Periodic Table

Chapter 6
Periodic Table & Periodic Law
6.1 Development of the Modern Periodic
6.2 Classification of the Elements
6.3 Periodic Trends
Section 6.1 Development of the Modern Periodic
The periodic table evolved over time as
scientists discovered more useful ways to
compare and organize the elements.
• Trace the development of the periodic table from the law of
octaves through the current table ordered by atomic number,
including the scientists who contributed to each stage of
• State the periodic law.
• Identify key features of the periodic table.
• Explain the common feature of elements within a group.
Section 6.1 Development of the Modern Periodic
• Identify the portion of the periodic table the terms
“representative elements”, “transition metals”, “inner
transition metals”, “alkali metals”, “alkaline earth metals”,
metalloids, halogens, “noble gases”, “lanthanide series”,
“actinide series”, and “transuranium elements” refer to and
be able to give examples of some characteristics of the
elements found in these regions.
• State the number of naturally occurring elements on Earth,
the total number of elements that are currently formally
recognized as existing, and the names of the 2 most recently
recognized elements.
Section 6.1 Development of the Modern Periodic
• State the typical characteristics of metals, nonmetals and
metalloids and be able to give an example of an element in
each of these categories
• Identify the states and colors of the 4 halogens at room
temperature and describe the trend in reactivity among the
• Identify what is different about copper, gold, and mercury
compared with other transition metals.
Section 6.1 Development of the Modern Periodic
Key Concepts
• The elements were first organized by increasing atomic
mass, which led to inconsistencies. Later, they were
organized by increasing atomic number.
• The periodic law states that when the elements are arranged
by increasing atomic number, there is a periodic repetition
of their chemical and physical properties.
• The periodic table organizes the elements into periods
(rows) and groups (columns); elements with similar
properties are in the same group.
Section 6.1 Development of the Modern Periodic
Key Concepts
• Elements are classified as either metals, nonmetals, or
History of Development
John Newlands
(~ mid 1860s)
• Octave rule for
ordered by
atomic mass
History of Development
Lothar Meyer, Dimitri Mendeleev
• Both demonstrated relationships
between elemental properties and
atomic mass
• Mendeleev given more credit for idea
• Published first table in 1872
Dimitri Mendeleev
First Periodic Table
• Ordered by atomic mass
Predicted existence/properties of
undiscovered elements
• Scandium (Sc)
• Gallium (Ga)
• Germanium (Ge)
Prediction of Germanium Properties
Atomic Mass
Density, g/cm3
Dens. Oxide
BP of chloride
Dens. Chloride
Dirty gray
EsO2: 4.7
EsCl4: < 100
EsCl4: 1.9
Grayish white
GeO2: 4.703
GeCl4: 86 C
GeCl4: 1.887
Henry Moseley (~1913)
Known that some elements in
wrong order
Used term atomic number (AN)
to indicate amount of charge in
nucleus (determined this charge
from positions of spectral lines)
6 years prior to proton
Arrangement by AN fixed
periodic table problems
Contributions to Classification of the
Elements (Table 6.2)
positive charge
Periodic Law
There is a periodic repetition of
chemical and physical properties of the
elements when they are arranged by
increasing atomic number
Elements – Periodic Table
Ordered by atomic number (number of
protons in nucleus)
Columns (vertical) = groups or families
Current IUPAC numbering system is 1-18
Elements in same group tend to have
similar chemical and physical properties
Rows (horizontal) = periods (currently 7)
Table “periodic” because pattern of variation
of chemical and physical properties repeats
in each period
Elements - Newest
Elements with atomic numbers 114 and 116
officially named 5/2012 as Flerovium (Fl)
and Livermorium (Lv), respectively.
Copernicium (Cn), atomic number 112
officially named 2/2010
Elements 113, 115, 117 (newest – April
2010), 118 have claimed to have been
made, but evidence not yet convincing
enough for official recognition by IUPAC
Official current total = 114 elements
Periodic Table of Elements
Each box shows atomic number and the element’s symbol
Newest elements Fl & Lv;
remainder claimed to have been
made but have not been officially
recognized as existing
Periodic Table – Fig. 6.5
Representative Elements
Modern Periodic Table
Representative Elements
• Groups 1-2, 13-18
• Wide range of chemical & physical
Transition Metals
• Groups 3 to 12
General Classifications
• Shiny when smooth
& clean
• Solid at room
• Good conductors of
heat & electricity
• Ductile
• Malleable
Metal Classifications
Alkali Metals – Group 1
Li, Na, K, Rb, Cs, Fr
Very reactive, soft, mostly exist as
compounds with other elements
Metal Classifications
Alkaline Earth Metals – Group 2
Be, Mg, Ca, Sr, Ba, Ra
Also reactive, but not as much as the
alkali metals
Ca & Mg important nutrients (as ions,
not as elements)
Mg alloys useful as light-weight
materials (bikes, laptop cases)
Transition and Inner Transition Metals
Inner transition metals:
Lanthanide & Actinide Series
Metal Classifications
Transition Metals / Inner Transition
Groups 3 to 12
Inner transition metals consist of two
• Lanthanide series - Ce through Lu
• Actinide series – Th through Lr
• All elements past U (AN 92) are
synthetic (man made)
Called transuranim elements
Some Transition Metals
Only transition metals not
having a silver/gray color
An Atypical Transition Metal Mercury
Only transition metal that is
liquid at room temperature
Semi-Metals / Metalloids
The “staircase” – dividing line
• Steps start between boron and aluminum
• Elements on either side of the dividing line
are metalloids except for Al
• Some texts do not include Po
• Most do not include At – highly radioactive,
estimated that total amount in Earth’s crust
<30 g at any time – hard to study
Silicon & germanium most important
Gases or brittle, dull-looking solids
Poor conductors of heat and electricity
Group 16 Nonmetals
O, S, Se, Te
Group 17 (Halogens): F, Cl, Br, I
All diatomic (F2, etc)
States @ RT: F2(g), Cl2(g), Br2(l), I2(s)
Colors: colorless, pale greeen, dark redbrown, very dark violet (almost black)
All reactive, but F2 most, I2 least
Group 18 - Noble Gases
He, Ne, Ar, Kr, Rn
Used in lasers, light bulbs,
signs and certain types of
welding (TIG, argon)
Chapter 6
Periodic Table & Periodic Law
6.1 Development of the Modern Periodic
6.2 Classification of the Elements
6.3 Periodic Trends
Section 6.2 Classification of the Elements
Elements are organized into different blocks
in the periodic table according to their
electron configurations.
• Explain why elements in the same group have similar
• Identify the four blocks of the periodic table based on their
electron configuration.
• Identify the two exceptions in normal orbital filling order in
the period 4 transition metals.
Section 6.2 Classification of the Elements
Key Concepts
• The periodic table has four blocks (s, p, d, f).
• Elements within a group have similar chemical properties.
• The group number for elements in groups 1 and 2 equals the
element’s number of valence electrons.
• The energy level of an atom’s valence electrons equals its
period number.
Periodic Table Organization
Basic organization of periodic table by
Mendeleev (~1872) was by recurring
trends in elemental properties
Electron not discovered until 1897
Schrodinger model atom was ~1927
Agreement between property based
table and electron configuration based
table demonstrates influence of
electron configuration on properties
Organizing Elements by
Electron Configuration
Group 1 – see table 6.3
Peri Elem
od ent
Na 1s22s22p63s1
N. Gas
Organizing Elements by
Electron Configuration
Group 1 – see table 6.3
Peri Element
N. Gas
# Valence
Configuration Electrons
Organizing Elements by
Electron Configuration
Elements in a group have similar
chemical properties because they have
the same number of valence electrons
• Group 1
• Group 2
• Group 13
• Group 14
Organizing Elements by
Electron Configuration
# of valence electrons = # of electrons
in highest principal energy level
= group number (groups 1 and 2)
= group number – 10 (groups 13 to 18)
He (18) exception (2 valence electrons)
For transition metals, # valence
electrons can vary and may not be
simply related to the group number
Organizing Elements by
Electron Configuration
By definition, energy level of an
element’s valence electrons = period of
[Ar]3d104s24p1 period 4
s-, p-, d-, and f- Block Elements
Block: section of the periodic table that
corresponds to the energy sublevel
being filled with valence electrons
• s, p, d, and f sublevels
s-block elements
Groups 1 and 2
• 1 s1
• 2 s2
Can only have 2 s-block groups
because s sublevel only holds 2
p-block elements
Groups 13 to 18 (except He)
s sublevel is already filled (s2)
13 p1
14 p2
15 p3
16 p4
17 p5
18 p6 completely filled s & p - very
Representative Elements
The representative elements are the
s- and p- block elements
d-block elements
Same as transition metals
In period n
• ns2 (filled s sublevel)
• Partially filled or filled d orbitals of level
• d sublevel can hold 10 electrons; dblock spans 10 groups on periodic table
d-block elements
Period 4 - Filling the n-1 d sublevel
• Sc [Ar]3d14s2
• Ti
Filling is more or less regular but
exceptions occur
• Cr and Cu irregular because of stability
of filled and half-filled d sublevel
Period 4, d Block Exceptions
Aufbau diagram works to vanadium, AN 23
Half-filled and fully-filled set of d orbitals
have extra energy stability, so chromium is
Cr [Ar]3d54s1 (1/2 filled d)
Not [Ar]3d44s2
Next exception is copper:
Cu [Ar]3d104s1 (filled d)
Not [Ar]3d94s2
f-block elements
f sublevel can hold 14 electrons; f-block
spans 14 columns of periodic table
Electrons don’t fill orbitals in a predictable
f-block Inner Transition Metals
For period n (n = 6 or 7)
• Filled ns (ns2); Filled or partially filled (n-2) f (4f
& 5 f)
• First member of lanthanides (La, AN 57) &
actinides (Ac, AN 89) not in Aufbau order –
have d1 configuration, expect f1
• f sublevel filled one element prior to last
member of either row
Aufbau Order & Periodic Table
Can “read” Aufbau order (page 160) directly
from periodic table and knowledge of how
blocks fill
Move left to right
Keep period number for representative
elements (groups 1-2, 13-18)
(Period number -1) for d block
(Period number -2) for f block (start after La
or Ac to get closer to actual order)
s, p, d, and f blocks
s, p, d, and f blocks
p block
s block
d block
f block
Practice problems, page 162
Problems 7-9
Section assessment, page 162
Problems 10-15
Chapter 6
Periodic Table & Periodic Law
6.1 Development of the Modern Periodic
6.2 Classification of the Elements
6.3 Periodic Trends
Section 6.3 Periodic Trends
Trends among elements in the periodic table
include their size and their ability to lose or
attract electrons
• Compare period and group trends of several properties.
• Explain the meaning of electronegativity
• Relate period and group trends in atomic radii, ionic radii
and electronegativity to electron configuration.
• Describe the roles of electron-electron repulsion, electronnucleus attraction, shielding (effective nuclear charge), and
the added stability of favored electron configurations (octet,
half filled and filled sublevels) in determining periodic
property trends.
Section 6.3 Periodic Trends
Trends among elements in the periodic table
include their size and their ability to lose or
attract electrons
• Use Coulomb’s law to explain shielding and its impact on
atomic radii.
• Predict and explain the change in size that occurs when an
atom forms either a cation or an anion.
• Predict and explain the relative changes in ionization energy
for the first, second, third, etc. ionizations of a given atom.
• Explain how departures from overall first ionization energy
trends for some elements can occur in terms of the specific
electron configurations of these elements.
Section 6.3 Periodic Trends
Key Concepts
• Atomic and ionic radii decrease from left to right across a
period, and increase as you move down a group.
• Core electrons are effective at shielding valence electrons
whereas other valence electrons are not effective shielders.
• Ionization energies generally increase from left to right
across a period, and decrease as you move down a group.
• The octet rule states that atoms gain, lose, or share electrons
to acquire a full set of eight valence electrons, which is a
particularly stable electron configuration. Filled and halffilled sublevels (particularly d) also impart extra stability.
• Electronegativity generally increases from left to right
across a period, and decreases as you move down a group.
Section 6.3 Periodic Trends
Key Concepts
• Electronegativity is a property related to bonding and
therefore is a measure of an interaction of the atom with
another atom.
• Disruptions of filled energy levels and of filled sublevels
require more energy than removing electrons from other
electron configurations.
Periodic Trends - Principles
Negative electrons are attracted to the
positive nucleus
Coulomb’s Law: F  (+q) (-q) / r2
F = attractive force
+q = charge on nucleus
-q = electron charge
r = distance between charge
Force vs distance for 1/r2
r (distance)
Periodic Trends - Principles
Coulomb’s Law: F  (+q) (-q) / r2
The greater the nuclear charge (+q),
the more strongly an electron is
The closer an electron is to nucleus
(smaller r), the more strongly it is
Periodic Trends - Principles
Coulomb’s Law: F  (-q)(-q)/r2
Force repulsive if charge same sign
Electrons repelled by other electrons in
an atom
The further away two electrons are
from each other, the weaker the
repulsive force between them
Electrostatic Forces in Atom
Periodic Trends - Shielding
If other electrons are between a valence
electron and nucleus, valence electron will
be less attracted to nucleus
Valence electron
Inner (core) electron
F1  (+3)(-1) / r2
F2  (< +3)(-1) / r2
Both forces are attractive, but F2 is smaller
than F1 due to shielding
Periodic Trends - Shielding
If other electrons are between a
valence electron and nucleus, valence
electron will be less attracted to
nucleus – nuclear charge is shielded
Effective nuclear charge less than full
nuclear charge due to shielding of
charge by negative core electrons
Other valence electrons not effective at
shielding a valence electron
Shielding Effect – Mg: [Ne]3s2
RED – attractive forces
BLUE – repulsive forces
Periodic Trends - Principles
Filled principal energy levels are very
stable (noble gas configuration)
Atoms prefer to add/subtract/share
valence electrons to completely fill a
principal energy level if possible (octet
Completely filled sublevels (s, p, d) also
have extra stability
Atomic Radius
Sizes mostly obtained from crystalline
form of element or from diatomic
In general, are averages of internuclear
separations observed in a variety of
Typically expressed in pm = picometer
= 10-12 m
May also see in
nm = nanometer
= 10-9 m
Atomic Radius
Sodium in
sodium atoms
372 pm
186 pm
Atomic Radius
Hydrogen in gaseous diatomic
74 pm
37 pm
Atomic Radius Trends: Figure 6.11
Atomic Radii Trends: Figure 6.12
Group trend – more shielding, more distant
orbitals with increasing n
Period trend – increasing effective nuclear
Trends – Atomic Radius
Period trend (left to right) – dominated by
increasing effective nuclear charge with no
change in n and with little additional
shielding provided by valence electrons
Group trend (top to bottom) – dominated by
increasing n (orbitals more distant) and
shielding of nuclear charge by core
electrons (from filled energy levels) –
effective nuclear charge decreases
1s, 2s, 3s Orbitals – Distance From
Trends – Ionic Radius
Ion is charged atom - electrons have
been added or removed
• Positive ion (+ ion) formed if electrons
removed – called cation
• Negative ion ( ion) formed if electrons
added – called anion
Trends – Ionic Radius: Cations
When atoms lose electrons and form
positively charged ions (cations), they
always become smaller for 2 reasons:
1. The loss of a valence electron can leave
an empty outer orbital resulting in a
small radius
2. Electrostatic repulsion decreases
allowing the electrons to be pulled
closer to nucleus
Trends –Ionic Radius
Ionization of elemental sodium
 Na+ +
eSodium atom
Sodium ion
Trends – Ionic Radius: Anions
When atoms gain electrons and form
negatively charged ions (anions), they
always become larger because electrostatic
repulsion increases, causing the electrons
to spread apart
If added electron placed in previously
unoccupied energy level, then average
distance from nucleus will be greater for
that electron
Trends –Ionic Radius
Ionization of atomic chlorine
Cl + e
ClChlorine atom
Chlorine ion
[Ne]3s23p6 or [Ar]
Radii in
Ionic Radius – Magnesium oxide
Trends – Ionic Radius
Formation of + ions always results in
decrease in radius
Formation of - ions always results in
increase in radius
Electron configuration of ions follows
same filling order as for neutral atoms
Within a set of + ion or – ions, radius
trends for groups/periods are the same
as for neutral atoms
Ionic Radius Trends
See figure 6.14
Ionic Radius Trends – Fig. 6.15
Within a category (+ or – ions), trend is
same as trend for atomic radius
Comparison of Atomic & Ionic Radii
Ionization Energy (IE)
Energy needed to remove electron from
neutral gaseous atom to form + ion
X(g)  X+(g) + e-
• Ionization potential is per atom value in eV (1
eV=1.602 x 10-19 J)
Must overcome electrostatic attraction
to remove electron
Low IE = easy valence electron loss –
element will readily form a cation
2nd Ionization Energy
Energy needed to remove electron from
ion with single positive charge
X+ (g)  X2+(g) + e-
Must overcome much stronger
electrostatic attraction to remove
second electron compared to the first
Trends – Ionization Energy
Group 1 lowest values and Group 18
highest values within a period
Group 1 likely to form M+1 but unlikely
to form M+2
• Difficult to remove electron from noble
gas configuration (filled p sublevel)
Trends – Ionization Energy (IE)
Period and group trends inverse of radii
trends – increased nuclear charge
pulls electron in tighter
Shape of trend for any given period is
irregular due to stability of filled / half
filled sublevels
Trends –
Ionization Energy
Periods 1-5 Fig. 6.16
Trends – 1st Ionization Energy Fig. 6.17
Trend opposite that for atomic radius
Reasons for trend same as for atomic radius
Trends – Ionization Energy (IE)
2nd IE always > 1st
3rd IE always > 2nd
• Removing electron from more + ion
Once valence electrons removed,
energy always takes big jump
• Must remove electron from filled level
Period 2 Successive Ionization
Energies (Table 6.5)
Trend Exceptions – (IE)
1st IE (values in kJ/mol)
Be (1s22s2) 900 vs B (1s22s22p1) 800
• B has inner 2s2 configuration which
effectively shields 2p1 electron
N (1s22s22p3) 1400 vs O (1s22s22p4)
• Half-filled p sublevel has extra stability
Trend Exceptions – IE
2nd IE
B (1s22s22p1) 2430 vs
C (1s22s22p2) 2350
• 2s2 effectively shields
O (1s22s22p4) 3390 vs
F (1s22s22p5) 3370
• Disrupts/creates 2p3 configuration
Trends - Electronegativity
Relative ability to attract electrons in a
chemical bond
Max 3.98 (F) to min 0.7 (Fr)
Elements with high EN tend to form
negative ions
• F-, Cl-, O2-
Noble gases not tabulated
• Very few compounds to get info from
Decreasing Electronegativity
Trends - Electronegativity
Increasing Electronegativity
electronegativity < 1.0
1.0  electronegativity < 2.0
2.0  electronegativity < 3.0
3.0  electronegativity < 3.9
electronegativity  3.9
Trends – Electronegativity vs AN
Trend Summary (ignore EA)

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