Download 05 sg Periodic Law

Study Guide:
THE PERIODIC TABLE
AND
PERIODIC LAW
STUDY GUIDE 6: PERIODIC LAW
CHEMISTRY
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TABLE OF CONTENTS
I.
THE PERIODIC TABLE - ORGANIZATION .................................................. 2
A.
B.
C.
Mendeleev ............................................................................................................... 2
Henry Moseley ........................................................................................................ 3
Other Periodic Tables ............................................................................................. 3
II. MODERN PERIODIC TABLE ........................................................... 4
A.
B.
Groups, Families & Periods, and Group Numbers ................................................. 4
Classification of Groups ......................................................................................... 4
1. Classification – Metals, Nonmetals, Metalloids ....................................................... 4
2. Classification – Blocks ................................................................................................ 5
3. Classification – Groups or Families .......................................................................... 5
III.
THE PERIODIC TABLE - TRENDS ......................................... 8
A.
Terms ...................................................................................................................... 8
Valence Electrons ....................................................................................................... 8
Octet Rule: .................................................................................................................. 8
Ion: ............................................................................................................................... 8
Anion: .......................................................................................................................... 8
Cation: ......................................................................................................................... 8
Isoelectric: ................................................................................................................... 8
Diagonal Relationship: ............................................................................................... 8
Effective Nuclear Charge: ......................................................................................... 8
B.
C.
D.
E.
IV.
Atomic Radius ........................................................................................................ 9
Ionization Energy .................................................................................................... 9
Electron Affinity ................................................................................................... 10
Electronegativity ................................................................................................... 10
EXPLAINING ELEMENT PROPERTIES USING
ARRANGEMENT OF ELECTRONS ...................................... 11
A.
Explaining the properties of elements using electron configuration .................... 11
1.
2.
3.
Number of Main Energy Levels:.......................................................................... 11
Number of Main Energy Levels:.......................................................................... 11
Net Effective Nuclear Charge: ............................................................................. 11
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I. THE PERIODIC TABLE - ORGANIZATION
A. Mendeleev
By the late 1790’s, the ‘father of chemistry’ Antoine Lavoisier had compiled a list of the 23 known
elements. During this time, scientists were performing rigorous investigations on electricity – e.g.,
Benjamin Franklin. Electricity was used to break apart chemical compounds into their component
elements. Although many elements had been known since prehistoric time (e.g., gold, silver), the number
of known elements expanded greatly during this period. In addition, the advent of the industrial revolution
led to the development of many chemistry-based industries (e.g., dyes, soaps, petrochemicals &
fertilizers) helped fuel the discovery of new elements. By the 1860’s, more than 60 elements were known.
With the relatively large number of known elements, scientists began to search for a systematic
understanding of the elements. In 1860 a group of chemists met at the First International Congress of
Chemists to discuss the use of atomic mass and other issues. At the meeting, the Italian chemist Stanislao
Cannizzaro presented an exacting way to measure atomic mass, enabling chemists to agree on standard
values for the average atomic mass of elements. In 1864, the John Newlands presented the “law of
octaves” – when he arranged the elements by increasing mass, the chemical and physical properties
repeated at every eighth element. However, he was met with considerable criticism – this pattern did not
work for all known elements and many scientists resented the analogy to music as being unscientific.
In 1869, the Russian scientist Dmitri Mendeleev presented his first periodic table.
Mendeleev created a table by creating a series of cards – on each card he wrote the name of
the element and a list of its chemical and physical
properties, and looked for repeating patterns or
trends. When arranged according to atomic mass,
Mendeleev noted a repeating pattern – periodic – in
the chemical and physical properties of the elements.
In fact, the reason that the periodic table is called ‘periodic’ results
from the organization of elements into observable periodicity, or
repeating pattern of chemical and physical properties. Although
Mendeleev ordered the elements by atomic mass, he reversed
iodine (atomic mass 127) with tellurium (atomic mass 128) due to
their pattern of chemical and physical properties. Additionally, he
left several empty spaces that predicted the existence and
properties of three undiscovered elements. The discovery of these
elements – scandium, gallium, and germanium – and that their
predicted properties matched those in reality - led to wide
acceptance of Mendeleev’s periodic table and earned him the title,
“Father of the Periodic Table.” It is interesting that Mendeleev’s
table had only 17 columns because the noble gases had not been
discovered. Their existence was not predicted, or expected, due to
their lack of reactivity.
The periodic table created by Mendeleev worked well but there were two nagging questions: (1) Why
could one organize most of the elements’ properties? (2) What was the basis for the periodicity in the
properties of the elements?
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B. Henry Moseley
In 1911, the English scientist Henry Moseley, a student of Ernst Rutherford, recognized that
the elements periodic trends resulted not from increasing mass but from increasing atomic
number – the number of protons.. This led to the periodic law: The physical and chemical
properties of the elements are functions of their atomic numbers.
C. Other Periodic Tables
The periodic table with which you are most familiar is not the only way to organize the elements. Many
scientists have developed alternative schemes but the most useful is still the conventional system. Notice,
also, that although the two lower tables are in languages other than English (left is Chinese; right is
Korean), the same element symbols are used throughout the world.
II. MODERN PERIODIC TABLE
The periodic table is the most important tool for organizing, remembering, and predicting chemical facts.
A. Groups, Families & Periods, and Group Numbers
Recall that the vertical columns of the periodic table are called groups or families; the horizontal
rows, periods. Elements in the same group have similar chemical and physical properties. There is
always some confusion about the numbering system used to label the groups. The currently accepted
system (IUPAC) labels the columns at the top of the periodic table from 1 to 18a.
B. Classification of Groups
There are many ways to classify the elements. Three common and useful ways are: (a) dividing
the periodic table into metals, nonmetals, and metalloids, (b) by the orbital shape of the valence
electrons, i.e., s-block, p-block, d-block, or f-block), and (c) the number of valence electrons,
which in turn determines much of the chemical and physical properties – by group or family
names.
1. Classification – Metals, Nonmetals,
Metalloids
Recall that the staircase on the right side of
the periodic table separates the metals from
the nonmetals, and many of the elements on
the staircase, itself, are metalloids.
Characteristic properties of the metals,
nonmetals, and metalloids are given in Table
1. Color in the elements with the appropriate
color designations. However, it is important
to keep in mind that these are only general
properties that may, or may not, be present in the material. For example, diamonds are made of
carbon atoms. Carbon is a nonmetal yet diamonds have luster. Identify the metals, nonmetals
and metalloids by coloring in the boxes on the above periodic table. Use the key at the lower left
to identify the group.
Table 1. Properties of Metals, Nonmetals, and Metalloids
Property
o conduct heat & electricity:
o luster (shiny)
o ductile (ability to be drawn out in a wire)
o malleable (ability to be pounded into shape)
o high tensile strength (ability to support mass)
• nonmetals:
• metalloids:
• metals:
Metals
good
high
high
high
high
Metalloids
moderate
moderate
moderate
moderate
moderate
Nonmetals
poor
low
low
low
low
§ most are gases at room temperature
§ typically gain electron(s) to form negative ions (anions)
. ⇒ oxidation states are typically negative
§ used as semiconductors in electronics (e.g., computers, memory chips)
§ 7 elements: boron (5), silicon (14), germanium (32), arsenic (33), antimony (51),
tellurium (52), polonium (84)
§ all are solids at room temperature except mercury (Hg)
§ the more reactive metals react with acids to form hydrogen gas and a salt
§ typically lose electron(s) to become positive ions (cations)
⇒ oxidation states are typically positive
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2. Classification – Blocks
Many properties of elements can be explained by the electron
configurations of the last electron. For example, many of the visually
attractive metals – e.g., iron, copper, cobalt, are associated with the
available d-block electrons of the transition metals. The s- and pblocks are grouped together into the main group elements; the dblock elements, the transition metals; the f-block elements, the inner transition metals.
s-Block elements have metallic characteristics and are softest elements (e.g., can be cut with a
knife). Metals of the p-block (e.g., lead and tin) are harder than s-block elements but typically
softer than d-block metals. Typically, p-block metals are too reactive to be found in free state in
nature (exception: bismuth).
3. Classification – Groups or Families
This is the way that elements are frequently classified.
Recall that the vertical columns of the periodic table are
called groups or families; the horizontal rows, periods.
Elements in the same group have similar chemical and
physical properties. Elements in the same group have
similar properties because they have the same
number of valance electrons.
a. Alkali Metals:
1
§ ns (denoting n as the number of the highest occupied energy level); lose the one
valence electron to form 1+ ions
§ silvery, lustrous appearance
§ soft enough to cut with a knife (even a butter knife)
§ most reactive of the metals; so reactive they are not found in the free (metallic) state in
nature
§ react vigorously with nonmetals
§ react vigorously with water to produce hydrogen gas and basic solutions (alkaline)
§ typically less dense than water
b. Alkaline Earth Metals:
2
§ ns (denoting n as the number of the highest occupied energy level); lose the two
valence electrons to form 2+ ions
§ properties similar to alkali metals but:
§ harder, denser, and stronger than alkali metals
§ higher melting points (alkali metals have typically low melting points)
c. Hydrogen:
1
§ Although hydrogen, like the other elements in group 1, has an ns electron
configuration, it is not an alkali metal – it is a unique element classified by itself.
Th
e
im
ag
e
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d. Transition Elements (or transition metals)
§ The electron configuration within this d-block sometimes deviates from the orderly
progression of filling orbitals. However, the sum of the s- and d- electrons equals the Group
number.
§ Typical properties of metals (e.g., conduct electricity
and
heat)
§ Some are so unreactive as to be found in the free
(metallic) state in nature – e.g., the ‘coinage metals’:
copper, silver, and gold. Platinum, palladium, and
gold are among the least reactive of all the elements.
e. Halogens
§ Most reactive nonmetals. Reactivity results from 7 valence electrons – one short of
noble-gas configuration.
§ React vigorously with most metals to form salts.
f.
§
§
Noble Gases:
John William Strutt (Lord Rayleigh) and Sir William Ramsay discovered argon in 1894.
This was remarkable because the noble gases, up to that time, had escaped detection
because they are unreactive and not easily observable (e.g., colorless gases). In 1868
helium had been detected on the sun, based on its line emission spectrum, but was not
demonstrated to exist on Earth until 1895 by Ramsay. Interesting note: although argon
makes up ~1% of air, it was not discovered until 1894 (Rayleigh and Ramsay). The
name for argon comes from the Greek work argos for ‘inactive.’
Noble gases are unreactive because they have a full shell of valence electrons.
(However classically unreactive, some noble gas compounds have been made. The first
–
+
was in 1962 Bartlett and Lohmann created O2 [PtF6] . )
g. Actinides & Lanthanides:
§ The f-block elements are wedged between Groups 3 and 4 on the periodic table. These
elements are typically grouped together because their properties are so similar.
Elements above neptunium (93) are all man-made.
§ The lanthanides (in the row beginning with lanthanide (57) are shiny metals with
reactivity similar to the alkaline-earth metals (Group 2).
§ The actinides (in the row beginning with actinide (89) are all radioactive. Only the first
four (Th, Pa, U, and Np) are naturally occurring.
h. Classification – States and Structures of Matter
The three common states of matter1 are solid, liquid and gas. All of the metals, except mercury
(Hg), are solids at room temperature. Mercury is a liquid. Most of the nonmetals are gases at
room temperature.
1
There are actually five known states of matter. The other two, plasma and Bose-Einstein condensate, are not
within the scope of this course.
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In addition, many elements can have more than one structural form or allotrope (Table 2). For
example, pure carbon can exist as several different allotropes including amorphous, diamond (like
the jewelry), graphite, or fullerene (Figure 1). The atomic arrangement accounts for the
properties we associate with each – as graphite, or pencil lead, sheets of carbon atoms are easily
shaved off and deposited on the paper as one writes; as diamond, each carbon atom is bonded to
five other atoms in fixed and immovable placement accounting for diamonds being one of the
hardest substances. Amorphous structures are without a definite or fixed shape or form’. In this
form, the atoms are not arranged in a regular pattern. For example, soot, which is the amorphous
form of carbon, the black residue from a wood fire) is easily spread or separated apart and is very
different from the carbon allotropes of graphite, diamond or fullerene.
Table 2. Some Classic Examples of Allotropes
Element
Forms
Carbon
amorphous, graphite, diamond, fullerene
Oxygen
diatomic oxygen (O2), ozone (O3)
Phosphorus white, red, and black. (White phosphorus is poisonous and can ignite
spontaneously when it comes in contact with air. Red phosphorus, used in safety
matches, is not as dangerous or poisonous as white phosphorus. Black phosphorus
is the least reactive form of phosphorus and has no significant commercial uses.)
Sulfur
rhombic (S8, the most stable sulfur allotrope), polymeric
Graphite
Diamond
Fullerene
Figure 1. Three Allotropes of Carbon. (source: http://www.nyu.edu/pages/mathmol/library/carbon/)
Some elements exist as diatomic molecules: nonmetallic elements that, when in their elemental
state, exist in pairs of atoms (Table 3). A helpful mnemonic for remembering the diatomic
molecules is HONClBrIF (pronounced hon-k’l-brif).
Table 3. Diatomic Molecules
Element
Diatomic Molecule
H2
Hydrogen (H)
O2
Oxygen (O)
N2
Nitrogen (N)
Cl2
Chlorine (Cl)
Bromine (Br)
Iodine (I)
Fluorine (F)
Br2
I2
F2
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III. THE PERIODIC TABLE - TRENDS
A. Terms
Valence Electrons: outermost electrons. As we will see, these are the electrons that are available to be
lost, gained, or shared to form chemical compounds. Valence electrons hold chemical
compounds together.
Octet Rule:
atoms tend to gain, lose, or share electrons in order to have a full set of eight valence
electrons. Elements on the right side of the periodic table tend to gain electrons,
becoming anions, in order to acquire a noble gas electron configuration. Elements on the
left side of the periodic table tend to lose electrons, becoming cations, in order to acquire
a noble gas electron configuration.
Ion:
a negatively or positively charged atom or group of atoms. Ionic charge results from the
difference between the number of electrons and the number of protons. For example, an
+
–
oxygen atom (all atoms are neutral) has eight protons and eight electrons (8 + 8 = 0).
However, an oxygen ion has eight protons (otherwise it wouldn’t be oxygen) and ten
+
–
–
electrons, producing a net charge of 2– (8 + 10 = 2 ).
negatively-charged ion. (Pronounced “an-eye-on”; help: the ‘t’ in cation looks like a
‘+’).
positively-charged ion. (Pronounced “cat-eye-on”).
Anion:
Cation:
Isoelectric:
Two or more atoms or groups of atoms having the same electron configuration.
An important consideration for determining, understanding, and predicting the reactivity of an element is
based on its electronic configuration – e.g., whether the atom readily forms a cation, an anion, or stays
neutral. The arrangement and movement (i.e., loss and/or gain) of electrons results from the atom
achieving a noble gas electronic configuration. For example, atomic fluorine has nine protons and nine
electrons. However, it does not have a complete octet of valence electrons. To achieve this, fluorine will
–
gain an electron, forming the ion fluoride (F ). The added electron produces a complete octet of valence
electrons and the entire ion has a noble gas configuration isoelectric with neon. An ion, however, doesn’t
exist by itself – positive cations and negative anions will combine with each other to obtain an overall net
+
–
charge of zero. Hence, Na ions will combine with Cl ions to form a NaCl structure with a zero net
charge.
Diagonal Relationship: The close relationship between elements in adjacent
groups of the periodic table (Figure at right). Often the lightest element in a
group has more in common with the element located diagonal to it – e.g., Li
and Mg, Be and Al, and B and Si – than the element directly below it.
Effective Nuclear Charge: The net positive charge experienced by an electron in a multi-electron atom.
It is involved in determining the atomic radius, ionization energy, etc. (see below), which determine the
atom’s (or ion’s) reactivity. Effective nuclear charge is essentially the charge of a nucleus seen by a
given electron in the atom. The term ‘effective’ is added because other (inner shell) electrons may act to
shield (= shielding effect) some of the effects of the nucleus’s positive charge on outer shell electrons.
This is an important consideration in understanding atomic radius, electronegativity, etc. given below.
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B. Atomic Radius
The atomic radius, like the radius of a circle, is one-half the distance across it.
In formal definition, the atomic radius is the one-half the distance between the
nuclei of identical atoms bounded together. The atomic radius is an important
consideration for the reactivity of an atom.
Atomic Radius Down Group 1
Atomic Radius Across Period 3
The radius of each atom gets larger as one goes down a
group (see left). This is expected: each row adds another
shell of electrons and electron repel electrons. However,
what causes the radius to become smaller as one goes from
left to right across a given row (see below)?
Element
Sodium
Magnesium
Aluminum
Silicon
Phosphorus
Sulfur
Chlorine
Argon
# Electrons
11
12
13
14
15
16
17
18
However, the data shows that the atomic radius decreases as one
goes across a period (see above table and graph for sodium through
argon). The basis for this results from the balance between (1) the
attraction between electrons and the positively charged nucleus and
(2) the repulsion between electrons. As one moves left to right
across a period, one adds not only electrons but also protons. The
increased pull by the more highly charged nucleus on the electrons
is stronger than the repulsion between the electrons. Thus, the
atomic radii across a period decrease.
Another point is the relative size of the atom and its corresponding ion.
When an atom becomes a cation, it loses electron(s), so the radius of
o
+
the atom is larger than the corresponding cation (Na > Na ).
Similarly, an atom becomes an anion, it gains electron(s), so the radius
o
–
of the atom is smaller than the corresponding anion (F < F ).
C. Ionization Energy
An ion is an atom or group of bonded atoms that has a positive or
negative charge. The ionization energy is the energy required to
remove one electron from a neutral atom of an element, or in other
words, the amount of energy required to form an ion.
+
–
0
A + energy → A + e
Radius (pm)
186
160
143
118
108
106
99
97
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This is important because the movement of electrons, much of which results from the formation of ions,
can explain a lot of chemistry. This results from the balance between having the number of electrons
surrounding an atom equal to the number of protons and the formation of a noble-gas configuration. The
trend in the ionization energy is opposite the trend in atomic radius – as the atom gets larger, it takes less
energy to remove an outer electron.
Ionization energy, unless otherwise stated, is assumed to be the first ionization energy. The second
ionization energy is the amount of energy to remove a second electron; the third ionization energy, the
third electron, and so on. Notice in Table 4 that the 1st ionization energy for sodium is the lowest on the
o
+
–
table. This is the easiest electron to remove: Na → Na + e .
It is also possible to remove a second electron from a positive ion. For example, Group 2 atoms typically
lose both s-orbital, valence electrons. Magnesium’s first two electrons are relatively easy to remove –
see, in Table 4, the jump in energy between the 2nd and 3rd electrons when the two valence electrons have
been removed and the 3rd electron represents removal of an electron from a noble-gas configuration of
electrons. Fluorine has a propensity to gain an electron to form a noble-gas electron configuration, so the
ionization energy for any of its electrons is high.
Table 4. Ionization Energies for Selected Elements.
Ionization Energy (kJ/mol)
Na
Mg
F
st
1
496
738
1,681
nd
2
4,562
1,451
3,374
3rd
6,912
7,733
6,050
th
4
9,544 10,540
8,408
5th 13,353 13,628
11,023
D. Electron Affinity
Sometimes thought of as the opposite of ionization energy.
Electron affinity is the change in energy associated with the
gaining of an electron by a neutral atom.
–
–
0
A + e → A + energy
By convention, the quantity of energy released is represented by a negative number; energy absorbed is
represented by a positive number (q.v., thermodynamic chapter).
Generally, more energy is released by the halogens (Group 17). The ease with which halogens gain
electrons accounts for their high reactivity. Generally, electron affinities become more negative (easier to
gain electrons; higher electron affinity) as one goes across a period. Group trends, going down a given
group, are not as regular.
Although the halogens never add more than one electron, it is common for some atoms to gain more than
2–
2 6
one electron. For example, oxygen gains two electrons to form the O ion ([He]2s 2p ), which is
3–
isoelectric with neon; nitrogen gains three electrons (N ) to become isoelectric with neon.
E. Electronegativity
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Electronegativity is an especially important concept in understanding chemical bonding (formation of
chemical compounds). It is the measure of the ability of an atom in a chemical compound to attract
electrons. Electrons are not distributed evenly throughout most chemical compounds. This uneven
distribution of charge has a significant effect on the chemical properties of compounds.
The most electronegative element is fluorine and is assigned an electronegativity value of 4. Other
significantly electronegative elements are nitrogen, oxygen, and the halogens. The alkali and alkaline
earth elements are the least electronegative. Generally, the trend in electronegativity is to decrease, or stay
the same, down a group. Across a period, electronegativity tends to increase.
IV. EXPLAINING ELEMENT PROPERTIES
USING ARRANGEMENT OF ELECTRONS
A. Explaining the properties of elements using electron configuration
Much of how elements behave – chemically and physically, can be explained and predicted by the
patterns of how electrons are arranged in an atom. For example, the overwhelming similarities between
1
1
lithium ([He]2s ) and sodium ([Ne]3s ) can be explained by electron configuration, but so can their
differences. The three factors influencing element behavior are explained below.
1. Number of Main Energy Levels:
As one goes down a group on the periodic table, the number of shells (main energy levels) increases.
Concomitantly, so does the size of the atoms. This means the valence electrons are further from the
nucleus which, in turn, reduces the attractive force holding the valence electrons to the nucleus. These
inner shell (non-valence) electrons produce a ‘shielding effect’ that blocks, or interferes with the
attraction between the positive nucleus and the outermost electrons. Thus, the valence electrons are easier
to remove.
2. Number of Main Energy Levels:
The noble gases (Group 18) have a full set of eight valence electrons2. This makes them very stable and
inert3. Elements react to gain, lose or share electrons in order to acquire eight valence electrons – the
Octet Rule. Atoms that have completely filled energy sublevels (e.g., d-orbitals), as well as those with
half-filled p-, d-, or f-sublevels, also have been found to have more stable electron arrangements.
3. Net Effective Nuclear Charge:
The Effective Nuclear Charge, or kernel charge, is a relative measure of the attraction between the
protons in the nucleus and the valence electrons. To calculate the kernel charge subtract the number of
lower energy level electrons from the number of protons. Going from left to right across a period on the
periodic table, the kernel charge increases from +1 to +8 pulling the outermost electrons closer to the
nucleus. Most transition elements have a +2 kernel charge. Although successive transition metals in a
period add another proton to the nucleus (e.g., lanthanide to mercury), they also add another inner-shell
electron and also increase the shielding effect experienced by the valence electrons.
2
3
Helium has only two valence electrons but it, too, is considered to have a full set of valence electrons because
only two electrons are allowed in the s-shell available to helium.
The noble gases do react but, for our purposes, they are considered inert.
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ENDNOTES:
a
There are three ways of numbering the groups of the periodic table, one using Arabic numerals and the other two
using Roman numerals. The Roman numeral names are the original traditional names of the groups; the Arabic
numeral names are those recommended by the International Union of Pure and Applied Chemistry (IUPAC) to
replace the old names in an attempt to reduce the confusion generated by the two older, but mutually confusing,
schemes.
There is considerable confusion surrounding the two old systems in use (old IUPAC and CAS) that combined the
use of Roman numerals with letters. In the old IUPAC system the letters A and B were designated to the left (A)
and right (B) part of the table, while in the CAS system the letters A and B were designated to main group
elements (A) and transition elements (B). The former system was frequently used in Europe while the latter was
most common in America. The new IUPAC scheme was developed to replace both systems as they confusingly
used the same names to mean different things. (http://en.wikipedia.org/wiki/Periodic_table_group; Nov. 05, 2005)
“The vertical columns of the periodic table are called groups. The way in which the groups are labeled is
somewhat arbitrary, and three labeling schemes are common in use... The [bottom] set of labels, which have A
and B designations, is widely used in North America. Roman numerals, rather than Arabic ones, are often
employed in this scheme. Group 7A, for example, is often labeled VIIA. Europeans use a similar convention that
numbers the columns from 1A through 8A and then from 1B through 8B, thereby giving the label 7B (or VIIB)
instead of 7A to the group headed by fluorine (F). In an effort to eliminate this confusion, the International Union
of Pure and Applied Chemistry (IUPAC) has proposed a convention that numbers the groups from 1 thorugh 18
with no A or B designations, as shown... in the table... (Brown, LeMay and Bursten. Chemistrry: The Central
Science AP Edition, 10th edition. p. 50)
In the European system, the A represents the main group or representative elements; B, the transition metals.
(Glencoe Chemistry: Matter and Change. ©2002. p.155)