in. (See figure 3.9)
Noble gases – Elements in the far right side of the periodic table. All are gases at
room temp and all are inert. They all have completely filled valence orbitals.
The elements can also be classified as: metals, nonmetals, and metalloids.
Metals – Found in the left two thirds of the periodic table. They have the following
properties:
• High thermal conductivity – transmit heat well
• High electrical conductivity – transmit electricity well
• Ductility – can be stretched in wires
• Malleability – can be hammered into thins sheets
• Metallic luster – shiny, “metallic” appearance
Nonmetals – Found in the right one third of the PT. Often brittle, powdery solids or
gases and have properties opposite those of metals.
Metalloids – Have properties between those of metals and nonmetals. They include:
B, Si, Ge, As, Sb, Te, At. They form a “staircase” separating metals from nonmetals
on the PT. (See figure 3.12)
3.6 Properties and Trends within the Periodic Table:
Elements become less metallic from left to right across a period and from bottom to
top in a group.
Elements decrease in size from left to right across a period and increase in size from
top to bottom down a group.
Each added proton results in a stronger attractive force between the nucleus and the
electrons.
So, each added proton pulls the e- cloud closer to the nucleus, shrinking the atom as
you move from left to right across the PT.
Scale Drawings and Atomic Radii of Various
Atoms Across the Periodic Table
Ionization energy – the energy
required to remove an electron from a
neutral atom.
Na
Na+ + e-
The higher the value of the ionization
energy, the harder it is to remove an e-
.
Ionization energy increases from left
to right across a period and decreases
in size from top to bottom down a
group.
As the size of the atom increases, the
e-s are farther away from the positive,
attractive forces in the nucleus, so they are held more loosely than e-s closer to the
nucleus and therefore easier to remove.
Note: When sodium loses 1 e-, the remaining valence e-s have a full octet and the
sodium cation has the same e- configuration as Neon (a noble gas). The Na+ cation is
very stable and forms many commonly used salts, such as NaCl (table salt) and
NaHCO3 (baking soda).
The tendency of metals to lose their e-s (and conduct electricity) is because by giving
a way a couple of e-s to an atom that is a few e-s short of its own octet, the metal can
have a completely filled valence and increase its stability.
This tendency to lose e-s results in metals holding their e-s less tightly, allowing e-s
(and therefore electricity) to flow easily from one atom of the metal to the other. It’s
almost like hot potato, no one atom wants the e-s, because it would rather lose it and
have a full octet, so it passes it down the line to the next atom and BOOM, electricity.
Electronegativity (EN) – the attraction an element has for e-s, the higher the EN, the
more e-s are pulled towards the atom.
Electronegativity increases from left to right across a period and decreases from top
to bottom down a group. Fluorine is the most electronegative atom. Francium is the
least electronegative.