1. ELEMENTARY
SUBSTANCES
ELEMENTS AND THE PERIODIC TABLE OF
ELEMENTS
ATOMS: THE SMALLEST PARTICLE OF MATTER
EVIDENCES OF SUBATOMIC PARTICLES
DETERMINATION OF ATOMIC MASSES
THE ELECTRONIC CONFIGURATION BY MAIN
ENERGY LEVEL AND BY ENERGY SUBLEVEL
PERIODICITY AND PERIODIC RELATIONSHIPS
2. ELEMENT
A PURE SUBSTANCE WHICH IS COMPOSED OF A SINGLE TYPE OF
ATOM
IT IS CONSISTS ONLY OF ATOMS THAT ALL HAVE THE SAME NUMBER
OF PROTONS IN THEIR ATOMIC NUCLEI
THE NUMBER OF PROTONS IN THE NUCLEUS IS THE DEFINING
PROPERTY OF AN ELEMENT, AND IS REFERRED TO AS ITS ATOMIC
NUMBER (z).
ALL ATOMS WITH THE SAME ATOMIC
NUMBER ARE ATOMS OF THE SAME
ELEMENT
IT CANNOT BE BROKEN DOWN INTO
SIMPLER SUBSTANCES BY CHEMICAL
MEANS.
WHEN DIFFERENT ELEMENTS UNDERGO
CHEMICAL REACTION, ATOMS ARE
REARRANGED INTO NEW COMPOUNDS
HELD TOGETHER BY CHEMICAL BONDS.
13p
14n
13e
3. ELEMENTS CAN BE EXPRESSED
USING CHEMICAL SYMBOLS.
BORON GOLD TITANIUM SULFUR TANTALUM
BARIUM URANIUM TITANIUM SULFUR TANTALUM
4. EXERCISES. DECIPHER THE FOLLOWING
USING CHEMICAL SYMBOLS
SODIUM
NEON
THORIUM
EINSTEINIUM
PROTACTINI
UM
RHENIUM
SULFUR
PRASEODYM
IUM
INDIUM
CERIUM
SULFUR
LANTHANUM
URANIUM
RHENIUM
NITROGEN
TELLURIUM
MOLYBDENU
M
NICKEL
COBALT
MOLYBDENU
M
NITROGEN
IODINE
CARBON
OXYGEN
CARBON
AMERICIUM
POLONIUM
IRIDIUM
IODINE
SULFUR
NICKEL
IODINE
POTASSIUM
POTASSIUM
IODINE
SULFUR
AMERICIUM
OXYGEN
NITROGEN
TELLURIUM
PROTACTINI
UM
RADIUM
IODINE
SULFUR
OXYGEN
SODIUM,
CARBON,
IODINE,
OXYGEN,
7. METALLIC ELEMENTS
Nearly all are
shiny and
grey-white in
color
Cu, Cs, and Au are
shiny and golden
Reflectivity is
from
intermediate
to typically
high.
Almost all are
solid in form.
Rb, Cs, Fr, Cs and
Hg are liquid at/near
STP
Densities are
generally
high.
with some
exceptions such as
alkali metals
Most metals
are ductile
and malleable.
Some are brittle like
Cr, Mn, Ge, Ru, W,
Os, and Bi
Most have
hexagonal
and cubic
structure at
freezing point.
Close-packed crystal structures
with high coordination numbers
Atomic radius is from
intermediate to very large.
112–298 pm (average 187 pm)
Around half form allotropes
Sn has a metalloid-like allotrope (grey Sn)
which forms below 13.2 °C.
Thermal conductivity is from
medium to high
Melting point is mostly high and
generally expands during
melting.
Enthalpy of fusion is from low
to high.
8. s, p, d, and f
periodic table
block.
Few number
of electrons in
outer s and p.
except for Pd (0),
Sn, Pb, and Fl (with
4), Bi (with 5) and
Po (with 6)
All has free
electrons that
facilitates
electrical and
thermal
conductivity.
Electrical
conductivity
is from
good to
high but
falls
gradually as
temperature
rises (as it
turns to
liquid)
Over-all
metallic
behavior
All metals tend to form cations.
Metals seldom form covalent
compounds (non-covalent bond)
Oxidation number is nearly
always positive
Ionization energy is relatively
low.
Electronegativity is usually low.
They are base-forming
Their names usually end in –um
or –ium with the following
exceptions: Bi, W, Mn, Ni, and
Co.
9. nonMETALLI
C
ELEMENTS
most are colourless or dull
red, yellow, green, or
intermediate shades
C, P, Se, and I are shiny and grey-white
Reflectivity is from zero
or low (mostly) to
intermediate.
most are gases.
C, P, S, Se, and I are solid while Br is liquid.
Their densities are often low.
They are brittle
when solid
Some (C, P, S, and Se)
have non-brittle forms.
Their crystalline structure at
freezing point ranges from
hexagonal (H, He, C, N, and Xe), to
cubic (O, F, Ne, P, Ar, Kr, Xe, and Rn)
and orthorhombic ( S, Cl, Br & I).
They have open structures
& low coordination
numbers.
Atomic radius is from very
small to intermediate
31–120 pm, average 76.4 pm
10. s and p
periodic
table block.
Most high
number of
electrons in the
outer s and p
except H and He.
no, few, or
directionally
confined free
electrons
(generally
hampering electrical
and thermal
conductivity)
Electrical
conductivity
is from poor
to good but
increases as
temperature
rises
Over-all
nonmetallic
behavior
All nonmetals
tend to form
anions.
Nonmetals form
many covalent
compounds
Oxidation number is either
positive or negative.
Ionization energy and
electronegativity are usually high.
Some form
allotropes
while some
(like graphitic C,
black P and grey
Se) are more
metalloidal
or metallic
in nature
Thermal
conductivity
is from almost
negligible to
very high.
They are acid-forming
Their names usually end in –on,
-gen, or –ine with the following
exceptions: S ends with –ur and
Te and He end with -ium
11. METALLOID
S
shiny and grey-
white in color
Reflectivity is
intermediate.
All are solid in
form
Their densities
are lower than
nearby metals but
higher than
nearby nonmetals
Their crystalline
structure at freezing
point ranges from
rhombohedral (B, As,
and Sb), to cubic (Si and
Ge) and hexagonal (Te).
relatively open crystal
structures and medium
coordination numbers[
Atomic radius is from
small to intermediate:
B, Si Ge, As, Sb, and Te
87–123 pm, average 115.5 pm
They are brittle.
all or nearly all
form allotropes
Some (red B and
yellow As) are more
nonmetallic in
nature
Si has high
thermal
conductivity but
mostly are
intermediate.
Melting point
is high and
some contract,
unlike (most)
metals.
12. p periodic
table block.
medium
number (3–7)
in the outer
level
valence
electrons less
freely
delocalized;
considerable
covalent
bonding
present
Electrical
conductivity
is from
intermediate
to good. Most
behaves like
metals.
Over-all
nonmetallic
behavior
Some have tendency to
form anions in water.
Solution chemistry
dominated by formation
and reactions of
oxyanions.
Form salts as well as
covalent compounds
Oxidation number is either
positive or negative.
Intermediate ionization energy and
electronegativity is close to 2 (between 1.9
and 2.2)
13.
14.
15. GIVE THE CHEMICAL NAMES AND
IDENTIFY WHETHER METAL,
NONMETAL OR METALLOID
1. C
2. Ca
3. Cl
4. Ce
5. Cf
6. Cm
7. Co
8. Cr
9. Cs
10. Cu
11. S
12. Sb
13. Sc
14. Sm
15. Sn
16. Sr
17. Mg
18. Mn
19. Mo
20. Mt
17. An element family is a set of elements
sharing common properties. Elements are
classified into families because the three
main categories of elements (metals,
nonmetals and metalloids) are very broad.
The characteristics of the elements in
these families are determined primarily by
the number of electrons in the outer energy
shell.
Element groups, on the other hand, are
collections of elements categorized
according to similar properties. Because
element properties are largely determined
by the behavior of valence electrons,
19. PERIODS
The rows of the table are known as periods. It is in
the succesive periods that we observe the
periodicity of properties of the elements. Each
period has the full range of properties. For example
more metallic elements occur to the left of a period,
and the less metallic elements to the right; or oxides
of the elements to the left are basic and acidic for
elements to the right. The periods are simply
numbered 1 though 7 from the top down
GROUPS
The columns of the table are known
as groups or families. All the elements in a group
have similar properties. Placing elements in groups
is one of the most important ways of classifying
them. There is some variation in properties within a
group, but the changes are relatively small as one
22. The periodic table groups are as follows (in the brackets are
shown the old systems: European and American):
Group 1 (IA): the alkali metals
Group 2 (IIA): the alkaline earth metals
Group 3 (IIIB)
Group 4 (IVB)
Group 5 (VB)
Group 6 (VIB)
Group 7 (VIIB)
Group 8 (VIIIB)
Group 9 (VIIIB)
Group 10 (VIIIB)
Group 11 (IB): the coinage metals (not a IUPAC-recommended
name)
Group 12 (IIB)
Group 13 (IIIA): the boron group
Group 14 (IVA): the carbon group
Group 15 (VA): the pnictogens (not a IUPAC-recommended
name) or nitrogen group
Group 16 (VIA): the chalcogens
23. ALKALI METALS are recognized as a group
and family of elements. These elements are
metals. Hydrogen is not considered an alkali
metal because the gas does not exhibit the
typical properties of the group. However,
under the right conditions of temperature and
pressure, hydrogen can be an alkali metal.
1 valence electron
Soft metallic solids
Shiny, lustrous
High thermal and electrical conductivity
Low densities, increasing with atomic mass
Relatively low melting points, decreasing with
atomic mass
Vigorous exothermic reaction with water to
GROU
P IA (CAS)
GROU
P 1 (IUPAC)
24. ALKALINE EARTH METALS or simply alkaline
earths are recognized as an important group
and family of elements. These elements are
metals. Examples include calcium and
magnesium.
2 valence electrons
Metallic solids, harder than the alkali metals
Shiny, lustrous, oxidize easily
High thermal and electrical conductivity
More dense than the alkali metals
Higher melting points than alkali metals
Exothermic reaction with water, increasing as
you move down the group; beryllium does not
react with water; magnesium reacts only with
steam
GROUP
IIA (CAS)
GROUP
2 (IUPAC)
25. TRANSITION METALS is the largest family of
elements. The center of the periodic table contains
the transition metals, plus the two rows below the
body of the table (lanthanides and actinides) are
special transition metals.
Groups 3-12
The d and f block metals have
2 valence electrons
Hard metallic solids
Shiny, lustrous
High thermal and
electrical conductivity
Dense
High melting points
Large atoms exhibit
GROUP
IIIB
GROUP 3
GROUP
VIIIB
GROUP 8,
9, 10
GROUP IB
/ IIB
GROUP
11, 12
26. BORON GROUP or earth metal family is not
as well-known as some of the other
element families.
3 valence electrons
Diverse properties, intermediate between
those of metals and nonmetals
CARBON GROUP is made up of elements
called tetrels, which refers to their ability
to carry a charge of 4.
4 valence electrons
Diverse properties, intermediate between those
of metals and nonmetals
GROUP
IIIA
GROUP
13
GROUP
IVA
GROUP
14
27. NITROGEN GROUP or pnictogens is a
significant
element family.
5 valence electrons
Diverse properties, intermediate
between those of metals and
nonmetals
OXYGEN GROUP or the chalcogens family.
6 valence electrons
Diverse properties, changing from
nonmetallic to metallic in moving
down the family
GROUP
VA (CAS)
GROUP
15 (IUPAC)
GROUP
VIA (CAS)
GROUP
28. HALOGEN FAMILY is a group of reactive nonmetals.
7 valence electrons
Reactive nonmetals
Melting points and boiling points increase
with increasing atomic number
High electron affinities
Change state as it moves down the family,
with fluorine and chlorine existing as
gases at room temperature while
bromine is a liquid and iodine is a solid
NOBLE GASES are a family of
nonreactive nonmetals.
8 valence electrons
Typically exist as monatomic gases,
although these elements do (rarely) form compounds
Stable electron octet makes nonreactive (inert) under
GROUP
VIIA (CAS)
GROUP
17 (IUPAC)
GROUP
VIIIA (CAS)
GROUP
18 (IUPAC)
29. ATOMS
• incredibly tiny
• numerous
• perpetually in motion
• ageless
• An atom takes part in
chemical reactions
independently.
• An atom can be
divided into a number
of sub-atomic particles -
electron, proton and
neutron.
31. • ELECTRONS
– discovered by Sir John Joseph Thomson in 1897
– are located in an electron cloud, which is the
area
surrounding the nucleus of the atom
– can abbreviated as e-
– have a negative charge that is equal in
magnitude to the positive charge of the protons
– mass is considerably less than that of a proton
or neutron
32. • PROTONS
– discovered by Ernest Rutherford in the year 1919, when he
performed his gold foil experiment.
– exist in a nucleus and have a positive nuclear charge.
– atomic number or proton number is the number of protons
present in an atom
– The atomic number determines an element (e.g., the element
of atomic number 6 is carbon).
33. • NEUTRONS
– were discovered by James Chadwick in 1932, when he
demonstrated that penetrating radiation incorporated beams
of neutral particles
– located in the nucleus with the protons
– Along with protons, they make up almost all of the mass of the
atom
– The number of neutrons is called the neutron number and can
be found by subtracting the proton number from the atomic
mass number.
34. • u is the SI symbol for atomic mass unit.
• The positive charge of protons cancels the negative charge of the
electrons. Neutrons have no charge.
• With regard to mass, protons and neutrons are very similar, and have
a much greater mass than electrons. Compared with neutrons and
protons, the mass of an electron is usually negligible.
• Spin is associated with the rotation of a particle. Protons, neutrons,
and electrons each have a total spin of 1/2.
35. RELATIONSHIP OF THE ATOMIC
MASS AND ATOMIC NUMBER
WITH THE SUBATOMIC PARTICLES
• A = p + n - Atomic mass (A) is the sum of the
number of protons (p) and neutrons (n) in the
nucleus of an atom.
• Z = p - Atomic number (Z) gives the number of
protons (p) in the nucleus. It is what
distinguishes the atoms of a given element.
• n = A – p - The number of neutrons can be found by
subtracting the proton number from the
atomic mass number.
37. ISOTOPES: Not all atoms are alike
• Isotopes are atoms of the same element
with different number of neutrons or
different atomic masses
• Isotopes are atoms with the same
number of neutrons. Isotopes are
different forms of a single element.
39. HOW IS ATOMIC MASS DETERMINED?
A = (IM X FA)1 + (IM X FA)2 + (IM X
FA)3 + .......
IM is the isotopic mass of a naturally occuring
element
FA is the isotopic abundance of a naturally
occuring element
What is the average atomic mass of
magnesium?
ISOTOPE ISOTOPIC
ABUNDANCE
ISOTOPIC MASS
Mg - 24 78.99 % 23.98501417 u
Mg - 25 10.00 % 24.98583692 u
Mg - 26 11.01 % 25.98259292 u
A = (IM X FA)Mg-24 + (IM X FA)Mg-25 + (IM X FA)Mg-26
A = (23.98501417 X 0.7899) + (24.98583692 X 0.1000) +
(25.98259292 X 0.1101)
A =
24.30502986 A = 24.31 u
40. Calculate the atomic mass of potassium given
the following:
K-39 isotopic abundance- 93.12% isotopic
mass- 38.964 u
K-41 isotopic abundance- 6.88% isotopic
mass- 40.962 u
A = (IM X FA)K-39 + (IM X FA)K-41
A = (38.964 X 0. 9312) + (40.962 X 0.0688)
A = 39.102 u A = 39.1 u
Use the atomic masses of each of the two isotopes of
chlorine along with their isotopic abundances to
calculate the average atomic mass of chlorine.
chlorine-35: atomic mass = 34.969 amu and %
abundance = 75.77%
chlorine-37: atomic mass = 36.966 amu and %
abundance = 24.23%
A = (IM X FA)Cl-35 + (IM X FA)Cl-37
A = (34.969 X 0.7577) + (36.966 X 0.2423)
A = 35.457 u A = 35.46 u
41. ELEMENTAL IONS: charged particles
• ELEMENTAL IONS ARE PARTICLES THAT
RESULTED FROM THE LOSS OR GAIN OF
ELECTRON.
• WHEN AN ATOM LOSES ELECTRON IT WILL
BECOME POSITIVELY CHARGED (CATION).
• Na has 11p, 11e-, and 12n. When one of its electron is
transferred to another atom, it will become positively
charged. The number of p and n are the same but the
number of electron becomes 10. The charge of the this ion
can be determined by subtracting the number of electrons
from the number of proton. Charge = p – e- = 11 – 10 = 1
• Na atom Na+ ion
42. • WHEN AN ATOM GAINS ELECTRON IT WILL
BECOME NEGATIVELY CHARGED (ANION).
• Cl has 17p, 17e-, and 18n. When one electron is accepted by
Cl , it will become negatively charged. The number of p and n
are the same but the number of electron becomes 18. The
charge of the this ion can be determined by subtracting the
number of electrons from the number of proton.
• Charge = p – e- = 17 – 18 = -1
• Cl atom Cl- ion
43. PARTICLE p e- n
F fluorine atom 9 9 10
O oxygen atom 8 8 8
N nitrogen atom 7 7 7
S sulfur atom 16 16 16
PARTICLE p e- n
F- fluoride ion 9 10 10
O2- oxide ion 8 10 8
N3- nitride ion 7 10 7
S2- sulfide ion 16 18 16
44. ELECTRON
CONFIGUR
ATION
The electron configuration of an atomic species
(neutral or ionic) allows us to understand the
shape and energy of its electrons. Many general
rules are taken into consideration when assigning
the location of the electron to its prospective
energy state, however these assignments are
arbitrary and it is always uncertain as to which
electron is being described. Knowing the electron
46. The protons and neutrons are found
inside the nucleus of an atom, while
the electrons are constantly moving
around the nucleus. The nucleus and
the electrons interact to form the most
stable arrangement possible, the
ground state, which requires the least
amount of energy.
There are four main concepts that
must remembered in electron
configuration. These are the energy
levels, the enrgy sublevel, the orbitals,
and number of electrons per sublevel.
47. Energy Levels and the
Atomic Model
The concept of energy levels is one part of the
atomic model that is based on a mathematical
analysis of atomic spectra. Each electron in an
atom has an energy signature that is
determined by its relationship with other
negatively charged electrons in the atom and
the positively charged atomic nucleus.
An electron can change energy levels, but
only by steps or quanta, not continuous
increments. The energy of an energy level
increases the further out from the nucleus it
is. The lower the number of a principal energy
level, the closer together the electrons are to
each other and to the nucleus of the atom.
During chemical reactions, it's more difficult to
remove an electron from a lower energy level
48. Rules of Principal Energy
Levels
A principal energy level may contain up to
2n2 electrons, with n being the number of each
level. The first energy level can contain 2(1)2 or
two electrons; the second can contain up to
2(2)2 or eight electrons; the third can contain up
to 2(3)2 or 18 electrons, and so on.
PRINCIPAL ENERGY LEVEL. n SHELL LETTER MAXIMUM NUMBER OF ELECTRONS, 2n2
1 K 2n2 = 2(1)2 = 2(1) = 2
2 L 2n2 = 2(2)2 = 2(4) = 8
3 M 2n2 = 2(3)2 = 2(9) = 18
4 N 2n2 = 2(4)2 = 2(16) = 32
5 O 2n2 = 2(5)2 = 2(25) = 50
6 P 2n2 = 2(6)2 = 2(36) = 72
7 Q 2n2 = 2(7)2 = 2(49) = 98
49. ENERGY
SUBLEVELS/ENERGY
SUBSHELLS
Each energy level, or shell, is divided into
sublevels. The terms sublevel and subshell
are used interchangeably. The sublevels are
represented by the letters s, p, d, and f.
Each energy level has certain sublevels.
ENERGY
LEVELS
ENERGY
SUBLEVELS
n = 1 s
n = 2 s, p
n = 3 s, p, d
n = 4 s, p, d, f
Energy levels that
are higher than four
would contain
additional sublevels
such as g and h. No
known atom, in its
ground state, would
have electrons that
50. ORBITALS
Each sublevel is made up of orbitals.
Each sublevel has a different number
of orbitals.
ENERGY
SUBLEVELS
ORBITALS
s 1
p 3
d 5
f 7
Orbitals are classified
by their
shape. s orbitals are
spherical in shape.
51. Orbitals found in the p subshell are dumbbell, or
figure-eight, shaped. Remember, in
each p subshell there are three orbitals, so there
are three dumbbell shapes. Atoms are three
dimensional objects, and the orbitals are oriented
around the different axes. The image below
shows the three p orbitals and the
complete p sublevel containing all three orbitals.
52. The first principal energy level has one sublevel
that contains one orbital, called the s orbital. The
s orbital can contain a maximum of two electrons.
The next principal energy level contains one s
orbital and three p orbitals. The set of three p
orbitals can hold up to six electrons. Thus, the
second principal energy level can hold up to eight
electrons, two in the s orbital and six in the p
orbital.
The third principal energy level has one s orbital,
three p orbitals, and five d orbitals, which can
each hold up to 10 electrons. This allows for a
maximum of 18 electrons.
The fourth and higher levels have an f sublevel in
addition to the s, p, and d orbitals. The f sublevel
contains seven f orbitals, which can each hold up
to 14 electrons. The total number of electrons in
53. PRINCIPAL
ENERGY LEVEL, n
NUMBER OF
ORBITALS
ORBITALS NUMBER OF
ELECTRONS PER
ORBITAL
1 1 s 2
2 2 s p 2 6
3 3 s p d 2 6 10
4 4 s p d f 2 6 10 14
Electron Notation
The notation used to indicate the type of energy level and the
number of electrons in that level has a coefficient for the
number of the principal energy level, a letter for the sublevel,
and a superscript for the number of electrons located in that
sublevel. For example, the notation 4p3 indicates the fourth
principal energy level, the p sublevel, and the presence of
54. An orbital diagram helps to determine the
electron configuration of an element. An
element’s electron configuration is the
arrangement of the electrons in the shells.
There are a few guidelines for working out
this configuration:
1. Each orbital can hold only two electrons.
Electrons that occur together in an orbital are
called an electron pair.
2. An electron will always try to enter the orbital
with the lowest energy.
3. An electron can occupy an orbital on its own,
but it would rather occupy a lower-energy orbital
with another electron before occupying a higher-
energy orbital. In other words, within one energy
55. THE BUILDING-UP
PRINCIPLE
The Aufbau principle determines an atom’s
electron configuration by adding electrons
to atomic orbitals following a defined set of
rules.
Although the nucleus of an atom is very dense, the
electrons around it can take on a variety of
positions which can be summarized as an electron
configuration. An element’s electron configuration
can be represented using energy level diagrams,
or Aufbau diagrams. The Aufbau principle (from
the German Aufbau meaning “building up,
construction”) describes a model-building method
in which an atom is “built up” by progressively
56. Filling in an Aufbau Diagram
The order in which orbitals are filled is given by the
Madelung rule. The rule is based on the total
number of nodes in the atomic orbital, n + ℓ, which
is related to the energy. In this context, n
represents the principal quantum number and ℓ
represents the azimuthal quantum number. The
values ℓ = 0, 1, 2, 3 correspond to the s, p, d, and f
labels, respectively. According to the principle,
electrons fill orbitals starting at the lowest
available energy states before filling higher states
(e.g., 1s before 2s).
The Madelung rule defines the order in which
atomic orbitals are filled with electrons. Electrons
fill orbitals starting at the lowest available energy
state before filling higher states.
58. The following steps detail how to draw
an Aufbau diagram
1. Determine the number of electrons that
the atom has.
2. Fill the s orbital in the first energy level
(the 1s orbital) with the first two electrons.
2. Fill the s orbital in the second energy level
(the 2s orbital) with the second two
electrons.
3. Put one electron in each of the three p
orbitals in the second energy level (the 2p
orbitals) and then if there are still electrons
remaining, go back and place a second
59. Limitations to Aufbau
The Aufbau principle is based on the idea
that the order of orbital energies is fixed—
both for a given element and between
different elements. This assumption is
approximately true—enough for the principle
to be useful—but not physically reasonable.
It models atomic orbitals as “boxes” of fixed
energy into which at most two electrons can
be placed. However, the energy of an
electron in an atomic orbital depends on the
energies of all the other electrons of the
atom.
60. Electron Configuration Standard
Notation
A special type of notation is used to write an
atom’s electron configuration. The notation
describes the energy levels, orbitals, and the
number of electrons in each. For example,
the electron configuration of lithium is
1s22s1. The number and letter describe the
energy level and orbital, and the number
above the orbital shows how many electrons
are in that orbital. Using standard notation,
the electron configuration of fluorine is
1s22s22p5.
61. Electrons will fill the lowest energy orbitals
first and then move up to higher energy
orbitals only after the lower energy orbitals
are full.. Although the implications are clear
for orbitals of different principal quantum
number (n), which are clearly of different
energy, the filling order is less clear for
degenerate sublevels. For example, for
boron through neon, the electron filling
order of the 2p orbitals follows Hund’s Rule.
Hund’s Rule states that:
Every orbital in a sublevel is singly occupied
before any orbital is doubly occupied.
All of the electrons in singly occupied
orbitals have the same spin.
62.
63. Diamagnetism
Any time two electrons share the same orbital,
their spin quantum numbers have to be different. In
other words, one of the electrons has to be “spin-
up,”, while the other electron is “spin-down,” This
is important when it comes to determining the total
spin in an electron orbital. In order to decide
whether electron spins cancel, add their spin
quantum numbers together. Whenever two
electrons are paired together in an orbital, or their
total spin is 0, they are called diamagnetic
electrons.
64. Paramagnetism
Electrons that are alone in an orbital are
called paramagnetic electrons. Remember
that if an electron is alone in an orbital, the
orbital has a net spin, because the spin of
the lone electron does not get canceled
out. If even one orbital has a net spin, the
entire atom will have a net spin. Therefore,
an atom is considered to be paramagnetic
when it contains at least one paramagnetic
electron. In other words, an atom could
have 10 paired (diamagnetic) electrons,
65. - Periodic trends are specific patterns that are
present in the periodic table that illustrate
different aspects of a certain element,
including its size and its electronic properties.
- Periodic trends, arising from the
arrangement of the periodic table, provide
chemists with an invaluable tool to quickly
predict an element's properties.
- These trends exist because of the similar
atomic structure of the elements within their
respective group families or periods, and
Periodic Trends
66. property describing an atom's ability
to attract and bind with electrons.
Because electronegativity is a
qualitative property, there is no
standardized method for calculating
electronegativity. However, the
most common scale for quantifying
electronegativity is the Pauling
scale, named after the chemist
Linus Pauling. The numbers
assigned by the Pauling scale are
dimensionless due to the qualitative
nature of electronegativity.
Electronegativity values for each
67. From left to right across a period of elements,
electronegativity increases. If the valence shell of an atom
is less than half full, it requires less energy to lose an
electron than to gain one. If the valence shell is more than
half full, it is easier to pull an electron into the valence
shell than to donate one.
From top to bottom down a group, electronegativity
decreases. This is because atomic number increases down
a group, and thus there is an increased distance between
the valence electrons and nucleus, or a greater atomic
68. Important exceptions of the above rules
include the noble gases, lanthanides, and
actinides. The noble gases possess a
complete valence shell and do not usually
attract electrons. The lanthanides and
actinides possess more complicated chemistry
that does not generally follow any trends.
Therefore, noble gases, lanthanides, and
actinides do not have electronegativity values.
As for the transition metals, although they have
electronegativity values, there is little variance
among them across the period and up and
down a group. This is because their metallic
properties affect their ability to attract electrons
as easily as the other elements.
69. Ionization Energy
- is the energy required to remove an
electron from a neutral atom in its
gaseous phase. Conceptually, ionization
energy is the opposite of
electronegativity. The lower this energy
is, the more readily the atom becomes a
cation. Therefore, the higher this energy
is, the more unlikely it is that the atom
becomes a cation. Generally, elements
on the right side of the periodic table
have a higher ionization energy
because their valence shell is nearly
filled. Elements on the left side of the
periodic table have low ionization
70. The ionization energy of the elements within a period
generally increases from left to right. This is due to
valence shell stability.
The ionization energy of the elements within a group
generally decreases from top to bottom. This is due to
electron shielding.
The noble gases possess very high ionization energies because of
their full valence shells as indicated in the graph. Note that helium has
the highest ionization energy of all the elements.
71. Electron Affinity
- is the ability of an atom to accept an electron.
Unlike electronegativity, electron affinity is a
quantitative measurement of the energy change
that occurs when an electron is added to a neutral
gas atom. The more negative the electron affinity
value, the higher an atom's affinity for electrons.
- Electron affinity increases from left to right within
a period. This is caused by the decrease in
atomic radius.
- Electron affinity decreases from top to bottom
within a group. This is caused by the increase in
72. ATOMIC RADIUS
- While the atomic radius can be defined in a number of
different ways, the general atomic radius trend across
the periodic table holds true. The atomic radius for
atoms of an element tends to go up as you move
down a group of elements in the table. The atomic
radius increases as you move down a column
because for every new row of the table a new electron
shell is added to the atom.
- Atomic radius decreases from left to right within a
period. This is caused by the increase in the number
of protons and electrons across a period. One proton
has a greater effect than one electron; thus, electrons
are pulled towards the nucleus, resulting in a smaller
radius.
Atomic radius increases from top to bottom within a
73. Melting Point
- is the amount of energy required to break a bond(s) to
change the solid phase of a substance to a liquid.
Generally, the stronger the bond between the atoms of
an element, the more energy required to break that
bond. Because temperature is directly proportional to
energy, a high bond dissociation energy correlates to a
high temperature. Melting points are varied and do not
generally form a distinguishable trend across the
periodic table.
Metals generally possess a high melting point.
Most non-metals possess low melting points.
The non-metal carbon possesses the highest melting
point of all the elements. The semi-metal boron also
possesses a high melting point.
74. The Metallic Character
- of an element can be defined as how
readily an atom can lose an electron.
From right to left across a period, metallic
character increases because the
attraction between valence electrons and
the nucleus is weaker, enabling an easier
loss of electrons. Metallic character
increases as you move down a group
because the atomic size is increasing.
When the atomic size increases, the
outer shells are farther away. The
principal quantum number increases and
average electron density moves farther
75. from left to right across a period.
This is caused by the decrease in
radius of the atom that allows the
outer electrons to ionize more
readily.
Metallic characteristics increase
down a group. Electron shielding
causes the atomic radius to
increase thus the outer electrons
ionizes more readily than electrons
in smaller atoms.
Metallic character relates to the
ability to lose electrons, and
nonmetallic character relates to the
78. ELECTRON CONFIGURATION BY
MAIN ENERGY LEVEL
THE DISTRIBUTION OR ARRANGEMENT OF ELECTRONS
IN THE ATOM’S SHELL.
ELECTRONS ARE DISTRIBUTED STARTING FROM THE
INNERMOST SHELL.
THERE ARE SEVEN SHELLS OR MAIN ENERGY LEVELS
WHICH ARE REPRESENTED BY EITHER SHELL LETTER
OR SHELL NUMBER.
93. GIVE THE ELECTRON
CONFIGURATION OF THE
FOLLOWING ELEMENTS, GIVE
THE PERIOD NUMBER AND
GROUP NUMBER TO WHERE IT
BELONGS THEN DETERMINE
WHETHER IT IS A METAL OR A
NONMETAL
