Credit to the owner
Atomic Structure
The text provides a historical perspective of how the internal structure
of the atom was discovered. It is certainly one of the most important
scientific discoveries of this century, and I recommend that you read
through it. However, we will begin our discussion of the atom from the
modern day perspective.
All atoms are made from three subatomic particles
 Protons, neutron & electrons.
These particles have the following properties:
Particle
Charge
Mass (g)
Mass (amu)
Proton
+1
1.6727 x 10-24 g
1.007316
Neutron
0
1.6750 x 10-24 g
1.008701
Electron
-1
9.110 x 10-28 g
0.000549
In the above table I have used a unit of mass called the atomic mass unit
(amu). This unit is much more convenient to use than grams for describing
masses of atoms. It is defined so that both protons and neutrons have a
mass of approximately 1 amu. Its precise definition will be given later.
The important points to keep in mind are as follows:


Protons and neutrons have almost the same mass, while the electron
is approximately 2000 times lighter.
Protons and electrons carry charges of equal magnitude, but
opposite charge. Neutrons carry no charge (they are neutral).
It was once thought that protons, neutrons and electrons were spread out
in a rather uniform fashion to form the atom (see J.J. Thompson’s plum
pudding model of the atom on page 42), but now we know the actual
structure of the atom to be quite different.
What does an atom look like?
Protons and neutrons are held together rather closely in the center of
the atom. Together they make up the nucleus, which accounts for nearly
all of the mass of the atom.
Electrons move rapidly around the nucleus and constitute almost the
entire volume of the atom. Although quantum mechanics are necessary to
explain the motion of an electron about the nucleus, we can say that the
distribution of electrons about an atom is such that the atom has a
spherical shape.
Atoms have sizes on the order of 1-5 � (1 angstrom = 1 � = 1  10-10 m)
and masses on the order of 1-300 amu.
To put the mass and dimensions of an atom into perspective consider the
following analogies. If an atom were the size of Ohio stadium, the nucleus
would only be the size of a small marble. However, the mass of that
marble would be ~ 115 million tons.
What holds an atom together?
The negatively charged electron is attracted to the positively charged
nucleus by a Coulombic attraction.
The protons and neutrons are held together in the nucleus by the strong
nuclear force.
How many electrons, protons and neutrons are contained in an atom?
Atoms in their natural state have no charge, that is they are neutral.
Therefore, in a neutral atom the number of protons and electrons are the
same. If this condition is violated the atom has a net charge and is called
an ion.
The number of protons in the nucleus determines the identity of the
atom. For example all carbon atoms contain six protons, all gold atoms
contain 79 protons, all lead atoms contain 82 protons.
Two atoms with the same number of protons, but different numbers of
neutrons are called isotopes.
How does the structure of the atom relate to its properties?
Chemical reactions involve either the transfer or the sharing of electrons
between atoms. Therefore, the chemical reactivity/ properties of an
element is primarily dependent upon the number of electrons in an atom
of that element. Protons also play a significant role because the tendency
for an atom to either lose, gain or share electrons is dependent upon the
charge of the nucleus.
Therefore, we can say that the chemical reactivity of an atom is
dependent upon the number of electrons and protons, and independent of
the number of neutrons.
The mass and radioactive properties of an atom are dependent upon the
number of protons and neutrons in the nucleus.
Note: The number of protons, neutrons and electrons in an atom
completely determine its properties and identity, regardless of how and
where the atom was made. So it is inaccurate to speak of synthetic atoms
and natural atoms. In other words a lead atom is a lead atom, end of
story. It doesn’t matter if was mined from the earth, produced in a
nuclear reactor, or came to earth on an asteroid.
Symbolism
There is a symbolism or shorthand for describing atoms which is
universally used across all scientific disciplines
Atomic Number (Z)  The # of protons
Mass Number (A)  [The # of protons] + [the # of neutrons]
The number of protons, neutrons and electrons in an atom are uniquely
specified by the following symbol
A
SyC
where:



Sy = The elemental symbol (i.e. C, N, Cr)  defines the # of protons
A = The mass number  [# of protons] + [# of neutrons]
C = The net charge  [# of protons] – [# of electrons]
Example
Lets start with a neutral boron 10 atom  10B
Since the atom is a boron atom the periodic table tells us that there
are 5 protons in the nucleus Z = 5.
The atom is neutral so that the number of electrons must balance the
number of protons, 5 electrons.
The mass number is 10, so that the number of neutrons is A - Z = 10 - 5
= 5 neutrons.
University of Colorado GEOLOGY 1010
Class Note 2
Atoms and Elements, Isotopes and Ions
Atoms are composed of protons, neutrons, and electrons. A proton has an electric
charge of +1 and a rest mass of 1.67 x 10-24 gm. A neutron has a charge of 0 and a
rest mass of 1.67 x 10-24 gm. (about the same as a proton). An electron has a charge
of -1 and a rest mass of 9.11 x 10-28 gm. (much, much less than a proton). The
important point here is that the electron mass is negligible relative to protons and
neutrons.
The heavy particles (protons and neutrons) are bound into the nucleus, whereas the
electrons form complex orbitals about the nucleus.
The chemical properties of an element depend on the number of protons (i.e. the net
electric charge) of the nucleus. The number of protons in the nucleus is known as
the atomic number of the element. Atomic numbers for natural element range from 1
(hydrogen) to 92 for uranium.
The number of protons plus neutrons in the nucleus is known as the mass number of
the atom. Atoms of a given element (atomic number) may have differing numbers of
neutrons. Atoms of the same element with different mass numbers are known
as isotopes. The mass numbers or isotopes of an element are denoted as preceding
superscripts. For example the stable isotopes of the element oxygen are
denoted 18O, 17O and 16O. Oxygen has an atomic number of 8 (eight protons). The
nucleus of 16O thus contains eight protons and eight neutrons. How many neutrons are
in the nucleus of 18O? (ans.: 10). Because elements may have several stable isotopes,
the average mass number of an element is the atomic weight and is commonly not an
integer.
Atoms may not change their atomic numbers or mass numbers except by very
energetic nuclear reactions. However atoms may gain or lose electrons in ordinary
chemical reactions. If an atom has the same number of electrons as protons, it is a
neutral atom. If it has a net charge, (more or less electrons than protons) it is an ion. If
it has more electrons than protons it has a net negative charge and is known as
ananion. If it has fewer electrons than protons it has a net positive charge and is
known as a cation. The ionic state may be denoted as a following superscript (e.g. O2, Fe2+). The common ionic states of a atom are known as its valences.
Because electrons arrange themselves in discreet orbitals about the nucleus and the
orbitals repeat in shells, the chemical properties of the elements tend to repeat as the
atomic number increases. This periodicity of properties gives rise to the periodic
table of the elements.
The elements H, He, and minor amounts of Li were formed in the original Big Bang.
All heavier elements were formed form the primordial H and He by
nuclear fusion reactions in stars. The fusion reaction proceeds in steps in stars
massive enough to undergo the full sequence. (Our sun is not massive enough to form
elements more massive than He by direct fusion and will die when all the H is
consumed.) First H is consumed to form He. When the H is consumed, the star
collapses until He is "ignited" to form Be and C. There are many free neutrons in
these reactors and nuclei will capture enough neutrons to stabilize themselves. Most
of the heavy elements are formed by neutron capture rather than by direct fusion. In
the last stage Fe is formed by direct fusion of Si and other light elements. This
reactions is rapid and results in an explosion. Our solar system condensed from the
remnants of one of these supernova explosions.
Chemical bonds may be either ionic, metallic, covalent or vander Waals (mirror
charge), and the bond type preferred by the various elements will determine their
geochemical affinity. Ionically bonded elements are termed lithophile and combine
with the most abundant element, O, and are enriched in the silicate and oxide minerals
(rocks). Metallically bonded elements are termed siderophile and combine with
native iron and are enriched in the core. Covalently bonded elements are
termed chalcophile and combine with sulfur and are enriched in ore minerals.
The atmophile elements form only very weak vander Waals bonds and did not
condense in the inner solar system. They are depleted in Earth and enriched in the
outer planets.
Here is a full Periodic Table of the Elements. Geological Periodic Table
GEOL 1010 Syllabus
GEOL 1010 Class Note 3
GEOL 1010 Class Note 1
Joe Smyth's Home Page
Atomic Structure: The Quantum Mechanical
Model
Two models of atomic structure are in use today: the Bohr model and the quantum mechanical model. The quantum
mechanical model is based on mathematics. Although it is more difficult to understand than the Bohr model, it can be
used to explain observations made on complex atoms.
A model is useful because it helps you understand what’s observed in nature. It’s not unusual to have more than one
model represent and help people understand a particular topic.
The quantum mechanical model is based on quantum theory, which says matter also has properties associated with
waves. According to quantum theory, it’s impossible to know the exact position and momentum of an electron at the
same time. This is known as the Uncertainty Principle.
The quantum mechanical model of the atom uses complex shapes of orbitals (sometimes called electron clouds),
volumes of space in which there is likely to be an electron. So, this model is based on probability rather than
certainty.
Four numbers, called quantum numbers, were introduced to describe the characteristics of electrons and their
orbitals:

Principal quantum number: n

Angular momentum quantum number: l

Magnetic quantum number:

Spin quantum number:
The principal quantum number
The principal quantum number n describes the average distance of the orbital from the nucleus — and the energy of
the electron in an atom. It can have positive integer (whole number) values: 1, 2, 3, 4, and so on. The larger the value
of n, the higher the energy and the larger the orbital. Chemists sometimes call the orbitals electron shells.
The angular momentum quantum number
The angular momentum quantum number l describes the shape of the orbital, and the shape is limited by the
principal quantum number n: The angular momentum quantum number l can have positive integer values from 0 to
n–1. For example, if the n value is 3, three values are allowed for l: 0, 1, and 2.
The value of l defines the shape of the orbital, and the value of n defines the size.
Orbitals that have the same value of n but different values of l are called subshells. These subshells are given
different letters to help chemists distinguish them from each other. The following table shows the letters
corresponding to the different values of l.
Letter Designations of the Subshells
Value of l (subshell)
Letter
0
s
1
p
2
d
3
f
4
g
When chemists describe one particular subshell in an atom, they can use both the n value and the subshell letter —
2p, 3d, and so on. Normally, a subshell value of 4 is the largest needed to describe a particular subshell. If chemists
ever need a larger value, they can create subshell numbers and letters.
The following figure shows the shapes of the s, p, and d orbitals.
As shown in the top row of the figure (a), there are two s orbitals — one for energy level 1 (1s) and the other for
energy level 2 (2s). The s orbitals are spherical with the nucleus at the center. Notice that the 2s orbital is larger in
diameter than the 1s orbital. In large atoms, the 1s orbital is nestled inside the 2s, just like the 2p is nestled inside the
3p.
The second row of the figure (b) shows the shapes of the p orbitals, and the last two rows (c) show the shapes of the
d orbitals. Notice that the shapes get progressively more complex.
The magnetic quantum number
The magnetic quantum number is designated as:
This number describes how the various orbitals are oriented in space. The value of this number depends on the value
of l. The values allowed are integers from –l to 0 to +l. For example, if the value of l = 1 (p orbital), you can write three
values for this number: –1, 0, and +1. This means that there are three different p subshells for a particular orbital. The
subshells have the same energy but different orientations in space.
The second row (b) of the figure shows how the p orbitals are oriented in space. Notice that the three p orbitals
correspond to magnetic quantum number values of –1, 0, and +1, oriented along the x, y, and z axes.
The spin quantum number
The fourth and final quantum number is the spin quantum number, designated as:
This number describes the direction the electron is spinning in a magnetic field — either clockwise or
counterclockwise. Only two values are allowed: +1/2 or –1/2. For each subshell, there can be only two electrons, one
with a spin of +1/2 and another with a spin of –1/2.
Energy Sub levels (Subshells) and
Oribtals
By merryann12
Energy Sublevels (Subshells) and Orbitals
The principal energy levels contain sublevels designated by the letters s, p, d, f. Just as successive
shells have higher energy, successive subshells also have higher energy. The region of space
where there is a significant probability of finding a particular electron is known as an orbital. It is
sometimes referred to as an electron cloud. The shape of this orbital is illustrated as a solid sphere,
All of the orbitals in a specific subshell have the same energy. For atoms with more than one
electron, the enery of the subshells within a shells is not the same and increases in order: s < p < d <
f. Notice that the 4s energy sublevel is lower in energy than the 3d. An electron spins on tis own axis
in one of only two direction, clockwise or counterclockwise. As a result, only two electrons can
occupy the same orbital contains two electrons. The maximum number of electrons that can exist in
the sublevels are tow in the s orbital, six in the three p orbital, ten in the five d orbital, and fourteen in
the seven f orbitals. This principle is identified as Paul's exclusion princple. Not all principal energy
levels contain each type of sublevel. The following are rules to determine what types of sublevels
occu in any given enegy level and the maximum number of electrons possible in that energy level.
(1) No more than two electrons can occupy one orbital. (2) Electrons occupy the lowest possible
energy sublevels; they enter a higher sublevel only when the lower sublevels are filled. (3) Orbitals in
a given sublevel of equal energy are each occupied by a single electron before a second electron
enters them.
Hund's rule tells us that when electrons have more than one equivalent orbital available, they will
half-fill each of the equivalent orbitals before filling the second half of each. Oxygen has eight
electrons, and tis electronic configuration is 1s2 2s2 2p4, In accordance with Hund's rule, the four 2p
electrons half fill each orbital before pairing up. The spdf atomic orbitals have definite orientations in
space. They are represented by particular spatial shaped. The s orbitals are spherical and the p
orbitals are dumbbell-shaped.
In atomic physics and quantum chemistry, the electron configuration is the distribution
of electrons of an atom ormolecule (or other physical structure) in atomic or molecular orbitals.[1] For
example, the electron configuration of theneon atom is 1s2 2s2 2p6.
Electronic configurations describe electrons as each moving independently in an orbital, in an
average field created by all other orbitals. Mathematically, configurations are described by Slater
determinants or configuration state functions.
According to the laws of quantum mechanics, for systems with only one electron, an energy is
associated with each electron configuration and, upon certain conditions, electrons are able to move
from one configuration to another by the emission or absorption of a quantum of energy, in the form
of a photon.
Knowledge of the electron configuration of different atoms is useful in understanding the structure of
the periodic table of elements. The concept is also useful for describing the chemical bonds that hold
atoms together. In bulk materials, this same idea helps explain the peculiar properties
of lasers and semiconductors.
Bonding & Molecular Structure
Structure & Bonding
The study of organic chemistry must at some point extend to the molecular level, for the
physical and chemical properties of a substance are ultimately explained in terms of the
structure and bonding of molecules. This module introduces some basic facts and principles
that are needed for a discussion of organic molecules.
Electronic Configurations
Electron Configurations in the Periodic Table
1A
2A
3A
4A
5A
6A
7A
1
H
1s1
8A
2
He
1s2
3
Li
1s2
2s1
4
Be
1s2
2s2
5
B
1s2
2s22p1
6
C
1s2
2s22p2
7
N
1s2
2s22p3
8
O
1s2
2s22p4
9
F
1s2
2s22p5
10
Ne
1s2
2s22p6
11
Na
[Ne]
3s1
12
Mg
[Ne]
3s2
13
Al
[Ne]
3s23p1
14
Si
[Ne]
3s23p2
15
P
[Ne]
3s23p3
16
S
[Ne]
3s23p4
17
Cl
[Ne]
3s23p5
18
Ar
[Ne]
3s23p6
The periodic table shown here is sever
There are, of course, over eighty other
A complete periodic table, having very
links has been created by Mark Winter
the right.
Other interactive periodic tables provid
data for each element, including nuclid
environmental and health factors, pres
languages and much more.
For comic relief you may wish to exam
linked to element references in comic b
Four elements, hydrogen, carbon, oxygen and nitrogen, are the major components of most
organic compounds. Consequently, our understanding of organic chemistry must have, as a
foundation, an appreciation of the electronic structure and properties of these elements. The
truncated periodic table shown above provides the orbital electronic structure for the first
eighteen elements (hydrogen through argon). According to the Aufbau principle, the
electrons of an atom occupy quantum levels or orbitals starting from the lowest energy
level, and proceeding to the highest, with each orbital holding a maximum of two paired
electrons (opposite spins).
Electron shell #1 has the lowest energy and its s-orbital is the first to be filled. Shell #2
has four higher energy orbitals, the 2s-orbital being lower in energy than the three 2porbitals. (x, y & z). As we progress from lithium (atomic number=3) to neon (atomic
number=10) across the second row or period of the table, all these atoms start with a filled
1s-orbital, and the 2s-orbital is occupied with an electron pair before the 2p-orbitals are
filled. In the third period of the table, the atoms all have a neon-like core of 10 electrons,
and shell #3 is occupied progressively with eight electrons, starting with the 3s-orbital. The
highest occupied electron shell is called the valence shell, and the electrons occupying this
shell are called valence electrons.
The chemical properties of the elements reflect their electron configurations. For example,
helium, neon and argon are exceptionally stable and unreactive monoatomic gases. Helium
is unique since its valence shell consists of a single s-orbital. The other members of group 8
have a characteristic valence shell electron octet (ns2 + npx2 + npy2 + npz2). This group
of inert (or noble) gases also includes krypton (Kr: 4s2, 4p6), xenon (Xe: 5s2, 5p6) and radon
(Rn: 6s2, 6p6). In the periodic table above these elements are colored beige.
The halogens (F, Cl, Br etc.) are one electron short of a valence shell octet, and are among
the most reactive of the elements (they are colored red in this periodic table). In their
chemical reactions halogen atoms achieve a valence shell octet by capturing or borrowing
the eighth electron from another atom or molecule. Thealkali metals Li, Na, K etc. (colored
violet above) are also exceptionally reactive, but for the opposite reason. These atoms have
only one electron in the valence shell, and on losing this electron arrive at the lower shell
valence octet. As a consequence of this electron loss, these elements are commonly
encountered as cations (positively charged atoms).
The elements in groups 2 through 7 all exhibit characteristic reactivities and bonding
patterns that can in large part be rationalized by their electron configurations. It should be
noted that hydrogen is unique. Its location in the periodic table should not suggest a kinship
to the chemistry of the alkali metals, and its role in the structure and properties of organic
compounds is unlike that of any other element.
Bonding & Valence
Chemical Bonding and Valence
As noted earlier, the inert gas elements of group 8 exist as monoatomic gases, and do not
in general react with other elements. In contrast, other gaseous elements exist as diatomic
molecules (H2, N2, O2, F2 & Cl2), and all but nitrogen are quite reactive. Some dramatic
examples of this reactivity are shown in the following equations.
2Na + Cl2
2NaCl
2H2 + O2
2H2O
C + O2
CO2
C + 2F2
CF4
Why do the atoms of many elements interact with each other and with other elements to
give stable molecules? In addressing this question it is instructive to begin with a very
simple model for the attraction or bonding of atoms to each other, and then progress to
more sophisticated explanations.
Ionic Bonding
When sodium is burned in a chlorine atmosphere, it produces the compound sodium
chloride. This has a high melting point (800 ºC) and dissolves in water to to give a
conducting solution. Sodium chloride is an ionic compound, and the crystalline solid has the
structure shown on the right. Transfer of the lone 3s electron of a sodium atom to the halffilled 3p orbital of a chlorine atom generates a sodium cation (neon valence shell) and a
chloride anion (argon valence shell). Electrostatic attraction results in these oppositely
charged ions packing together in a lattice. The attractive forces holding the ions in place can
be referred to as ionic bonds.
By clicking on the NaCl diagram, a model of this crystal will be displayed and may be
manipulated.
Covalent Bonding
The other three reactions shown above give products that are very different from sodium
chloride. Water is a liquid at room temperature; carbon dioxide and carbon tetrafluoride are
gases. None of these compounds is composed of ions. A different attractive interaction
between atoms, called covalent bonding, is involved here. Covalent bonding occurs by a
sharing of valence electrons, rather than an outright electron transfer. Similarities in physical
properties (they are all gases) suggest that the diatomic elements H2, N2, O2, F2 & Cl2 also
have covalent bonds.
Examples of covalent bonding shown below include hydrogen, fluorine, carbon dioxide and
carbon tetrafluoride. These illustrations use a simple Bohr notation, with valence electrons
designated by colored dots. Note that in the first case both hydrogen atoms achieve a
helium-like pair of 1s-electrons by sharing. In the other examples carbon, oxygen and
fluorine achieve neon-like valence octets by a similar sharing of electron pairs. Carbon
dioxide is notable because it is a case in which two pairs of electrons (four in all) are shared
by the same two atoms. This is an example of a double covalent bond.
These electron sharing diagrams (Lewis formulas) are a useful first step in understanding
covalent bonding, but it is quicker and easier to draw Couper-Kekulé formulas in which each
shared electron pair is represented by a line between the atom symbols. Non-bonding
valence electrons are shown as dots. These formulas are derived from the graphic
notations suggested by A. Couper and A. Kekulé, and are not identical to their original
drawings. Some examples of such structural formulas are given in the following table.
Common Name
Molecular Formula
Methane
CH4
Ammonia
NH3
Ethane
C2H6
Methyl Alcohol
CH4O
Lewis Formula
Kekulé Formula
Ethylene
C2H4
Formaldehyde
CH2O
Acetylene
C2H2
Hydrogen Cyanide
CHN
Multiple bonding, the sharing of two or more electron pairs, is illustrated by ethylene and
formaldehyde (each has a double bond), and acetylene and hydrogen cyanide (each with a
triple bond). Boron compounds such as BH3 and BF3 are exceptional in that conventional
covalent bonding does not expand the valence shell occupancy of boron to an octet.
Consequently, these compounds have an affinity for electrons, and they exhibit exceptional
reactivity when compared with the compounds shown above.
Valence
The number of valence shell electrons an atom must gain or lose to achieve a valence octet
is called valence. In covalent compounds the number of bonds which are characteristically
formed by a given atom is equal to that atom's valence. From the formulas written above,
we arrive at the following general valence assignments:
Atom
H
C
N
O
F
Cl Br
I
Valence
1
4
3
2
1
1
1
1
The valences noted here represent the most common form these elements assume in
organic compounds. Many elements, such as chlorine, bromine and iodine, are known to
exist in several valence states in different inorganic compounds.
Charge Distribution
Charge Distribution
If the electron pairs in covalent bonds were donated and shared absolutely evenly there
would be no fixed local charges within a molecule. Although this is true for diatomic
elements such as H2, N2 and O2, most covalent compounds show some degree of local
charge separation, resulting in bond and / or molecular dipoles. A dipole exists when the
centers of positive and negative charge distribution do not coincide.
Formal Charges
A large local charge separation usually results when a shared electron pair is donated
unilaterally. The three Kekulé formulas shown here illustrate this condition.
In the formula for ozone the central oxygen atom has three bonds and a full positive charge
while the right hand oxygen has a single bond and is negatively charged. The overall
charge of the ozone molecule is therefore zero. Similarly, nitromethane has a positivecharged nitrogen and a negative-charged oxygen, the total molecular charge again being
zero. Finally, azide anion has two negative-charged nitrogens and one positive-charged
nitrogen, the total charge being minus one.
In general, for covalently bonded atoms having valence shell electron octets, if the number
of covalent bonds to an atom is greater than its normal valence it will carry a positive
charge. If the number of covalent bonds to an atom is less than its normal valence it will
carry a negative charge. The formal charge on an atom may also be calculated by the
following formula:
Polar Covalent Bonds
Electronegativity Values
for Some Elements
H
2.20
Li
Be
B
C
N
O
F
Because of their differing nuclear
0.98
1.57
2.04
2.55
3.04
3.44
3.98
charges, and as a result of shielding
by inner electron shells, the different
atoms of the periodic table have
Na
Mg
Al
Si
P
S
Cl
different affinities for nearby electrons.
0.90
1.31
1.61
1.90
2.19
2.58
3.16
The ability of an element to attract or
hold onto electrons is
called electronegativity. A rough
K
Ca
Ga
Ge
As
Se
Br
quantitative scale of electronegativity
values was established by Linus
0.82
1.00
1.81
2.01
2.18
2.55
2.96
Pauling, and some of these are given
in the table to the right. A larger
number on this scale signifies a greater affinity for electrons. Fluorine has the greatest
electronegativity of all the elements, and the heavier alkali metals such as potassium,
rubidium and cesium have the lowest electronegativities. It should be noted that carbon is
about in the middle of the electronegativity range, and is slightly more electronegative than
hydrogen.
When two different atoms are bonded covalently, the shared electrons are attracted to the
more electronegative atom of the bond, resulting in a shift of electron density toward the
more electronegative atom. Such a covalent bond is polar, and will have a dipole (one end
is positive and the other end negative). The degree of polarity and the magnitude of the
bond dipole will be proportional to the difference in electronegativity of the bonded atoms.
Thus a O–H bond is more polar than a C–H bond, with the hydrogen atom of the former
being more positive than the hydrogen bonded to carbon. Likewise, C–Cl and C–Li bonds
are both polar, but the carbon end is positive in the former and negative in the latter. The
dipolar nature of these bonds is often indicated by a partial charge notation (δ+/–) or by an
arrow pointing to the negative end of the bond.
Although there is a small electronegativity difference between carbon and hydrogen, the C–
H bond is regarded as weakly polar at best, and hydrocarbons in general are considered to
be non-polar compounds.
The shift of electron density in a covalent bond toward the more electronegative atom or
group can be observed in several ways. For bonds to hydrogen, acidity is one criterion. If
the bonding electron pair moves away from the hydrogen nucleus the proton will be more
easily transfered to a base (it will be more acidic). A comparison of the acidities of methane,
water and hydrofluoric acid is instructive. Methane is essentially non-acidic, since the C–H
bond is nearly non-polar. As noted above, the O–H bond of water is polar, and it is at least
25 powers of ten more acidic than methane. H–F is over 12 powers of ten more acidic than
water as a consequence of the greater electronegativity difference in its atoms.
Electronegativity differences may be transmitted through connecting covalent bonds by
an inductive effect. Replacing one of the hydrogens of water by a more electronegative
atom increases the acidity of the remaining O–H bond. Thus hydrogen peroxide, HO–O–H,
is ten thousand times more acidic than water, and hypochlorous acid, Cl–O–H is one
hundred million times more acidic. This inductive transfer of polarity tapers off as the
number of transmitting bonds increases, and the presence of more than one highly
electronegative atom has a cumulative effect. For example, trifluoro ethanol, CF 3CH2–O–
H is about ten thousand times more acidic than ethanol, CH3CH2–O–H.
Excellent physical evidence for the inductive effect is found in the influence of electronegative
atoms on the nmr chemical shifts of nearby hydrogen atoms.
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Functional Groups
Functional Groups
Functional groups are atoms or small groups of atoms (two to four) that exhibit a
characteristic reactivity when treated with certain reagents. A particular functional group will
almost always display its characteristic chemical behavior when it is present in a compound.
Because of their importance in understanding organic chemistry, functional groups have
characteristic names that often carry over in the naming of individual compounds
incorporating specific groups. In the following table the atoms of each functional group are
colored red and the characteristic IUPAC nomenclature suffix that denotes some (but not
all) functional groups is also colored.
Functional Group Tables
Exclusively Carbon Functional Groups
Group Formula
Class Name
Specific
Example
Alkene
H2C=CH2
Ethene
Ethylene
Alkyne
HC≡CH
Ethyne
Acetylene
Arene
C6H6
Benzene
Benzene
IUPAC Name Common Name
Functional Groups with Single Bonds to Heteroatoms
Group Formula
Class Name
Specific
Example
Halide
H3C-I
Iodomethane
Methyl iodide
Alcohol
CH3CH2OH
Ethanol
Ethyl alcohol
Ether
CH3CH2OCH2CH3
Diethyl ether
Ether
Amine
H3C-NH2
Aminomethane
Methylamine
Nitro Compound
H3C-NO2
Nitromethane
IUPAC Name Common Name
Methyl
mercaptan
Thiol
H3C-SH
Methanethiol
Sulfide
H3C-S-CH3
Dimethyl sulfide
Functional Groups with Multiple Bonds to Heteroatoms
Group
Formula
Class Name
Specific
Example
IUPAC Name
Common Name
Nitrile
H3C-CN
Ethanenitrile
Acetonitrile
Aldehyde
H3CCHO
Ethanal
Acetaldehyde
Ketone
H3CCOCH3
Propanone
Acetone
Carboxylic
Acid
H3CCO2H
Ethanoic Acid
Acetic acid
Ester
H3CCO2CH2CH3
Ethyl ethanoate
Ethyl acetate
Acid Halide
H3CCOCl
Ethanoyl chloride
Acetyl chloride
Amide
H3CCON(CH3)2
Acid
Anhydride
(H3CCO)2O
N,NN,NDimethylethanamide Dimethylacetamide
Ethanoic anhydride
Acetic anhydride
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to whreusch@msu.edu.
These pages are provided to the IOCD to assist in capacity building in chemical education.
05/05/2013