Electrons and Periodicity
Wow, that periodic table is useful.
Development of the Periodic Table
Mendeleev
In 1869, Dmitri Mendeleev published a table of the elements
organized by increasing atomic mass. He grouped elements
with similar chemical properties together. Mendeleev left
blank spaces in the table because there were no known
elements with the appropriate properties and masses.
Mendeleev was able to predict the physical and chemical
properties of many of the missing elements. Eventually these
elements were discovered and were found to have properties
similar to those predicted.
Mendeleev cont.
However, not all of the elements in Mendeleev's table were in
order of increasing atomic mass. Mendeleev arranged the
elements tellurium and iodine, and cobalt and nickel, out of
order by atomic mass so that they could be placed in the
groups with which they share chemical properties Mendeleev
assumed that the atomic masses for these elements had been
incorrectly determined. However, new mass measurements
simply confirmed the original masses.
Moseley
In the early 1900s, Henry Moseley used x-ray diffraction to
determine the number of protons in the nucleus of an atom.
Moseley showed that arranging the elements by increasing
atomic number was consistent with Mendeleev’s ordering of
the periodic table by properties rather than strictly by atomic
mass.
Seaborg
One of the most recent changes in the traditional periodic
table was made by Glenn Seaborg. In 1944, Seaborg
rearranged the structure of the periodic table based on
the properties of elements he discovered. He removed
thorium, protactinium, and uranium from the body of
the table. These elements became the beginning of the
actinide series. His new arrangements allowed him to
predict the properties of even more elements. His ideas
were later verified when the elements were produced
artificially. He created plutonium, americium, curium,
berkelium, and californium. The element seaborgium
was named for him in honor of his accomplishments.
Arrangement of the Modern Periodic
Table
 In the modern periodic table, the elements are organized into
groups (vertical columns) and periods (horizontal rows) in
order of increasing atomic number.
 Elements that have similar chemical properties are in the
same group.
 Elements in the periodic table are classified as metals,
nonmetals, or metalloids.
Properties of Metals
1.
2.
3.
4.
5.
6.
They are malleable, ductile, and have luster.
They are good conductors of heat and electricity.
They have relatively high densities.
They are solids at standard temperature and pressure
(STP), except for mercury, which is a liquid.
They have relatively high melting points, except for
mercury and gallium.
They don’t combine chemically with other metals. Metals
combine physically to create alloys. Examples of alloys
include brass, steel, and solder.
Properties of Nonmetals
 They are dull and brittle.
 They don’t conduct heat and electricity well.
 They have relatively low boiling and freezing points.
 They exist in all phases at STP, but most are gases.
Properties of Metalloids
 Metalloids have properties of both metals and nonmetals.
 They are semiconductors at temperatures high than room
temperature. This means that they have a conductivity
between that of a metal and that of a nonmetal, and can even
change energy input and output by using electrical forces.
 They are all solids at STP.
Special Groups of Elements
Alkali Metals
The alkali metals are found in group 1 (1A) of the periodic
table.
 They do not occur in nature as elements (in other words,
they’re always in a compound).
 They form ionic compounds.
 They are good conductors of heat and electricity, ductile,
malleable, and soft enough to be cut with a knife.
 They have a silvery luster, low density, and low melting
point.
 They are the most reactive metals.
Alkaline Earth Metals
The alkaline earth metals are found in group 2 (2A) of the
periodic table.
Alkaline earth metals have similar characteristics to alkali
metals, except that they are less reactive.
Transition Metals
The transition metals are found in groups 3 through 12. They
have similar characteristics to the other metals except that:
 They are usually harder and more brittle than the metals in
group 1 and 2.
 They often form colored compounds. (Think of the blue
copper sulfate we’ve worked with.)
Halogens
The halogens are found in group 17 (7A) of the periodic table.
 They are all nonmetals and occur in combined form in nature,
mainly as metal halides (an example is salt, sodium chloride).
 They exist in room temperature as gases (F2 and Cl2), a liquid
(Br2) and solids (I2 and At).
 They are the most reactive nonmetals. Fluorine is the most
reactive of all nonmetals (as anyone who has seen Breaking Bad
discovered).
Noble Gases
The noble gases are found in group 18 (8A) of the periodic
table.
 They are colorless and odorless.
 They have very low boiling and freezing points.
 They rarely combine with other elements and are considered
to be inert (nonreactive chemically).
Inner Transition Metals
The inner transition metals are the two rows at the bottom of
the periodic table. They include the lanthanide and actinide
series. Here are some characteristics of the inner transition
metals.
 They are very dense compared to most other metals.
 Many of the inner transition metals are radioactive.
 Many of the actinides are synthetic.
 They are sometimes called the “rare earth elements”.
Target Check
How many groups are there on the periodic table?
How many periods are there on the periodic table?
Use the periodic table to identify by name and symbol the elements that have
the following locations.
1.
2.
3.
1.
2.
Group 1, Period 4
Group 17, Period 3
Where are the metals located on the periodic table?
Where are the nonmetals located on the periodic table?
Where are the metalloids listed on the periodic table?
Classify each of the following elements as metals, nonmetals, or metalloids.
4.
5.
6.
7.
1.
2.
3.
4.
Silicon
Mercury
Carbon
Neon
Target Check cont.
Match each choice on the right with the correct term on the
left. An answer may be used more than once.
Element Classification
Property
1. Malleable and ductile
a. Metal
2. Brittle
3. Dull in appearance
b. Nonmetal
4. Shiny
5. Good conductors of heat and
electricity
c. Metalloid
6. Nonconductors of heat and
electricity
7. Semiconductors of electricity
Target Check cont.
Identify the classes of elements to which each of the following
elements belongs. Each element is described by more than
one term. Choose from the following classes: metal,
nonmetal, metalloid; alkali metal, alkaline earth metal,
transition metal, inner transition metal (lanthanide or actinide),
halogen, noble gas; solid, liquid, gas (at room temperature).
Example: argon is a nonmetal, noble gas, gas
1.
2.
3.
4.
5.
Uranium
Barium
Selenium
Boron
Chlorine
Electron Arrangement in Atoms
The Bohr Model
In 1913, Niels Bohr came up with a new atomic model. He proposed
that electrons are arranged in concentric circular paths, or orbits,
around the nucleus. Bohr’s model, patterned after the motions of
the planets around the sun, is often referred to as the planetary
model. According to the Bohr model of the atom, electrons exist
on fixed energy levels and cannot exist between energy levels. To
more from one energy level to another, an electron must gain or
lose just the right amount of energy.
Bohr’s model of the atom was consistent with the emission spectrum
produced by the hydrogen atom, but the model does not work for
more complicated atoms. Although the Bohr model is not the
current model of the atom, his model is still commonly used to
help explain chemical behavior and periodic trends.
Drawing Bohr Models for Atoms
The electrons are drawn on energy levels.
1.
a)
b)
c)
d)
2.
The first energy level can hold 2 electrons.
The second energy level can hold 8 electrons.
The third energy level can hold 18 electrons.
The fourth energy level can hold 32 electrons.
Although the third and fourth energy level can hold more
than 8 electrons, an outer energy level must never have
more than eight electrons on it. This will affect the
elements with atomic numbers greater than 8.
Example 1
Let’s look at an example of the Bohr model for the oxygen-16
atom.
 How many protons and electrons does an oxygen atom have?
 How many neutrons does an oxygen atom have?
Two electrons will go on the first energy level and six electrons
will go on the second energy level.
An easier way of
writing this is 2.6
Example 2
Let’s look at an example of the Bohr model for aluminum-27
atom.
How many protons and electrons does an aluminum atom have?
How many neutrons?
Two electrons will go on the first energy level. Eight electrons
will go on the second energy level and three electrons will go
on the third energy level.
An easier way of writing this is 2.8.3
Example 3
Now let’s look at an example of a Bohr model for calcium-40. Calcium has more
than 18 e-.
How many protons and electrons does a calcium atom have?
How many neutrons?
Two electrons will go on the first energy level. Eight electrons will go on the
second energy level. Although the third energy level can hold up to 18
electrons, only 8 electrons will be placed in it. Remember: only eight
electrons can be on an outer energy level. The remaining electrons go
on the fourth energy level.
An easier way of writing this is 2.8.8.2.
Bohr Model Note
If you needed to add more electrons for a higher number
element, you would put them on the third level until you
reached 18. For example, the element zinc has 30 electrons.
There would be 2 electrons on the first energy level, 8
electrons on the second energy level, 18 electrons on the
third energy level, and then 2 electrons the fourth energy
level. The shorthand way of writing this would be 2.8.18.2.
Target Check
Draw a Bohr model for chlorine-35. Have your teacher check
your answer before you continue.
Target Check cont.
Draw a Bohr model for titanium-48. Have your teacher check
your answer before you continue.
Beyond Bohr
Since Bohr’s model of the atom only worked for hydrogen,
other methods needed to be investigated in order to
determine the structure of the atom with regards to the
electrons. One proposal was to examine the electromagnetic
radiation being emitted or absorbed by a substance. This area
of chemistry is referred to as spectroscopy. In order to study
the electromagnetic radiation of substance, we will be
treating this energy as waves.
Parts of a Wave
With regards to electrons as waves, an electron is pushed in one
direction and then in the opposite direction. It will then
continue to do so over and over again.
Frequency is the number of wave cycles per second that pass
through a given point in units of Hertz (Hz). (Note: Hz=1/s)
Wavelength (symbolized by the Greek letter λ) is the peak-topeak distance measured in meters (m) or also typically in
nanometers (nm) for visible light. (Note: 1 nm=109 nm)
Amplitude is the height of the wave above the center line.
The relationship involving a wave is
frequency x wavelength = speed of light or
f λ=c
Where c=2.998 x 108 m/s, wavelength is in meters, and
frequency is in Hertz.
Example
Suppose we want to find the wavelength of blue light of
frequency 6.4 x 1014 Hz. Then,
Planck’s Constant
Light can also be described as made up of particles, or photons.
German physicist Max Planck suggested that each photon of
light exchanges a certain amount of energy with its
surroundings. The amount of energy depends on the
frequency of the light according to the equation:
E=hf
Where E is the amount of energy released in Joules (J), and h is
Planck’s constant, 6.626 x 10-34 J•s. This is the amount of
energy produced per photon.
Example
What is the energy released of a photon of blue light of
frequency 6.4 x 1014 Hz?
Target Check
 Calculate the wavelength of a green light at a traffic signal
assuming the frequency is 5.75 x 1014 Hz.
 A green line from the emission spectrum of hydrogen has a
wavelength of 486 nm. Calculate the energy given off by each
photon of this green line.
Spectroscopy
When white light is passed through a prism, a continuous
spectrum of light results. However, when the light emitted
from excited atoms of an element are passed through a
prism, only certain spectral lines appear as opposed to the
entire spectrum. Each element produces a unique, brightline atomic spectrum that can be used to identify that
element.
Quantum Mechanical Model
The current model of the atom is known as the quantum
mechanical model of the atom, which is a mathematically
based model. Like the Bohr model, the quantum mechanical
model of the atom restricts the energy of electrons to certain
values, but does not define the exact path an electron takes.
The quantum mechanical model pictures energy levels as
regions of space where there is a high probability of finding
an electron. Each energy level is actually made from a certain
number of sublevels. Each sublevel has one or more atomic
orbitals within a specific 3-dimensional shape. An atomic
orbital is a region in space where there is a high probability of
finding an electron.
Target Check
 How are the Bohr model and the quantum mechanical model
alike?
 How are energy levels depicted according to the quantum
mechanical model?
Electron Configuration and Orbital
Diagrams
Electron configuration and orbital diagrams can be written to help
illustrate the arrangement of the electrons around the nucleus of an
atom. In an orbital diagram, the orbital is represented by a box. The
electrons are represented by arrows. Here is an example of an orbital
diagram and an electron configuration for the element hydrogen.
The 1 in front represents the energy level.
The s represents the type of sublevel.
Hydrogen’s single electron is represented by a single arrow in the orbital
diagram and the superscript 1 in the electron configuration.
Sublevels
There are four types of sublevels. They are s, p, d, and f. Each
sublevel is made of atomic orbitals.
 The s sublevel contains one atomic orbital (one box).
 The p sublevel contains three atomic orbitals (three boxes).
 The d sublevel contains five atomic orbitals (five boxes).
 The f sublevel contains seven atomic orbitals (seven boxes).
They can be compared to an apartment building. The energy level
tells you what floor of the apartment building the electron lives
on. The sublevel tells you what type of apartment the electron
lives in—1, 3, 5, or 7 bedroom. The superscripts tell you how
many electrons live in the apartment. The arrows tell you how
many electrons live in each room of the apartment.
Pauli Exclusion Principle
There are three rules governing the filling of atomic orbitals by electrons within
the principal energy levels. These rules are the Pauli exclusion principle, the
aufbau principle, and Hund’s rule.
The Pauli exclusion principle states that only 2e- can occupy an atomic orbital and
they must have opposite spin.
a)
b)
c)
d)
The s sublevel contains only 1 atomic orbital. How many electrons can
occupy the s sublevel?
The p sublevel contains three atomic orbitals. How many electrons can
occupy a p sublevel?
The d sublevel contains five atomic orbitals. How many electrons can occupy
a d sublevel?
The f sublevel contains seven atomic orbitals. How many electrons can
occupy a f sublevel?
The aufbau principle
 The aufbau principle states that electrons enter orbitals of
lowest energy first.
 Here is a diagram illustrating the order in which the
electrons enter the orbitals.
 Here is a list of the atomic orbitals in order of increasing
energy through 8s.
Hund’s Rule
Hund’s rule states that when electrons occupy orbitals of equal
energy, one electron occupies each orbital until all orbitals
contain one electron with parallel spins. In other words, for
each sublevel, there must be one electron placed in each box
before you “double up” the electrons.
Examples
 Orbital diagram and electron configuration for Magnesium.
Magnesium has 12e-.
 Orbital diagram and electron configuration for Arsenic.
Arsenic has 33e-.
Target Check
 Draw an orbital diagram and write an electron configuration
for silicon. How many electrons occupy each level?
 Draw an orbital diagram and write an electron configuration
for selenium. How many electrons occupy each level? How
many half-filled orbitals are there?
Photoelectric Effect
 The electron configurations and orbital diagrams that you
have been drawing represent the atom in its ground state.
When the electrons absorb energy, they move to an excited
state. Here is an example of an electron configuration for an
atom of neon in an excited state: 1s22s22p64s1
 Notice that the last electron is in a higher energy level than it
should be. The electron will not remain in this higher energy
level. It will return to its normal ground state. When the
electron falls back down to its ground state, it emits energy
in the form of colored light. An example of this is neon signs
and fireworks.
Target Check
 Identify each of the following elements given the
configurations below.
 1s22s23p3
 1s22s22p63s23p64s23d6
 For each of the following electron configurations of neutral
atoms, determine if the configuration as written is the
ground state, an excited state, or if it is an impossible
configuration.
 1s22s22p2
 1s22s22p7
 1s22s22p64s1
Using the Periodic Table to Write
Electron Configurations
With a few exceptions, correct electron structures for atoms
can be derived from examining the element’s position on the
periodic table.
Examples
 Chlorine:




Reading across the first row it is 1s2.
Reading across the second row it is 2s22p6.
Reading across the third row it is 3s23p6.
The electron configuration for chlorine 1s22s22p63s23p6.
 Cadmium:






Reading across the first row it is 1s2.
Reading across the second row it is 2s22p6.
Reading across the third row it is 3s23p6.
Reading across the fourth row it is 4s23d104p6.
Reading across the fifth row it is 5s24d10.
The electron configuration for cadmium is
1s22s23s23p64s23d104p65s24d10.
Target Check
 Use the periodic table to write electron configurations for
the following elements.
1. Oxygen
2. Potassium
3. Zirconium
Noble Gas Notation
A shorter way of writing electron configurations is to use a
noble-gas notation. Study the examples below to see if you
can determine how to write noble-gas notations for
elements.
1. The noble-gas notation for sodium would be [Ne]3s1.
2. The noble-gas notation for potassium would be [Ar]3s1.
3. The noble-gas notation for iodine would be [Kr]5s24d105p5.
4. The noble-gas notation for argon would be [Ne]3s23p6.
Target Check
 Write the noble-gas notations for each of the following
elements.
1. Aluminum
2. Zinc
3. Krypton
4. Rubidium
Valence Electrons
The outermost electrons play the largest role in determining
the chemical properties of the elements. These are the
electrons that can be gained, lost or shared in the formation
of chemical compounds. These electrons are known as
valence electrons. The valence electrons for the representation
elements (elements in groups 1, 2, and 13-18) are the electrons
filling the s and p sublevels of the highest occupied energy level.
Target Check
 Write noble-gas notations to determine the number of
valence electrons for each of the following elements. The first
one has been done for you.
Sodium
Magnesium
Aluminum
Silicon
Phosphorus
Sulfur
Chlorine
Argon
Group 1
[Ne]3s1
1
Target Check
Compare the group number to the number of valence
electrons. Describe an easier way to determine the number
of valence electrons.
Octet Rule
The octet rule states that elements will gain, lose, or share electrons
in order to obtain 8 electrons in their outermost energy level. The
noble gases (except for helium) have 8 electrons in their
outermost energy level. Noble gas configurations are considered
to be very stable. Since other elements tend to favor stability as
well, their atoms often gain, lose, or share electrons to obtain a
configuration matching that of a noble gas.
When an element loses an electron, it becomes a positive ion
(cation). Metals tend to lose electrons and form positive ions
(cations). When an element gains an electron, it becomes a
negative ion (anion). Nonmetals tend to gain electrons and form
negative ions (anions).
Octet Rule cont.
Using the Octet Rule and ending electron configurations,
predict the charge on the stable ion of each of the following
elements. The first one has been done for you.
Iodine
Calcium
Phosphorus
[Kr]5s24d105p5
-1
Target Check
 Complete the following table.
Group
Number
1
2
13
14
15
16
17
18
Number of
Valence
Electrons
Charge of
Common Ion
Target Check cont.
 Write the symbol for the ion formed by each of the following
elements and indicate the number of electrons gained or lost.
The first one has been done for you.
Calcium
Phosphorous
Potassium
Ca2+
Loses 2e-
Isoelectronic
 The noble-gas configuration for Chlorine is [Ne]3s23p5. According




to the octet rule, chlorine will gain one electron in order to obtain
a full outermost energy level. When chlorine gains one electron, it
will become the chloride ion (Cl-). The electron configuration for
the chloride ion is [Ne]3s23p6. That is the same as the electron
configuration for argon. When two species have the same electron
configuration, they are said to be isoelectronic.
What happens when lithium becomes an ion?
Write an electron configuration for the lithium ion (Li+).
Which neutral element has the same electron configuration?
Name five ions that are isoelectronic with Neon.
Periodicity
 When the elements are arranged in order of increasing
atomic number, there is a periodic reoccurrence of
properties that leads to the group of elements in the periodic
table. This periodic recurrence is known as periodicity.
 The position of the element in the periodic table can also be
used to compare periodic trends in atomic radii,
electronegativity, ionization energy, and ionic radii.
Nuclear Charge and the Shielding
Effect
All the periodic trends can be understood in terms of three basic rules.
1. Electrons are attracted to the protons in the nucleus of an atom.
a) The closer an electron is to the nucleus, the more strongly it is attracted.
b) The more protons in the nucleus, the more strongly an electron is attracted
to the nucleus. This is known as nuclear charge or force.
2. Electrons are repelled by other electrons in an atom. So if other electrons are
between a valence electron and the nucleus, the valence electron will be less
attracted to the nucleus. The tendency for the electrons in the inner energy
level to block the attraction of the nucleus for the valence electrons is known as
the shielding effect.
3. Completed p sublevels are very stable. Atoms prefer to add or subtract valence
electrons to create completed p sublevels if possible.
Nuclear charge plays an important role in determining periodic trends. The
shielding effect plays an important role in determining group trends.
Atomic Radii
The atomic radius is half the distance between nuclei in two adjacent
atoms.
 How do the values for atomic radii change as you move from left to
right across a period?
Moving from left to right across a period, the atomic number increases.This means
that more protons are added to the nucleus , so the valence electrons are more
strongly attracted to the nucleus.
 How do the values for atomic radii change from top to bottom with a
group?
Each successive element down a group has its valence electrons in a principal energy
level further away from the nucleus.This means that there is a greater distance
between the nucleus and the valence electrons.The combined effect of a greater
distance from the nucleus and increased shielding results in a larger atomic
radius.
Target Check
Atomic radii do what as they move from left to right across
a period? What do they do when they go down a group?
2. Which element would have the smallest atomic radii?
3. Which element would have the largest atomic radii?
4. Which atom in each pair would have the larger atomic
radii?
1.
a)
b)
c)
Li or Cs?
Li or F?
K or Br?
Ionization Energy
Electrons are attracted to the nucleus of an atom, so it takes energy to remove an
electron. The energy required to remove an electron from an atom is called the
first ionization energy. Once an electron has been removed, the atom becomes a
positively charged ion. The energy required to remove the next electron from
the ion is the second ionization energy, and so on.
 How do the values for ionization energy change as you move from left to right
across a period?
As you go from left to right across a period, the nuclear charge increases. Because
there are more protons in the nucleus and the energy level of the valence
electrons remains constant, the valence electrons are more strongly attracted to
the nucleus. This increases the amount of energy required to remove them.
 How do the values for ionization energy change from top to bottom within a
group?
As you go from top to bottom within a group, the shielding effect increases as
more energy levels are added. The valence electrons are not as strongly
attracted to the nucleus. This decreases the amount of energy required to
remove them.
Comparing Successive Ionization
Energies
 The energy required to remove an electron increases as more
electrons are removed.
 Write the electron notification for a sodium atom. For Na,
the second izionization energy is much larger than the first.
The first electron is being removed from the s sublevel. The
second electron would have to be removed from a completed
p sublevel. A completed p sublevel is also more stable.
 Write the electron configuration for a magnesium atom. The
first and second ionization energies are comparable, but the
third ionization energy is much larger than the second.
Explain why this is so.
Target Check
 How does ionization energy change in periods and groups?
 Which group of elements would have the highest first
ionization energy?
 Which group of elements would have the lowest first
ionization energy?
 Which atom in each pair would have the larger first
ionization energy?




Ca vs Br
Ca vs Ba
Na vs Cs
Na vs P
Ionic Radii
 The ionic radius is the radius of a cation or anion. When the
atom loses aor gains electrons, the resulting ion changes in
size from the original atom.
 Metals tend to lose electrons and form cations.
 Nonmetals tend to gain electrons and form anions. The names
of monatomic anions end in –ide. Example: sulfide, phosphide,
fluoride.
Cations
 How does the ionic radius of cations compare to their respective atomic radius?
 Let’s look more closely at the sodium atom versus the sodium ion. How many
electrons does sodium lose when it becomes an ion?
 Generally, when electrons are removed from an atom to form a cation, the
outer energy level is lost, making the cation smaller than the atom. The proton
to electron ratio also changes. In a sodium atom the proton to electron ratio is
11p+ to 11e-. In a sodium ion the proton to electron ratio is 11p+ to 10e-.
Cations are smaller than the neutral atom from which they came due to a higher
proton to electron ratio.
Anion
 How does the ionic radius of anions compare to their respective
atomic radius?
 Let’s look more closely at the nitrogen atom versus the nitride
ion. How many electrons does nitrogen gain when it becomes an
ion?
 The nitrogen atom as a proton to electron ratio of 7p+ to 7e-. The
nitride ion has a proton to electron ratio of 7p+ to 10e-. Anions
are larger than the neutral atom from which they came due to the
lower proton to electron ratio.
Target Check
 The two ions K+ and Ca2+ both have 18 electrons surrounding
their nucleus. Which would expect to have the smaller ionic
radius? Why?
 Cations are always (what?) than the neutral atom from which they
came.
 Anions are always (what?) than the neutral atom from which they
came.
 Which particle in each pair would be larger?




Lithium atom vs lithium ion
Fluorine atom vs fluorine ion
Sodium atom vs magnesium ion
Oxide ion vs fluoride ion
Electronegativity
 Electronegativity refers to the tendency for an atom to
attract electrons to itself when it is chemically combined
with another element.
 How do the values for electronegativity change as you move
from left to right across a period?
 How do the values for electronegativity change from top to
bottom within a group?
 As you go from top to bottom within a group, the shielding
effect increases. The valence electrons are not as strongly
attracted to the nucleus.
Target Check
 The noble gases do not generally have electronegativity




values. Why do you think this is so?
Which group of elements has the highest electronegativity
values?
Which group of elements has the lowest electronegativity
values?
Arrange the elements gallium, bromine, and calcium in order
of increasing electronegativity.
Arrange the elements iodine, chlorine, and bromine in order
of increasing electronegativity.
Summary of Periodic Trends