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Atoms and the Periodic Table
Interesting Links:






WebElements (online Periodic Table at http://www.webelements.com/webelements/scholar/index.html)
It's Elemental: The Periodic Table (http://pubs.acs.org/cen/80th/elements.html)
Exact Masses of the Elements and Isotopic Abundances (http://www.sisweb.com/referenc/source/exactmas.htm)
Spectra of Gas Discharges (http://astro.u-strasbg.fr/%7Ekoppen/discharge/)
Atomic Spectra (http://www.colorado.edu/physics/2000/quantumzone/lines2.html)
Interactive Chart of Nuclides (http://www.nndc.bnl.gov/chart/)
Chapter Objectives:
After completing this chapter, you should, at a minimum, be able to do the following. This information can be found in my
lecture notes for this and other chapters and also in your text.
1.
2.
Correctly answer all of the questions suggested above and in the quiz for this chapter.
Define basic terms such as proton, neutron, electron, ion, isotope, atomic number, atomic mass, atomic weight, atomic mass unit (amu),
cation, anion, octet rule, quantum number, shell, subshell, orbital, ground state, excited state, electron configuration, full electron
configurations, Noble gas (or valence shell) electron configurations, electron spin diagrams, inner shell electrons, outer shell electrons,
valence electrons, atomic size, metallic character, ionization energy, electron affinity, electronegativity.
3. Be familiar with Dalton's atomic theory.
4. Be familiar with the fundamental properties of protons, neutrons, and electrons.
5. Assign the correct number of protons, neutrons, and electrons to any isotope or ion.
6. Correctly calculate the atomic weight of any element given the masses and abundances of its isotopes.
7. Predict the number of electrons an element will typically gain or lose using the octet rule and its position on the periodic table.
8. Be able to assign full electron configurations, Noble gas (or valence shell) electron configurations, and electron spin diagrams to atom or
ion.
9. Understand the relationship between an atom's electron configuration and its position in the periodic table.
10. Predict how many inner shell, outer shell, and valence electrons an atom has.
11. Describe in general terms the periodic trends.
A brief history of atomic theory
In the 5th Century B.C. Leucippus and his student Democritus postulated that all matter is made of invisible,
indivisible particles called "atomos," which in Greek means "indivisible." At roughly the same time another
Greek philosopher, Empedocles, proposed that everything is made up of four basic elements: earth, air, water,
and fire. This notion was later popularized by the work of Aristotle (abt. 200 B.C.), and the ideas of Democritus
were largely forgotten. In the 1st Century B.C. the rejected and mostly forgotten ideas of Democritus were
popularized by the Roman poet Lucretius in his work "De Rerum Naturae."
By the time of the Renaissance the original writings of Democritus no longer existed, but his ideas were
rediscovered in the late 16th Century as scholars read Lucretius. By the 17th Century physicists such as Robert
Boyle, Robert Hooke, and Isaac Newton began to reject the Aristotlean notion of matter, and to suggest that the
physical properties of matter might be better explained by forces acting upon a ultimate building block of
matter. In his great work "Optiks" Newton wrote: "it seems probable to me that God in the beginning formed
matter in solid, massy, hard, impenetrable, moveable particles. . . ." .
Between 1803-1807 the English scientist John Dalton proposed a theory that still serves as the basis of our
understanding of matter. He proposed:
1.
2.
3.
4.
5.
Each element is composed of extremely small particles called atoms
All of the atoms in an element are identical
The atoms of an element are different from the atoms in any other element
The atoms of two or more elements can combine to form compounds
The law of constant composition: compounds always have the same ratios of atoms regardless of where
the compound is found
6. Atoms are not created or destroyed in chemical reactions
While we now take these ideas as a basic statement of fact, they were not immediately accepted by the
scientific community. Some were argued for nearly 100 years. But in the end, they were proved and accepted. It
is nothing short of amazing to realize that our notions of matter were initially postulated two hundred years ago.
While people began to accept the notion of atoms as the basic building blocks of matter during the 19th
Century, they still did not know what they were.
In 1897 J. J. Thomson discovered the electron. In 1909 Millikan discovered both the charge and mass of
electrons.
In 1896, at roughly the same time as Thomson was working on his electrons, Henri Becquerel discovered the
"radioactivity" of a uranium-bearing compound, pitchblende. We will discuss radioactivity in greater detail
later. Around 1898 the French chemists Marie and Pierre Curie discovered two additional radioactive elements,
radium and polonium.
In 1909 Ernest Rutherford, a chemist from New Zealand who had studied under Thomson, discovered that
every atom has a nucleus at its center. His work demonstrated that the nucleus of the atom contains 99.95% of
the mass of the atom but only 10-15 (one-thousand trillionth!) of its volume. In other words, an atom is mostly
just empty space. Based on his work we can calculate that if the nucleus of an atom is the size of a football
stadium, the atom itself would be about the size of the earth! In 1919 Rutherford discovered the proton, a
particle found in the nuclei of all atoms. This was the first identified building block of all atomic nuclei. The
name derives from the Greek word "protos," meaning "first." In 1932 Rutherford's student, James Chadwick,
discovered the neutron, the other particle found in nearly all atomic nuclei.
I do not require you to remember dates in this class. However, after reading this section you will hopefully gain
insight into the way that science creeps forward, step by step, sometimes rapidly, sometimes slowly. Science
rarely advances by enormous leaps and bounds.
The structure of the atom
Atoms are made of three types of subatomic particles, which are summarized in the table below:
charge
relative mass
absolute mass (kg)
proton
+1
1
1.6726 x 10-27
neutron
0
1
1.6749 x 10-27
electron
-1
1/1800
9.109 x 10-31
Protons and neutrons are held together in the nucleus at the center of the atom. Under normal circumstances
particles with similar electrical charges repel each other, but protons are held together in the nucleus by the
Strong Force, which is about 100 times stronger than the electrostatic repulsion experienced by similarly
charged particles. The Strong Force, while not a particularly clever name, is the strongest physical force in the
known universe. Neutrons also help to mitigate the electrical repulsion that arises when positively charged
protons are jammed in an atomic nucleus. It is the number of protons in an atomic nucleus that determines the
identity of the atom.
Electrons orbit the nucleus relatively far away. As
we stated above, the atom is mostly empty space.
An electrostatic attraction, or, electromagnetic attraction,
occurs between the positively charged nucleus and the
negatively charged electrons as they move around it. The
number of electrons, or, more correctly, they way that they are
arranged around the nucleus, determines an atom's chemical
reactivity.
Atoms and protons
The number of protons in the nucleus is unique for the atoms of each element. The number of protons in an
atom is invariable in chemical reactions. If the number of protons did change, the atom would be converted into
an atom of a different element. As we just said, and let us reiterate for emphasis: in chemical reactions, the
number of protons in reacting atoms never changes. In nuclear reactions, which are not chemical reactions, the
number of protons in the nuclei of reacting atoms may change. When this occurs, the reacting atoms are
changed to atoms of a different element through a process called transmutation.
If we examine a periodic table we find three pieces
of information for every element.
The top number is the atomic number. In the above
figure we see that this number, six, is the number of
protons in the nucleus of all carbon atoms. The atomic
number is sometimes referred to as the"Z" number of
an element.
Atoms and neutrons
While the number of protons for a given element is invariable, the number of neutrons in a nucleus can
differ. In other words, it is possible for atoms to have the same numbers of protons but different numbers of
neutrons. We call these isotopes, which are atoms with the same number of protons but different numbers of
neutrons. All elements have at least one isotope. Most elements have two or more isotopes.
The sum of the protons and neutrons in any isotope is known as its atomic mass. We can indicate the atomic
number and the atomic mass (or, as see in the figure below, the mass number) of any isotope using it's
elemental
symbol:
The atomic number of an element is indicated as a subscript to
the left of the elemental symbol. The atomic mass is indicated
as superscript to the left of the elemental symbol.
If we know the atomic mass of an isotope we can calculate the number of neutrons found in its nucleus. The
number of neutrons in any isotope is equal to the difference between the isotope's mass number and its atomic
number (number of protons). As examples, carbon has three isotopes and chlorine has two isotopes:
12
13
6C
14
6C
6C
atomic mass
12
13
14
number of protons
6
6
6
number of neutrons
6
7
8
35
17Cl
37
17Cl
atomic mass
35
37
number of protons
17
17
number of neutrons
18
20
The mass of individual atoms is expressed in amu, atomic mass units. One amu is equal to 1/12 the mass of
single a 12C isotope, or to 1.661 x 10-27 kg. In simpler terms, we can state that protons and neutrons each have a
mass of 1 amu when calculating atomic mass. This means that the three isotopes of carbon have atomic masses
of 12 amu, 13 amu, and 14 amu respectively, while we say that the atomic masses of the two isotopes of
chlorine weigh 35 amu and 37 amu respectively.
The bottom number in the Periodic Table for each element is the atomic weight, i.e. the weighted average of
the atomic masses of all of the isotopes. In the figure above, the atomic weight of carbon is given as 12.011
amu. A weighted average is calculated by taking into consideration both the mass and the abundance of each of
the isotopes of an element. We calculate the atomic weight of an element by multiplying the atomic mass of
each isotope by its fractional abundance and then adding all of the resulting numbers for all of an element's
isotopes together.
The abundance of 12C is 98.9%, of 13C is 1.1%, and of 14C is around ~0%. In other words, if we take a
million carbon atoms and examine them, we will find that roughly 989,000 of them are 12C isotopes, 11,000 are
13
C isotopes, and only about 1 of every million atoms is a 14C isotope (even though 14C is not very abundant it is
still useful, since it is the isotope examined when objects are carbon dated in archeological studies).
(0.989)(12) + (0.011)(13) + (0)(14) = 12.011 amu
which is the same as the atomic weight reported for carbon on your periodic table. Other examples of atomic
weight calculations:


The relative abundance of 35Cl is 75.78% and of 37Cl is 24.22%. The atomic weight of chlorine is equal
to:
(0.7578)(35) + (0.242)(37) = 35.48 amu
The relative abundance of 50Cr is 4.345%, of 52Cr is 83.789%, of 53Cr is 9.501%, and of 54Cr is 2.315%.
The atomic weight of chromium is equal to:
(0.0.4345)(50) + (0.83789)(52) + (0.09501)(53) + (0.02315)(54) = 52.03 amu
Atoms that have two or more isotopes will always have an atomic weight that is not an integer (i.e.,
an integer is not a fraction, i.e. there is no decimal place in the number). Check the following periodic
table and determine how many elements have an atomic weight that is a whole number:
What does this suggest about the number of isotopes for each of these elements?
Atoms and electrons
In their elemental (natural) state elements are electrically neutral. In other words, in each atom the number of
electrons is equal to the number of protons. If an atom gains or loses electrons (remember: atoms never gain or
lose protons in chemistry!), an electrical imbalance will be created and the atom will become electrically
charged. Atoms with an electrical charge are called ions. We speak of two categories of ions in chemistry.
Cations have lost electrons, and since (#electrons < #protons), cations have a net positive charge. Anions have
gained electrons, and since (#electrons > #protons), anions have a net negative charge.
Charge is represented as a superscript to the right of an element's symbol for ions. The sign tells us whether
electrons have been lost or gained. Something that has lost electrons has more protons than electrons and
becomes positively (+) charged. Something that has gained electrons has more electrons than protons and
becomes negatively (-) charged. The number tells us how many electrons have been lost or gained. This number
is always a whole number. We do not lose or gain parts or electrons. As examples:



Mg has 12 protons (how do we know? hint: what is its atomic number?) and therefore 12 electrons.
Mg2+ has 12 protons but only 10 electrons.
As has 33 protons and therefore 33 electrons. As3- has 33 protons but 36 electrons.
NH4+ and SO42- are examples of molecular ions. The "+" on the ammonium ion (NH4+) tells us that the
entire molecule has lost one electron, i.e., if we add up all of the protons and all of the electrons in the
molecule we will find that there is one fewer electron than protons. The "2-" charge on the sulfate ion
(SO42-)
tells
us
that
the
entire
molecule
has
gained
two
electrons.
(note: by convention we write the number and then the sign when representing charge. Mg2+ is correct,
Mg+2 is incorrect. If an ion has lost or gained an electron, we use only a "+" or "-" sign and do not
write "1+" or "1-". Na+ is correct, Na1+ is incorrect. If an atom is electrically neutral there is no
superscript to the right of the symbol. This absence tells us that the number of protons and electrons are
equal.)
The octet rule
Why would an atom want to gain or lose electrons? When two or more substances have the same number of
electrons, they are said to be isoelectronic. We will see that substances with the same numbers of electrons also
have the same arrangement of those electrons around their nucleus.
It turns out that atoms are ultimately more stable when they are isoelectronic with the nearest Noble Gas.
This is the basis for what is known as the octet rule. Atoms become isoelectronic with the nearest Noble Gas by
either gaining or losing electrons, or by forming chemical bonds (which we will discuss later). The number of
electrons an atom gains or loses depends on how far away it is (on the Periodic Table) from the nearest Noble
Gas. This behavior of becoming isoelectronic with the nearest Noble Gas is the driving force in many chemical
reactions.
For many elements we can predict how many electrons an atom may gain or lose based on the element's
position in the Periodic Table. Bear in mind that it takes energy for each electron lost or gained, so the atoms of
an element will be inclined to gain or lose the smallest possible number of electrons necessary to become
isoelectronic with the nearest Noble Gas.
All alkali metals (Group 1A elements) find it easiest to become isoelectronic with the nearest Noble Gas by
losing one electron. When lithium loses an electron, it becomes isoelectronic with helium. When sodium loses
an electron, it becomes isoelectronic with neon. When potassium loses an electron it becomes isoelectronic with
argon. And so on.
All alkali earth metals (Group 2A elements) find it easiest to become isoelectronic with the nearest Noble
Gas by losing two electrons. When beryllium loses two electrons it becomes isoelectronic with helium. When
magnesium loses two electrons, it becomes isoelectronic with neon. When calcium loses two electrons it
becomes isoelectronic with argon. And so on.
The Group 3A elements find it easiest to become isoelectronic with the nearest Noble Gas by losing three
electrons. In other words, we often find B3+, Al3+, Ga3+, and so on.
The Group 5A elements find it easiest to become isoelectronic with the nearest Noble Gas by gaining three
electrons. In other words, we often find N3-, P3-, As3-, and so on.
The chalcogens (Group 6A elements) find it easiest to become isoelectronic with the nearest Noble Gas by
gaining two electrons. In other words, we often find O2-, S2-, Se2-, and so on.
The halogens (Group 7A elements) find it easiest to become isoelectronic with the nearest Noble Gas by
gaining one electron. In other words, we often find F-, Cl-, Br-, and so on.
Trying to predict the charge of the transition metals (see the Periodic Table above) or the Group 4A elements
from the Periodic Table is not easily done and will not be discussed here.You should also know that, as a rule,
metals typically lose electrons in chemical reactions and non-metals typically gain electrons in chemical
reactions.
You will find it incredibly useful to memorize the relationship between charge and position in the Periodic
Table as quickly as possible. You do not need to worry about unusual atomic or chemical behavior in this class.
If can predict in general how things will behave, you have achieved my expectations of you.
Fundamentals of electrons in atoms
In neutral atoms (i.e. atoms in their elemental state) the number of protons is equal to the number of
electrons. Electrons do not orbit the nucleus in planetary fashion as we once believed. The position of an
electron is described in terms of statistical probability and may be calculated using quantum mechanics (which I
will not ask you to do in this class). Electrons cannot exist just anywhere with respect to the nucleus. They can
only be found at certain discrete (specific) distances from the nucleus, which again may be calculated using
quantum mechanics. These distances correspond to energies. The further an electron is from the nucleus the
greater its energy. Areas of high probability (90%) of finding an electron with a specific energy are called
orbitals.
The discrete distances at which electrons can be found from the nucleus can be broken down into shells,
subshells, and orbitals.
The shell in which an electron is found describes its distance from the nucleus. Shell numbers, known as "n"
quantum numbers or principal quantum numbers, range from 1 to 7 in chemistry. In other words, there are only
seven average distances from the nucleus in which electrons may be found in atoms. The shell closest to the
nucleus is n=1, the next nearest is n=2, and so on to the shell furthest away from the nucleus at n=7. As n
increases, the distance from the nucleus increases, as does the energy of the electrons found in that shell. The
spacing between shells is not linear. This has an important consequence, as we will shortly see.
Each shell contains subshells, represented by "l" (this character is a lower-case L) the angular momentum
quantum numbers. Each shell contains as many subshells as its principal quantum number, i.e., the shell with
n=1 has 1 subshell, the shell with n=2 has 2 subshells, the shell with n=3 has 3 subshells, and so on. There are
seven different types of subshells, s, p, d, f, g, h, and i subshells. Each subshell differs from the others as to the
distance from the nucleus, the number of orbitals it contains, the shape of its orbitals, and the orientation of its
orbitals with respect to the nucleus.
subshell
# orbitals
s
1
p
3
d
5
f
7
g
9
h
11
i
13
The shapes and orientations of of s and p orbitals appear as follows:
The lowest energy subshell is the "s" subshell, and subshell energy follows the order s < p < d < f < g < h < i.
There is an implication to this. If the subshells differ in energy, this means they must also differ somewhat in
their distance from the nucleus. The differences are not great, but they are significant.
The orbitals in the various subshells have magnetic quantum numbers (ml) that describe their orientation with
respect to a Cartesian coordinate system with its origin at the nucleus of the atom. Any orbital, regardless of is
shell and subshell, can only hold a maximum of two electrons. Electrons within the same orbital have opposite
spin quantum numbers (ms). Electrons can have one of two spin quantum numbers, spin up or spin down.
We can summarize what we know about shells, subshells, and the total number of electrons in a shell in a
table:
shell (n) subshell (l)
# orbitals
#electrons/orbital
total
electrons
1
s
1
2
2
2
s, p
1+3=4
2
8
3
s, p, d
1+3+5=9
2
18
4
s, p, d, f
1 + 3 + 5 + 7 = 16
2
32
5
s, p, d, f, g
1 + 3 + 5 + 7 + 9 = 25
2
50
2
72
2
98
6
7
s, p, d, f, g, h 1 + 3 + 5 + 7 + 9 + 11 = 36
s, p, d, f, g,
h, i
1 + 3 + 5 + 7 + 9 + 11 + 13
= 49
There are a few interesting features in the table we should point out. Notice that the total number of orbitals in
any shell is equal to n2. The first shell holds one orbital, the second shell holds four orbitals, the third shell holds
nine orbitals, and so on. Also note that the total number of electrons in any given shell is equal to 2n2.
I expect you to understand what shells, subshells, and orbitals are, and how many orbitals are in each shell and
subshell.
Ground states and excited states
Electrons in their lowest energy state are said to be in the ground state. When an electron absorbs energy
(either thermal or electromagnetic) it may become excited and can move to higher unoccupied orbitals. When
an electron becomes excited and moves to an orbital higher in energy than its ground state orbital is is said to be
in an exited state. We mentioned g, h, and i subshells above, and we should point out that they are virtual
subshells, only populated by excited electrons. In the ground state electrons are never found in these subshells.
The difference in energy between the various energy levels of shells and subshells correspond to specific
amounts of energy. These, in turn, correspond to specific wavelengths of light since there are mathematical
relationships between the energy, frequency, and the wavelength of radiation.
If an atom is heated or bombarded with electromagnetic radiation (with a light or a laser for example),
electrons may become excited and jump from low-energy ground state orbitals to empty higher energy orbitals.
The electrons cannot remain excited for very long. They usually only remain in their higher energy orbitals for a
fraction of a fraction of a second. But, to return to lower energy orbitals they must give off energy. The amount
of energy given off must correspond to the exact difference in energy between the high energy and low energy
orbitals. This excess energy is usually given off as electromagnetic radiation, and often it is emitted as light in
the visible portion of the spectrum. In other words, all of the light by which we see, regardless of source (from
the sun, a light bulb, etc.), is generated by energy given off by excited electrons returning to their ground state.
Electron configurations: full configurations
Of what practical value is all of this to us as we study chemistry? We are interested in being able to predict
how things will behave in chemical reactions. Believe it or not, this is a skill you will develop in this class. One
of the most important predictive tools we have is the knowledge that the electron configuration of atoms has a
direct bearing on their behavior in chemical reactions. In other words, if we know the electron configuration of
an element, we can often predict how it will react.
The electron configuration of an element is a description of the shells, subshells, and orbitals in which its
electrons are found in the ground state. While this might seem like an imposing task, it is made easier if we
remember what we said above about the relationship between shells, subshells, orbitals, and numbers of
electrons. A few other useful tidbits:



Electrons always occupy the lowest energy shells and subshells first. This means that the lowest energy
electron is all elements is found in the 1st shell, in its s subshell.
No orbital can ever hold more than two electrons, and electrons paired up in the same orbital must have
opposite spins (spin up and spin down)
When assigning electrons to subshells, all orbitals must hold at least one electron before any orbital may
hold two electrons. In other words, electrons will never pair up in an orbital together if there is an empty
orbital available (can you guess why?).
We will discuss three types of electron configurations in this class, full configurations, Noble Gas
configurations, and arrow diagrams (also known as orbital diagrams).
Full configurations are assigned to elements by determining how many electrons they have and then by
placing them in orbitals in shells and subshells, starting with those orbitals of the lowest energy and working
upward.

Hydrogen
has
one
electron.
It's
full
configuration
is
1s1.
The first number tell us the shell in which the subshell is found (1), the letter tells us which subshell (s) is









being used, and the superscripted "1" tells us that the 1s shell and subshell is occupied by a single
electron.
Helium has two electrons. It's full configuration is 1s2. In other words, there are two electrons in the
single orbital of the s subshell of the 1st shell.
Lithium has three electrons. The first two electrons are assigned to the lowest available shell and
subshell (i.e., the orbital in the lowest energy shell and subshell) and have the same configuration as
helium, 1s2. At this point we have filled both the subshell and also the shell, so we have to move to the
second shell and it's lowest energy subshell, 2s to place our third electron. The full configuration of
lithium is 1s22s1.
Beryllium has four electrons. The first two electrons are assigned to the lowest available shell and
subshell and have the same configuration as helium, 1s2. The remaining two electrons are placed in the
2s shell and subshell. The full configuration of beryllium is 1s22s2.
Boron has five electrons. The first two electrons are assigned to the lowest available shell and subshell
and have the same configuration as helium, 1s2. The next two are assigned to the 2s shell and subshell.
Since any s subshell only contains a single orbital and can therefore only hold two electrons, where do
we place boron's last electron? Remember that any shell holds has many subshells as its principal
quantum number. In the second shell, n=2. That means that the second shell contains two subshells, an s
subshell and a p subshell. The p subshell contains 3 orbitals. The remaining electron is placed in one of
the three 2p orbitals. The full configuration of boron is 1s22s22p1.
Carbon has six electrons. The first two electrons are assigned to the lowest available shell and subshell
and have the same configuration as helium, 1s2. The next two are assigned to the 2s shell and subshell.
The last two are assigned to separate 2p orbitals. The full configuration of carbon is 1s22s22p2.
Nitrogen has seven electrons. The first two electrons are assigned to the lowest available shell and
subshell and have the same configuration as helium, 1s2. Two electrons are assigned to the 2s shell and
subshell, and the remaining three electrons are assigned their own orbital in the 2p subshell. The full
configuration of nitrogen is 1s22s22p3.
Oxygen has eight electrons. The first two electrons are assigned to the lowest available shell and
subshell and have the same configuration as helium, 1s2. Two electrons are assigned to the 2s shell and
subshell, and the remaining four electrons are assigned to orbitals in the 2p subshell. At this point, since
each of the 2p orbitals contains at least one electron, one orbital must accept a second electron. The full
configuration of oxygen is 1s22s22p4.
Fluorine (note: fluorine is an element. Flourine - I don't know what flourine is, perhaps an element made
up of ground-up wheat? Be careful with spellings) has nine electrons. The first two electrons are
assigned to the lowest available shell and subshell and have the same configuration as helium, 1s2. Two
electrons are assigned to the 2s shell and subshell, and the remaining five electrons are assigned to
orbitals in the 2p subshell. At this point, since each of the 2p orbitals contains at least one electron, two
orbitals must accept a second electron. The full configuration of fluorine is 1s22s22p5.
Neon has ten electrons. The first two electrons are assigned to the lowest available shell and subshell
and have the same configuration as helium, 1s2. Two electrons are assigned to the 2s shell and subshell,
and the remaining six electrons are assigned to orbitals in the 2p subshell. At this point, each of the 2p
orbitals contains two electrons. The full configuration of neon is 1s22s22p6.
At this point we've filled all of the orbitals in all of the subshells in the first two shells. Where will we put
additional electrons? If you guessed the third shell and it's subshells, you're doing great!
Before we examine electron configurations in the 3rd shell, we should point out that if we add up the total of
all of the exponents in a full configuration, it tells us the number of electrons we've accounted for in our
configuration. Take neon for example. The sum of the the exponents in its configuration is 2 + 2 + 6 = 10. How
many electrons does neon have? This is a good way to check the correctness of any configurations you assign.

Sodium is the first element in the third period (the 3rd row down). It has 11 electrons. The first 10
electrons have the same configuration as neon, 1s22s22p6. The eleventh electron must be placed in the
third shell and it the lowest energy subshell, the 3s subshell. The full configuration of sodium is
1s22s22p63s1.


Magnesium has 12 electrons. The first 10 electrons have the same configuration as neon, 1s22s22p6. The
eleventh and twelfth electrons must be placed in the orbital in the 3s subshell. The full configuration of
magnesium is 1s22s22p63s2.
Aluminum has 13 electrons. The first 10 electrons have the same configuration as neon, 1s22s22p6. The
eleventh and twelfth electrons must be placed in the orbital in the 3s subshell. The 13th electron must be
placed in the next available orbital, which falls in the 3p subshell. The full configuration of aluminum is
1s22s22p63s23p1.
The next five elements, silicon, phosphorus, sulfur, chlorine, and argon also have electrons that occupy the
orbitals in the 3p subshell.





The full configuration of silicon is 1s22s22p63s23p2
The full configuration of phosphorus is 1s22s22p63s23p3
The full configuration of sulfur is 1s22s22p63s23p4
The full configuration of chlorine is 1s22s22p63s23p5
The full configuration of argon is 1s22s22p63s23p6
Does each of these configurations account for the correct number of electrons? How do you know?
The next element is potassium, which has 19 electrons (how do we know it has 19 electrons?). We find
potassium in the 4th period (the 4th row down). The first 18 electrons have the same configuration as argon,
1s22s22p63s23p6. How do we assign the 19th electron? On the one hand we need to remember that the 3rd shell
has 3 subshells, 3s, 3p, and 3d. We have yet to place any electrons in the orbitals in the 3d subshell. On the
other hand, potassium is found in the 4th period. Is the correct full configuration of potassium
1s22s22p63s23p63d1 or 1s22s22p63s23p64s1? We will see in a moment.
Electron configurations: Noble Gas configurations
Ok, let's admit it. Writing full electron configurations, especially of elements with many electrons, is a real
drag. Let's learn a short-cut. We can substitute the symbol of the Noble Gas in the previous period for the
configuration of lower energy electrons in a full electron configuration. We must place the symbol for the Noble
Gas in square brackets to correctly indicate what we're doing. Let's see how this works.
element
full configuration
Noble Gas in
previous period
Noble Gas
configuration
H
1s1
-
1s1
He
1s2
-
1s2
Li
1s22s1
He
[He]2s1
Be
1s22s2
He
[He]2s2
B
1s22s22p1
He
[He]2s22p1
C
1s22s22p2
He
[He]2s22p2
N
1s22s22p3
He
[He]2s22p3
O
1s22s22p4
He
[He]2s22p4
F
1s22s22p5
He
[He]2s22p5
Ne
1s22s22p6
He
[He]2s22p6 or [Ne]
Na
1s22s22p63s1
Ne
[Ne]3s1
Mg
1s22s22p63s2
Ne
[Ne]3s2
Al
1s22s22p63s23p1
Ne
[Ne]3s23p1
Si
1s22s22p63s23p2
Ne
[Ne]3s23p2
P
1s22s22p63s23p3
Ne
[Ne]3s23p3
S
1s22s22p63s23p4
Ne
[Ne]3s23p4
Cl
1s22s22p63s23p5
Ne
[Ne]3s23p5
Ar
1s22s22p63s23p6
Ne
[Ne]3s23p6 or [Ar]
Electron configurations: arrow diagrams (orbital, spin, or
electron spin diagrams)
Full electron configurations and Noble Gas configurations account for all of the electrons in an atom, but
there are times when it is helpful to display electron configurations in a slightly different way. Arrow diagrams,
also known orbital diagrams, spin diagrams, and electron spin diagrams, use dashes to represent each of the
orbitals in a subshell and arrows to represent electrons. Arrow diagrams may be drawn vertically or
horizontally. The energy of electrons in the diagram increases as we move from lower to higher levels or as we
move from left to right, depending on if the arrow diagram is drawn vertically or horizontally. Remember that
electrons repel each other (why?) and to minimize the repulsion, electrons that share an orbital have different
spins, spin up or spin down.
The arrow diagram for hydrogen is
Note the use of a single dash for the 1s orbital, and also that the shell and subshell are
placed as a label beneath the dash.
The arrow diagram for helium is
Note that the 1s orbital is now filled, and the two electrons, represented by the arrows,
point in opposite directions since they have opposite spins.
The arrow diagram for lithium is
We now have two dashes, one for the orbital in the 1s shell and subshell, and
one for the 2s orbital. Again, we always label the orbitals (dashes) as to which shell and subshell they belong.
The arrow diagram for beryllium is
At this point we have filled the 2s orbital. The next six elements will have
electrons in the three 2p orbitals. The first three elements with 2p electrons are boron, carbon, and nitrogen:
Did you notice that each 2p electron occupied an empty orbital? These halffilled orbitals will be forced to hold an extra electron as we progress through the arrow diagrams for the next
three elements, oxygen, fluorine, and neon.
And so it goes. We draw a single dash to represent the 3s orbital, and then,
when necessary, three dashes to represent the three 3p orbitals. As an example, the arrow diagram for
phosphorus is
Electron configurations and the Periodic Table
What we need is a tool to help us remember the order in which subshells are filled. You may be amazed to
learn that the periodic table is just such a device. The periodic table not only provides information about the
number of protons and neutrons in each of the elements, it also gives information about the highest energy
electron for each element.
The periodic table can be divided into four regions. Groups 1A and 2A are known are the "s" block, because
the highest energy electron of each element in these two groups is an "s" electron. Groups 3A, 4A, 5A, 6A, 7A,
and 8A are known as the "p" block, because the highest energy electron of each element in these six groups is
a"p" electron. The transition metals (Groups 1B - 8B) are found in the "d" block, because the highest energy
electron of each element in these ten
groups is a"d" electron. The
lanthanides and actinides are found in
the "f" block, because the highest
energy electron of each element in
these 14 groups is a"f" electron.
Did you notice that the "s" block
contains two groups? Did you
remember that an s subshell can hold a
maximum of two electrons? The "p"
block contains six groups, and a p
subshell can also hold six electrons.
The "d" block holds 10 groups, and a d
subshell can hold up to 10 electrons.
The "f" block holds 14 groups, and a f
subshell can hold 14 electrons?
Coincidence? I think not. Remarkably
clever? Absolutely!
How do we use the periodic table to assign electron configurations to elements? We must find the position of
the element of interest, notice which period it is in and which block it is in, and observe how many groups to the
right of the left-most group in the block the element is. Then we work backwards to account for the rest of the
electrons. Let's choose a few examples and see how this works:

Carbon is in the second period, and in the second group from the left in the "p" block (use a proper
periodic table, not the diagram immediately above, to work these problems). Therefore, carbon's highest
energy electron must have the configuration 2p2. Which blocks are found in the periodic table before the
2p-block? The 2s-block and the 1s-block. Carbon has the electron configuration 1s22s22p2.



Sulfur is found in the third period, and in the fourth group from the left in the "p" block. Therefore,
sulfur's highest energy electron must have the configuration 3p4. The groups found in the periodic table
before 3p are 3s, 2p, 2s, and 1s. Carbon has the electron configuration 1s22s22p63s23p4.
Let's return to the electron configuration of potassium. We find it located in the 4th period, and in the
left-most column of the "s" block. This makes the configuration of potassium's highest energy electron
4s1. This is potassium's 19th electron. What is the configuration of the other 18? 1s22s22p63s23p6 The
full configuration of potassium is 1s22s22p63s23p64s1.
Calcium has 20 electrons. It is located in the 4th period, second to the left in the "s" block. Calcium's
highest energy electron is 4s2. note that this actually accounts for both 4s electrons The rest of calcium's
configuration is as potassium's, so calcium's full configuration is 1s22s22p63s23p64s2.
Next we come to scandium, which has 21 electrons. Scandium is located in the forth period, in the left-most
group of the "d" block. Does this mean that scandium's highest energy electron has a 4d 1 configuration? Not
quite. Remember that the 3rd shell has 3 subshells, s, p, and d. Recall as well that the spacing between shells is
not linear. An important consequence of this non-linearity is that as shells move further from the nucleus, their
subshells begin to overlap with those in higher shells. The first d electrons we encounter in the periodic table are
3d electrons, meaning that scandium's highest energy electron has a 3d1 configuration. Scandium's full
configuration is 1s22s22p63s23p64s23d1.
Because of overlap, the d electrons in every period are from one shell lower than the period in which they
reside, while the s and p electrons are from the same shell as the period. In other words, 4s3d4p, 5s4d5p,
6s5d6p, and 7s6d7p.
Where do we find elements with f electrons in the periodic table? Note the order of the elements in the 6th
period. The atomic number of cesium (Cs) is 55, the atomic number of barium (Ba) is 56, the atomic number of
lanthanum (La) is 57, and the atomic number of hafnium (Hf) is 72. Where are the elements with atomic
numbers from 58 to 71? In the "f" block, in the lanthanides and actinides, which are found in a strip below the
periodic table. Why are these elements placed separately from the main body of the periodic table rather than
integrated in their proper place? There is a reason, although it is not necessarily a very good one. By including
the 14 f-block groups in their proper place a periodic table we nearly double it's width. This means that it is
difficult to fit the entire table on an regular-sized piece of paper in print that is large enough to read clearly. A
periodic table with the f-block in its proper place looks like this:
Remember that scandium (Sc), yttrium (Y), lanthanum (La), and actinium (Ac) are all part of the d-block in the
above table, even though they are separated from the rest of the d-block by the f-block.
To which f-block do elements 58 through 71 belong? The first shell with an f subshell is the 4th shell.
Elements 58 through 71 have electrons with a 4f configuration. The f subshells overlap even more severely with
higher level subshells than the d subshells.
In the 7th period there is a gap between actinium (Z=89) and Rutherfordium (Z=104), with the missing
elements, those with atomic numbers from 90 to 103, again found in the f-block. These are 5f elements.
To summarize, s and p electrons always have the same number as the period in which they are found. d
electrons always have a number one less than the period in which they reside, and f electrons, 2 less.
Let's do a few examples:



Arsenic (As) has 33 electrons. It is found in the 4th period, 3 groups from the left in the p-block.
Arsenic's highest energy electron has the configuration 4p3. By using the periodic table to remind us
which subshells lie beneath 4p in energy, and their order, we assign arsenic a full electron configuration
of 1s22s22p63s23p64s23d104p3. note that 2 + 2 + 6 + 2 + 6 + 2 + 10 + 3 = 33
Palladium (Pd) has 46 electrons. It lies in the 5th period, 8 groups from the left. Palladium has the full
electron configuration 1s22s22p63s23p64s23d104p65s24d8. note that 2 + 2 + 6 + 2 + 6 + 2 + 10 + 6 + 2 +
8 = 46
Polonium (Po) has 84 electrons. It lies in the 6th period, four groups from the left in the p-block.
Polonium has the full electron configuration 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p4. note
that 2 + 2 + 6 + 2 + 6 + 2 + 10 + 6 + 2 + 10 + 6 + 2 + 14 + 10 + 4 = 84
A few final notes. First, you do not need to memorize the periodic table. I will always provide you with one
on exams and permit you to use one on quizzes. Second, there are periodic tables available at bookstores that
offer electron configurations that are correct but not in the order provided by the periodic table. As an example,
you might find the electron configuration of arsenic given as 1s22s22p63s23p63d104s24p3, rather than as we have
written it aboveFinally, the electron configurations of d and f-block elements can be tricky things. Sometimes
their full electron configurations differ slightly from those we might predict using the periodic table. There are
good reasons for these deviations, but we will not discuss them in this class, and I do not require you to know
them. If I ask for the electron configuration of a d-block element I will accept as correct the configuration
derived from the periodic table, regardless of whether it is absolutely correct or not. You will not be asked about
the electron configuration of f-block elements.
Inner shell, outer shell, and valence electrons
Noble gas configurations provide us with a short-hand way of writing electron configurations. They also
help us categorize an atom's electrons into one of two types, inner and outer shell electrons.
All electrons that can be incorporated into those denoted by the Noble Gas (i.e., all electrons found in lower
periods than the period on which the element is found) are called inner shell electrons. Inner shell electrons do
not participate in chemical reactions.
All electrons found in the same period as the element are known as outer shell electrons. Outer shell
electrons are often called valence electrons, but there is actually a slight distinction between the two. Valence
electrons are those outer shell electrons available to participate in chemical reactions.
Completely filled p, d, and f subshells are extremely stable, to the extent that these electrons behave as
though they are inner shell. The electrons of filled p, d, and f subshells do not participate in chemical reactions.
This means that all of an atom's outer shell electrons may or may not be valence electrons.
If we examine, for example, aluminum, we find its noble gas configuration to be [Ne]3s23p1. This means
aluminum has 13 total electrons, of which 10 are inner shell and three are outer shell electrons, all three of
which are valence electrons.
Arsenic has the full configuration 1s22s22p63s23p64s23d104p3and the noble gas configuration [Ar]4s23d104p3.
Arsenic has 33 total electrons, of which there are 18 inner shell electrons and 15 outer shell electrons. Of the 15
outer shell electrons there are only 5 valence electrons, the 4s and 4p electrons. The other 10 are tied up in the
filled 3d subshell and are unreactive.
element
Noble Gas configuration
inner shell
electrons
outer shell
electrons
valence
electrons
commonly formed
ions
H
1s1
0
1
1
H+
He
1s2
0
2
0
none
Li
[He]2s1
2
1
1
Li+
Be
[He]2s2
2
2
2
Be2+
B
[He]2s22p1
2
3
3
B3+
C
[He]2s22p2
2
4
4
C4-
N
[He]2s22p3
2
5
5
N3-
O
[He]2s22p4
2
6
6
O2-
F
[He]2s22p5
2
7
7
F-
Ne
[Ne]
2
8
0
none
Na
[Ne]3s1
10
1
1
Na+
Mg
[Ne]3s2
10
2
2
Mg2+
Al
[Ne]3s23p1
10
3
3
Al3+
Si
[Ne]3s23p2
10
4
4
Si4-
P
[Ne]3s23p3
10
5
5
P3-
S
[Ne]3s23p4
10
6
6
S2-
Cl
[Ne]3s23p5
10
7
7
Cl-
Ar
[Ar]
10
8
0
none
K
[Ar]4s1
18
1
1
K+
Ca
[Ar]4s2
18
2
2
Ca2+
Sc
[Ar]4s23d1
18
3
3
Sc3+
Ti
[Ar]4s23d2
18
4
4
Ti4+, Ti3+
V
[Ar]4s23d3
18
5
5
V5+, V4+, V3+, V2+
Cr
[Ar]4s23d4
18
6
6
Cr6+, Cr3+, Cr2+
Mn
[Ar]4s23d5
18
7
7
Mn7+, Mn6+, Mn4+,
Mn3+, Mn2+,
Fe
[Ar]4s23d6
18
8
8
Fe3+, Fe2+
Co
[Ar]4s23d7
18
9
9
Co3+, Co2+
Ni
[Ar]4s23d8
18
10
10
Ni3+, Ni2+
Cu
[Ar]4s23d9
18
11
11
Cu2+, Cu+
Zn
[Ar]4s23d10
18
12
2
Zn2+
Ga
[Ar]4s23d104p1
18
13
3
Ga3+
Ge
[Ar]4s23d104p2
18
14
4
Ge4+
As
[Ar]4s23d104p3
18
15
5
As3-
Se
[Ar]44s23d104p4
18
16
6
Se2-
Br
[Ar]4s23d104p5
18
17
7
Br-
Kr
[Ar]4s23d104p6
18
18
0
none
Generally there is a very good correlation between the number of valence electrons an element has and the
ion it forms to satisfy the octet rule. This is especially true of the s and p block elements found in the first three
periods. The behavior of the transition metals is not well described by the octet rule, and we will not discuss
why they behave the way they behave in this class.
You should be able to accurately predict for each s and p block element whether it forms a cation or an anion
and its charge. This is based on position in the periodic table, as we have seen.
The electron configurations of ions
When two substances have the same electron configuration (because they have the same number of
electrons), they are said to be isoelectronic. The Noble Gas configuration of neon is [Ne] and of sodium is
[Ne]3s1. When an atom loses electrons, it loses outer shell electrons. This means that when sodium becomes
sodium ion, Na+, it loses the 3s electron. This makes sodium ion isoelectronic with neon. how do we know that
sodium always tends to lose one electron? where is Na found in the periodic table? and remember: metals tend
to lose electrons in chemical reactions, nonmetals tend to gain electrons in chemical reactions.
Magnesium has a Noble Gas configuration of [Ne]3s2. Magnesium nearly always exists as Mg2+ in chemical
reactions. how do we know this? where is Mg found in the periodic table? Mg2+ has lost its two 3s electrons,
and has become isoelectronic with both neon and Na+ in the process.
Fluorine has the electron configuration 1s22s22p5. It tends to gain an electron in reactions. how do we know
this? where is F found in the periodic table? When an atom gains electrons, they are always added to empty or
half-filled orbitals in the outer shell. The electron configuration of F- is 1s22s22p6, which makes fluoride ion
isoelectronic with neon.
Oxygen has the electron configuration 1s22s22p4. Oxygen tend to gain two electrons. how do we know this?
where is O found in the periodic table? The electron configuration of O2- is 1s22s22p6, which makes it
isoelectronic with neon and F-.
Note that more than two substances can be isoelectronic. In our examples in this section we find that Na+,
Mg2+, Ne, F-, and O2- are all isoelectronic. What are examples of other ions that are isoelectronic with neon?
Al3+ and N3- are two examples
Atomic properties and periodic trends
There are four basic properties of elements, the extent of which can be predicted in a relative sense from
position in the periodic table. These are called periodic trends. The properties are atomic size, metallic
character, ionization energy, and electronegativity.
Atomic size is the actual size of the atoms. Within a group, atoms become larger as we move down the
group. Since moving down a group involves moving to electron shells that are further away from the nucleus, it
should not surprise you that atomic size increases. As an example, fluorine is smaller than chlorine, which is
smaller than bromine, which is smaller than iodine. Atomic size also increases as we move from right to left
within a period. This is a bit surprising but not difficult to explain. Within a period, as we move from left to
right, the number of electrons increases. Lithium has fewer electrons (3 electrons) than beryllium (4 electrons),
which has fewer electrons than boron (5 electrons), which has fewer electrons than carbon (6 electrons), and so
on. While there is a balance between the numbers of protons and electrons in each of these elements, as the
number of protons in the nucleus increases, electrons at the same average distance from the nucleus (i.e., in the
same period) are attracted more forcefully to the nucleus. As a result, the size of the atom becomes smaller.
If we are interested in
comparing the size of the atoms of
two or more elements, without
actually knowing any hard
information about the size of the
atoms, we can still qualitatively
predict which atoms are the largest
and which are the smallest.
Compare the atoms of barium (Ba)
and chlorine. Which are larger?
barium Compare the atoms of
cesium (Cs), ruthenium (Ru), and
sulfur and arrange them in order of
increasing size. S (smallest) < Ru < Cs (largest)
Metallic character involves the extent to which a substance behaves as a metal. Remember from chapter 1
that metals are elements characterized by good thermal and electrical conductivity and are usually lustrous,
malleable, and ductile. Metals are also characterized by the tendency to give up electrons in reactions (i.e. form
cations).
It is the "looseness" with which valence electrons are held that determines how metallic an element is. The
further down a group, the further the valence electrons are found from the nucleus and the more loosely they are
held. As we stated above, as we move from left to right within a period,
We can qualitatively predict how
metallic elements are relative to
each based on this information.
Compare the atoms of cesium (Cs),
ruthenium (Ru), and sulfur and
arrange them in order of increasing
metallic character. S (least metallic,
in fact a non-metal) < Ru < Cs
(most metallic)
Ionization energy is the energy
required to remove an electron from
a neutral atom, thus making it a
cation. Whenever an atom loses an
electron, it takes energy. The energy required to remove an electron increases as they are held more tightly. This
means that the closer electrons are found to the nucleus (i.e., going up a group or from right to left within a
period), the more energy
required to remove them from
an atom.
We can qualitatively predict
how ionization energies are
relative to each based on this
information. Compare the
atoms of cesium (Cs),
ruthenium (Ru), and sulfur and
arrange them in order of
increasing ionization energy.
Cs (lowest ionization energy) <
Ru < S (highest ionization
energy)
There is a relationship between ionization energy and ion size. When an electron is removed from an atom
and a cation is formed, an imbalance between electrons and protons results. As a consequence, the remaining
electrons are pulled nearer to the nucleus then they were in the electrically neutral atom. In other words, cations
are smaller than the neutral atom from which they originated. It is possible to remove more than one electron
from an atom and to form cations with +2, +3 and higher charges. This loss can be thought of as occurring in a
stepwise process, meaning if an atom loses three electrons, it doesn't lose all three simultaneously, but rather, in
three steps that follow each other in rapid succession. Since each loss of an electron results in the electrons
being held more tightly by the nucleus, it should not surprise you to read that it takes more energy to remove a
second electron from an atom than it did the first, far more still to remove a third electron than it did to remove
the second, and so on. At the same time, with the loss of each electron, the cation becomes smaller and smaller
as it holds its remaining electrons more and more tightly. So, as an example, if we compare the size of
aluminum with that of its cations: Al > Al+ > Al2+ > Al3+
Electron affinity is the energy associated with the addition of a single electron to a neutral atom, thus
making it an anion. Whenever an atom gains an electron, it may result in the release of energy (i.e. an
exothermic process) if a stable anion is formed, or it may require an input of energy (i.e. an endothermic
process) to make it happen if an unstable anion is formed. Electron affinity does not have a consistent periodic
trend, but it is related to electronegativity which we shall soon discuss and which does have a consistent
periodicity to it.
As with cations, the size of anions differs from that of neutral atoms. When an electron is added to an atom
and an anion is formed, an imbalance between electrons and protons results. As a consequence, the electrons are
not held quite as tightly as in the neutral atom. In other words, anions are larger than the neutral atom from
which they originated. It is possible to add more than one electron to an atom and to form anions with -2, -3 and
higher charges. This addition can be thought of as occurring in a stepwise process, meaning if an atom gains
three electrons, it doesn't gain all three simultaneously, but rather, in three steps that follow each other in rapid
succession. With the addition of each electron, the anion becomes larger and larger as it holds its electrons more
and more loosely. So, as an example, if we compare the size of nitrogen with that of its anions: N < N- < N2- <
N3Electronegativity is the tendency of an atom to attract electrons to itself. As we think about the elements,
we should recall that metals tend to lose electrons in reactions and that nonmetals tend to gain electrons in
chemical reactions. Therefore, as we move across a period from left to right, from metals to nonmetals, we
would expect that electronegativity should increase, and it does. Electronegativity also increase as we move up
a group. This is related to the tendency of smaller atoms to cling to their electrons more tightly. The noble gases
do not have electronegativity, i.e, they have an electronegativity of zero. why? Electronegativity is commonly
ranked on a scale from 0 to 4. The most electronegative element is fluorine, with an electronegativity of 4.
Francium (Fr) is the least
electronegative element with a
value of 0.7.
We can qualitatively predict
how
electronegativities
are
relative to each based on this
information. Compare the atoms
of cesium (Cs), ruthenium (Ru),
and sulfur and arrange them in
order
of
increasing
electronegativity. Cs (lowest
electronegativity) < Ru < S
(highest electronegativity)