Transcript The Periodic Table

The Periodic Table
An introduction to the Periodic
Table- some history, key points,
element properties and
characteristics
Background Terminology
The Periodic Table
 The Periodic Table is divided into horizontal
rows called periods and vertical columns
called groups or families
 Elements in the same group have very
similar chemical properties.
 As one moves across a period, there is a
gradual change from active metal, to metal,
inactive metal, to metalloid, to nonmetal to
and finally noble gas.
Organization of the
Elements on Periodic Table
 Color Coded Periodic Table
Organization of Elements
on the Periodic Table
 Color Coding
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Alkali metals- blue
Alkaline earth metals- red
Transition Metals-bright yellow
Metalloids- dark pink/purple
Halogens- light pink
Nonmetals- green, pink
Inactive Metals- light blue
Noble Gases- Orange
Lanthanide/Actinide Series- Dark Yellow
Hydrogen
 The element hydrogen on the table is
colored green- however keep in mind it
is normally considered an element in
its own group.
 In terms of its physical properties it
is most like a nonmetal, however in
terms of its chemical properties it is
more like a metal.
The Beauty of the Periodic
Table
 The Periodic Table is a tool. Its
arrangement allows one to make
predictions about elements based on
neighboring elements. It is set up in a
systematic fashion in order to view
patterns and trends in element
activity and properties.
Who is the father of the
Periodic Table?
 Dimitri Mendeleev- Russian Chemist
that did most of his work on the
Periodic Table around 1870
 Many people consider his work to be a
key pioneering effort in the
development of modern day chemistry
Earlier groupings
 Mendeleev’s work followed that of:
 J. Dobereiner – triads (e.g. Ca, Sr, Ba)
 J. Newlands – “law of octaves”

Developed a 7 x 7 grid
Mendeleev’s Table
 One of his first versions
Mendeleev’s Early Logic
 Mendeleev initially based his Periodic
Table on how elements reacted with
oxygen. He noticed a pattern, as
illustrated by some aspects of his
reasoning illustrated on the next slide.
Pattern in the Oxides of
Elements
 Note the pattern in the chemical formulae
Mendeleev’s Periodic
Table
 Elements were grouped based on properties.
Elements in the same group had the same
properties and there was a gradual change in
magnitude of the property down the group.
 There was a gradual change in properties as
one moved across the table in chemical
properties (metals, metalloids, nonmetals
and noble gases) and also physical properties
such as density, melting point, and chemical
formula.
Mendeleev’s Periodic
Table
 Criticisms of Mendeleev’s table
included:

There existed some “apparent” gaps or
openings in his table
Mendeleev’s Periodic
Table
 Mendeleev believed that the gaps were
a result of undiscovered elements

Mendeleev went so far as to predict the
general properties of these elements
 How was Mendeleev able to predict
these properties so well?
Predicting Properties Missing
Elements- Ekasilicon
 Note the missing element and the neighboring ones
The Element Ekasilicon
 Mendeleev‘s
Predictions
 Atomic weight 72
 Density 5.5 g/cm3
 Melting pt. 825 oC
 Oxide form EsO2
 Chloride form EsCl4
 Density oxide 4.7
g/cm3
 Germanium’s Actual
Properties
 Atomic weight 72.61
 Density 5.32 g/cm3
 Melting pt. 938 oC
 Oxide Actual
Properties
 form GeO2
 Chloride form GeCl4
 Density of oxide 4.70
g/cm3
The Periodic Law
 Mendeleev summarized his work in the
form of the Periodic Law. The
Periodic Law describes how elements
are arranged on the Periodic Table.
 Properties of the elements when
arranged in order of increasing atomic
mass are periodic functions of their
atomic masses.
The Periodic Law
 What are the two critical attributes
of the Periodic Law?
 Arrangement in order of increasing
atomic mass.
 Properties of elements are periodic
functions .
What is a periodic
function?
 A periodic function simply means that
properties of elements follow a
regularly repeating pattern.
 In the case of the Periodic Table, the
pattern begins with the group IA
element of a given row and ends with
the Noble Gas (Group VIII A) element
of that row.
Periodic Functions
 In period 3, the periodic function
repeats itself every __________
elements.
 In which periods is the periodic
function every 18 elements?_______
 In period 6, the pattern begins with
_____ and ends with ____________.
Periodic Functions
 What kind of change occurs as one
moves across the Periodic Table to the
nature of the elements?
 Is this change more or less consistent
(periodic)?
Problems with Mendeleev’s
Table
 Mendeleev’s table worked extremely well in
most cases. There was a definite repeating
pattern of elements going across a period
and elements in the same group had similar
properties.
 Additionally, Mendeleev was able to predict
amazingly well the properties and
characteristics of then undiscovered
elements based on the properties of nearby
elements.
 Most of the time that is………….
Problems with Mendeleev’s
Table
 In period 4 of today’s Periodic Table,
cobalt and nickel are not in order of
increasing atomic mass- that is Nickel
should precede Cobalt according to
Mendeleev and it does not.
 Likewise, Iodine should precede
Tellurium and it does not.
 Why??
Problems with Mendeleev’s
Table
 HINT: Sodium reacts with iodine to
form NaI, and it reacts with chlorine
to form NaCl. Likewise, sodium reacts
with tellurium to form Na2Te, and with
sulfur to form Na2S.
 Any ideas? Would the same idea most
likely apply for cobalt and nickel?
Problems with Mendeleev’s
Table
 Hopefully, you realized that the
properties of iodine are like those of
chlorine and therefore should be in
that group. Likewise, the properties
of tellurium are like those of sulfur
and it should be in that group.
 Also, the cobalt is more like rhodium,
and nickel is more like palladium.
Problems with Mendeleev’s
Table
 Therefore, Mendeleev decided it was more
important to arrange based on similarity of
properties within a group instead of strict
adherence to increasing atomic mass.
 Mendeleev was never able to determine with
certainty the cause for these problems. He
believed that possibly the atomic masses of
the elements in question were inaccurate.
Moseley’s Modification
 The atomic masses of cobalt and nickel are
very close to each other, as are the atomic
masses of tellurium and iodine. It so
happens that the proportions of isotopes of
these elements caused an unusual atomic
mass pattern.
 Moseley used X-ray diffraction technology
to determine the atomic numbers of the
elements in 1915
Moseley’s Modification
 Therefore, Moseley arranged the elements
in order of increasing atomic number.
 Interestingly enough, the order of the
elements on Moseley’s Periodic Table was
identical to that of Mendeleev’s, except
cobalt, nickel, tellurium and iodine.
 Moseley modified the Periodic Law to read,
properties of the elements are periodic
functions of their atomic numbers.
Moseley’s Modification
 Basically, when arranged in order of atomic
number, cobalt and nickel were in the
correct sequential order in addition to the
correct location based on properties.
 The same is true for tellurium and iodine.
 In summary, by switching to sequencing by
atomic number the four problems caused by
sequencing by atomic mass were corrected.
Moseley : A Historical
Anecdote
 When Moseley did his work (1913)
(discovering atomic numbers and arranging
the Periodic Table by atomic number) he was
only 20.
 Shortly after that England became involved
in World War I, and Moseley was killed at
the battle of Gallipoli at age 22 in 1915.
 As a result of this loss, to this day England
does not permit their scholars to fight in
battle.
Modern Considerations
 Quantum Mechanics and the Wave
Mechanical Model of the atom lead
scientists to develop the actual electron
arrangements in atoms (electron
configurations).
 It was noted, that elements in the same
group have the same valence shell electron
configuration.

The valence shell of an atom is the outermost E
level containing at least one electron
Electron Configuration
Notation and the PT
 Notice the same valence e-config within a group
Electron Dot Notation and
the Periodic Table
 Notice the same electron dot notation for elements
in the same group
Modern Considerations
 Elements in the same group have the
same chemical properties because they
have the same valence shell electron
configuration.
 For example, LiCl, NaCl, KCl, RbCl, and
CsCl, or Li2O, Na2O, K2O, Rb2O, Cs2O
Groups of Elements
 Information regarding the various groups of
elements: alkali metals, alkaline earth
metals, transition metals, chalcogens,
halogens, noble gases, lanthanide and
actinide series (rare earth elements) can be
found in the textbook or at the web sites
below:
 http://www.webelements.com/
 http://www.chemicalelements.com/
PART II
Periodic Trends
Effective Nuclear Charge
Atomic Radius
Shielding Effect
Ionization Energy
Electronegativity
Enduring Understanding
 Many properties of the elements deal with
how well the nucleus holds onto the electron.
 Since the nucleus is positive and electrons
are negative, there is a natural electrostatic
force of attraction that holds the electron
in the atom. (Coulomb’s law)
 Different element characteristics influence
how well an electron is attracted to the
nucleus.
A Simple Analogy
 Picture two magnets- the attraction for the
north pole of one magnet for the south pole
of another magnet is influenced by three
things.
 1.) the strength of the magnet
 2.) the distance between the two poles of
the magnet
 3.) If there is anything between the two
magnets that blocks their attraction.
From Analogy to Atomic
Structure
 The strength of the magnet- corresponds to
how strong the nucleus is- or the number of
protons in the nucleus (Z) - called the
nuclear charge
 The distance between the two polescorresponds to the size of the atom- or the
distance between outermost electron(s) and
the nucleus- called atomic radius

What is the relationship between Coulombic
force and distance?
From Analogy to Atomic
Structure- con’t
 Something that is between the two
magnets that blocks the attractive
force corresponds to layers of
innershell electrons between the
nucleus and outermost electron-this is
called shielding effect
Atomic Radius
 The atomic radius is essentially the
distance between the nucleus and the
outermost electron.
 The greater this distance the less the
attraction between nucleus and
outermost electron.
Shielding Effect
 In any atom, there is an outermost
energy level (known as the valence
shell) and energy levels that are
between the outermost energy level
and the nucleus. (excepting H & He)
 These inner levels of electrons act to
shield the outermost electrons from
the nucleus- thereby decreasing the
force of attraction between nucleus
and outermost electron(s).
Effective Nuclear Charge
 Effective nuclear charge(ENC/Zeff) is
a measure of how well (effective) a
nucleus is at attracting electrons to it.
 As the number of protons increases
across a period- the ENC increases.
The analogy here would be increase in
number of protons means stronger
magnet. If a nucleus has more protons
it is better (more effective) at
attracting electrons.
Factors that Impact ENC
 The effective nuclear charge is based
on the number of protons in the
nucleus.
 Additionally the shielding effect and
the atomic radius impact the effective
nuclear charge.
Periodic Trends in ENC
Across a Period
 Across a period ENC increasesmeaning across a period a nucleus
becomes more effective at attracting
electrons.
 This is so, because the number of
protons increases. This means that the
electrons are pulled closer to the
nucleus making the atom smaller
(radius decreases).
Periodic Trends in ENC
Across a Period
 Also moving across a period- the shielding
effect remains the same (no new energy
levels are added across- meaning number of
inner shells remains constant).
 Therefore, it is easier for an atom to hold
onto and attract other electrons as you
move to the right. On the left hand side of
the periodic table (relatively low number of
protons for the row) it is more difficult for
an atom to hold onto its electrons.
Summary ENC
 Since ENC is impacted by the shielding
effect and radius- it is sometimes
simpler to evaluate ENC based on
shielding effect (SE) and atomic
radius (AR).
 Therefore, we evaluate how well a
nucleus holds onto / attracts electrons
by evaluating ENC.
Atomic radii
 The atomic radii of atoms decreases moving
across a period (l to r) due to increased
nuclear charge and fairly constant shielding.
 The atomic radii increases moving down a
family due to increased shielding and
increased distance from the nucleus.

The increased shielding cancels out the increase
in nuclear charge (Z).
Atomic radius
Trends in periods and groups
Ionic radii
 When atoms lose electrons to form
cations, the atomic radius decreases

There is less electron-electron repulsion
as well as the possible loss of an energy
level which moves the valence shell closer
to the nucleus
Ionic radii cont’d
 If atoms gain electrons in forming
anions, the size of the atom increases

This results from an increase in electron –
electron repulsion which forces the
electrons apart
Chart of ionic radii
Ionic vs Atomic Radii
Cation vs atom radius
Element
Li
Atomic
radius (nm)
0.123
Ionic
Radius (nm)
0.068
Na
0.157
0.095
K
0.2025
0.133
Rb
0.216
0.148
Cs
0.235
0.169
Radius vs Z
Group 1 - alkali metals
Anion vs atom radius
Element
F
Atomic radius Ionic radius in
in nm
nm
0.064
0.136
Cl
0.099
0.181
Br
0.1142
0.196
I
0.1333
0.216
Radii vs atomic
Group 17 - halogens
Ionization Energy
 The energy required to remove the
outermost electron (least tightly held
valence electron) from a gaseous atom.
(First ionization energy)
 The more strongly held the electron- the
higher the ionization energy
 Likewise, the less tightly held the electron
the lower the ionization energy.

M(g) + I.E.1 -------> M(g)+ + e-
Ionization EnergyFactors that Influence
 The further the electron is from the
nucleus (larger atomic radius) the less
tightly held and lower the ionization energy.
 The greater the shielding effect, the less
tightly held the electron, and the lower the
ionization energy.
 The greater the effective nuclear chargethe more tightly held the electron and the
higher the ionization energy.
Periodic Table TrendsIonization Energy
 View the table below
Active MetalsLow Ionization Energy
 Most active metals are in the lower lefttherefore minimal ENC, large SE, and large
AR
 All of these factors cause the electron to
be less tightly held- which causes the metal
to be more active because it loses electrons
more easily.
 Active metals have low Ionization Energies
Active NonmetalsHigher Ionization Energies
 The highest ionization energies are in the
upper right of the Periodic Table.
 The SE is lowest as is the AR- meaning the
electrons are near the nucleus Additionally, the ENC is greatest for the
period.
 All of these factors indicate that the
electron will be tightly held meaning a lot of
energy is required to remove it.
PT Trends- A Closer Look
Ionization Energy
 The figure below gives a nice 3D look
for the representative elements
PT Trends- A Closer Look
Ionization Energy
 The entire Periodic Table
Anomalies in Trend
 An anomaly in trend is a fancy way of
saying an irregularity in trend.
 Unfortunately, there are occasional
anomalies in trends- the Periodic Table
is not as regular as we would like it to
be.
 There are two anomalies in the
Ionization Energy trend.
Anomaly in Ionization
Energy
 Can you find the anomaly?
The Anomalies in
Ionization Energy
 There is a constant drop in IE at
Group IIIA and then again at Group
VIA.
 The atomic structure of these
elements in Group IIIA and VIA is
such that makes it slightly easier for
the element to lose an electron than
the previous element.
Configurational Stability
 Certain electron configurations
provide additional stability to the atom

These configurations are:

Filled valence shells – eight electrons for most
atoms or an octet – s2p6
 The exception is ?
Filled energy sublevels
 ½ filled energy sublevels

Multiple Ionizations
Electron Affinity
Electron affinity is the amount of energy lost or
gained when a gaseous atom acquires (gains) an
electron.
In general, provides an idea of an atom’s desire to
acquire an electron
X(g) + e- ------> X(g)-
+ E
When the electron affinity value is negative, the
capture of the electron is favorable.
Electron affinity cont’d
If the electron affinity value is positive
or “0,” then the atom does not have a
desire to capture or acquire an
electron.
X(g) + e- + E ------> X(g)-
Trend of E. A.
 As can be seen in the table, nonmetals
have a greater desire to capture
electrons than do most metals
 The more negative the value the
greater the desire
 The trend here is not as clear and
regular as that of ionization energy
Trends in electron affinity
E. A. of the first three
periods
Electronegativity
Electronegativity describes the
strength with which an atom attracts
electrons in a bonded situation.
This concept is often applied when
establishing the character of a bond –
extent to which it is ionic or covalent
Ionic vs Covalent
 Ionic bonds result when the
electronegativity of one atom is so
much greater than another atom’s as
to cause a transfer of an electron(s)
 Covalent bonds result when atoms have
somewhat similar electronegativities
so that they share electrons
Electronegativity Trends
The electronegativity generally
increases moving left to right across a
period.
The electronegativity decreases moving
down a group.
E. N. Trends cont’d
One E. N. Scale
Overall PT TrendsSummary
 The ENC/Zeff increases moving left to
right across a period and decreases
moving down a family or group
Ionization energy
 For the most part the IE increases (l
to r) across a period especially for the
representative elements. This is
especially noticeable with periods 2
and 3. The ionization energy for
transition metals, and the lanthanide
and actinide series stays fairly
constant for the most part.
I. E. cont’d
 The ionization energy generally
decreases moving down a family or
group
Atomic size/radius
 Atomic radius decreases moving left
to right across a period
 Atomic radius increases moving down a
family or group
 This trend runs opposite to the other
trends
Electron Affinity
 Nonmetals tend to have greater
electron affinities than the metals
 Many metals have an electron affinity
value of “0” or one that is positive

What does this indicate about metals?
Electronegativity
 Electronegativity increases moving left
to right across a period
 Electronegativity decreases moving
down a family or group


What situation would likely produce an
ionic bond?
What situation would likely produce a
covalent bond?