6.1 - The Periodic Table:
A History
http://www.woodrow.org/teachers/ci/1992/MENDELEEV.GIF
Jöns Jakob Berzelius 1828
Swedish chemist
- developed a table of
atomic weights
- introduced letters to
symbolize elements

made the task easier
33 elements known by 1800
Jöns Jakob Berzelius
Johann Döbereiner 1829
• German Chemist
• Triads
• 53 known
elements
http://www.glogster.com/media/1/6/49/85/6498532.jpg
History of the Periodic Table
I.
Döbereiner
a) - described triads of elements
(e.g. Cl, Br, I; Ca, Ba, Sr; Li, Na, K)
- first indication that elements are related to one
another
- atomic mass is related to chemical properties –
the mass of the center element was halfway
between the masses of the other two
elements, all three have similar properties
History of the Periodic Table
History of the Periodic Table
1848
57 elements
1860
Karlsruhe Congress
(big Chemistry Conference)
Germany
John Newlands 1865
• English Chemist
• Arranged elements
by atomic mass
• Described the “Rule
of octaves”
• 62 elements
http://www.rsc.org/education/teachers/learnnet/periodictable/scienti
sts/newlands.jpg
Lothar Meyer 1870
• German Chemist
• Arranged elements
based on atomic mass
• Discovered periodic
properties related to
atomic volume
• Established concept of
valency
http://www.chemistrydaily.com/chemistry/Lothar_Meyer
Meyer’s Data
It’s in the Cards Pre-Lab
• Ionization energy =
the amount of energy, in J or kJ, required to remove 1
electron from an atom in the gaseous state
• Atomic radius =
the distance between the nuclei of two adjacent atoms of
the same kind, divided by 2, measured in pm
• Melting point =
the temperature at which a solid becomes a liquid,
measured in oC
It’s in the Cards Pre-Lab
• Average atomic mass =
the weighted average of the masses of all known
isotopes of an element, measured in amu (or g)
• Density =
ratio of mass divided by volume, g/mL or g/cm3
• Electronegativity =
a measure of the relative ability of an atom to
attract electrons in the context of a chemical
bond, Paulings or none
Dmitri Mendeleev
Dmitri Mendeleev 1869
• Russian chemist
• Wrote elements and
properties on notecards
• Arranged by atomic mass
and properties
• Noted repetition of
properties every 8 or 18
elements
http://anhso.net/data/69/X_kun/571478/mendeleev18371.jpg
Dmitri Mendeleev 1869
• Predicted properties of 3
elements!
– eka-aluminum, eka-boron,
eka-silicon
• Problems: Ar/K, Te/I,
Co/Ni
– First element of each pair
has greater atomic mass
Properties of Some Elements
Predicted by Mendeleev
Predicted
Element
Element
and year
discovered
Properties
Predicted
Properties
Observed
Properties
Ekaaluminum
Gallium,
1875
Density of
metal
6.0 g/mL
5.96 g/mL
Melting point Low
30oC
Oxide
formula
Ga2O3
Ea2O3
Properties of Some Elements
Predicted by Mendeleev
Predicted
Element
Element
and year
discovered
Properties
Predicted
Properties
Observed
Properties
Eka-boron
Scandium,
1877
Density of
metal
3.5 g/mL
3.86 g/mL
Oxide
formula
Eb2O3
Sc2O3
Solubility of
oxide
dissolves in
acid
dissolves in
acid
Properties of Some Elements
Predicted by Mendeleev
Predicted
Element
Element
and year
discovered
Properties
Predicted
Properties
Observed
Properties
Eka-silicon
Germanium,
1886
Melting point
High
900oC
Density of
metal
5.5 g/mL
5.47 g/mL
Color of metal Dark gray
Grayish
white
Oxide formula EsO2
GeO2
Density of
oxide
4.7 g/mL
4.70 g/mL
Chloride
formula
EsCl4
GeCl4
Review
• Döbereiner 1829
– Arranged by atomic mass
– Triads: [Cl Br I], [Ca Ba Sr], [Li Na K]
• Newlands 1865
– Arranged by atomic mass
– Rule of Octaves
• Meyer 1870
– Arranged by atomic mass, periodic trend with atomic volume
– Established concept of valency
• Mendeleev
– Arranged by atomic mass
– Repetition every 8 or 18 elements
– Predicted 3 elements not yet discovered: eka-aluminium - gallium, ekasilicon - germanium and eka-boron - scandium
Discovery of the Noble Gases
1890s
• Lord Rayleigh (physicist) and Sir William Ramsay
(chemist)
• 1894 - Argon “the lazy one”, discovered when Ramsay
was trying to isolate nitrogen
• 1895 - Helium – found on earth in uranium minerals
(found in the sun in 1868)
• 1898 - Neon “the new one”
Krypton “the hidden one”
Xenon “the alien one”
• 1910 – Radon
Properties:
Largely unreactive
8 electrons in valence shell
Low boiling and melting points
Nucleus discovered – 1910
Rutherford predicted that
the charge of an atom is
proportional to its mass
Henry Moseley 1913
• English Physicist
• worked with Rutherford –
was given the task of
testing his prediction
about charge vs. mass
• Periodic Law: Properties
of elements are periodic
functions of their atomic
numbers
http://www.explicatorium.com/images/Personalidades/Henry_Moseley.jpg
History of the
Periodic Table
n of emitted X-rays
corresponded to # protons
 atomic number
“Do other properties match
atomic numbers?” Yes!
 arranged the periodic
table by atomic #’s, not
mass
Law of Atomic Numbers
- the properties of elements are periodic
functions of their atomic numbers (not atomic
mass)
 corrected incorrect placement of cobalt and
nickel, and iodine and tellurium
Glenn Seaborg 1940’s
• American Scientist at UC
Berkeley
• Nobel Prize in Physics,
1951
• Discovered 7 elements
beyond U
• Developed actinide series
and added it to PT
• Seaborgium the only
element publicly named
after a living person
Letter to Seaborg
Trends of the Periodic Table
“periodic” = repeating pattern
• Overall theme = electrons’ positions
relative to each other and the nucleus
determine the following properties.
6.3 - Periodic Trends
Trends of the Periodic Table
“periodic” = repeating pattern
Electron configuration
(  reactivity and bonding)
1. Atomic radius
2. Ionization energy
3. Electronegativity
Periodic Trends
The position of a valence electron and the
ability to remove it from an atom are
related to
• the number of protons in the nucleus
• the extent to which the valence electron is
shielded from the positively-charged nucleus by
the negatively-charged core electrons
1. Atomic Radius
Trend across a period: smaller
– Add e- to valence shell, add p+, stronger pull
from nucleus draws e-’s closer.
– Shielding effect is constant across period
– Not as noticeable with heavier elements
Atomic Radius
Atomic Radius
Atomic Radius
1. Which groups and periods of elements are shown in the table of
atomic radii?
2. In what unit is atomic radius measured? Express this unit in m.
3. What are the values of the smallest and largest atomic radii shown?
What elements have these atomic radii?
4. What happens to atomic radii within a period as the atomic number
increases?
5. What accounts for the trend in atomic radii within a period?
6. What happens to atomic radii within a group?
7. What accounts for the trend in atomic radii within a group?
8. a) Divide the atomic radius of Cs by the atomic radius of Li and
round to 2 significant figures. Cs:Li
b) Divide the atomic radius of Cs by the atomic radius of Rn and
round to 2 significant figures. Cs:Rn
c) Summarize your findings about a) and b) here:
Atomic Radius
1.
2.
3.
4.
5.
6.
7.
8.
Which groups and periods of elements are shown in the table of atomic radii? groups 1A-8A;
periods 1-6
In what unit is atomic radius measured? pm Express this unit in m. 10-12 m
What are the values of the smallest and largest atomic radii shown?
What elements have these atomic radii?
31 pm – helium; 265 pm - cesium
What happens to atomic radii within a period as the atomic number increases? The atomic radius
of the elements within a period generally decreases as the atomic number of the elements
increases.
What accounts for the trend in atomic radii within a period? With increasing atomic number, the
increased positive charge of the nucleus pulls more strongly on the outermost electrons, pulling
them closer to the nucleus. The size of the shield stays the same, so becomes less effective.
Consequently, the atomic radius decreases.
What happens to atomic radii within a group? The atomic radius within a group generally increases
as the atomic number of the elements increases.
What accounts for the trend in atomic radii within a group? With increasing atomic number, the
increased pull by the larger positive charge of the nucleus is offset by the outer electrons’ larger
orbitals and by shielding by inner electrons. Consequently, the atomic radius increases.
a) Divide the atomic radius of Cs by the atomic radius of Li and round to 2 significant figures.
Cs:Li 1.7 X
b) Divide the atomic radius of Cs by the atomic radius of Rn and round to 2 significant figures.
Cs:Rn 1.9 X
c) Summarize your findings about a) and b) here:
2. Ionization Energy
•
•
•
the energy required to remove an
electron from an atom in the gas phase
(in J or kJ)
there is a series of ionization energies for
each atom (since > 1 electron can be
removed)
removing each subsequent electron
requires more energy
Diagram from Document Camera
Ionization Energy
Successive ionization energies (kJ/mol)
Element First
Second
Third
Fourth
Na
496
4,562
6,912
9,543
Mg
738
1,451
7,733
10,540
Al
578
1,817
2,745
11,577
Successive Ionization Energies
Successive Ionization Energies
1. What happens to the values of the
successive ionization energies of an
element?
2. How is a jump in ionization energy
related to the valence electrons of the
element?
Successive Ionization Energies
1. What happens to the values of the successive
ionization energies of an element?
The values of the successive ionization
energies increase.
2. How is a jump in ionization energy related to
the valence electrons of the element?
The jump occurs after the valence electrons
have been removed.
First Ionization Energy
Ionization Energies
1. What is meant by first ionization energy?
2. Which element has the smallest first ionization energy?
The largest? What are their values?
3. What generally happens to the first ionization energy of
the elements within a period as the atomic number of the
elements increases?
4. What accounts for the general trend in the first ionization
energy of the elements within a period?
5. Based on the graph, rank the group 2A elements in
periods 2-5 in decreasing order of first ionization energy.
8. What generally happens to the first ionization energy of
the elements within a group as the atomic number of the
elements increases?
9. What accounts for the general trend in the first ionization
energy of the elements within a group?
Ionization Energies
First ionization energy is the energy
required to remove the first electron from a gaseous atom.
1. What is meant by first ionization energy?
2.
3.
4.
5.
6.
7.
Which element has the smallest first ionization energy? The largest? What are their
values? rubidium – about 400 kJ/mol; helium – about 2375 kJ/mol
What generally happens to the first ionization energy of the elements within a period
as the atomic number of the elements increases? The first ionization energy of the
elements within a period generally increases as the atomic number of the elements
increases.
What accounts for the general trend in the first ionization energy of the elements
within a period? With increasing atomic number, the increased positive charge of the
nucleus produces an increased hold on the valence electrons. Consequently, the
first ionization energy increases.
Based on the graph, rank the group 2A elements in periods 1-5 in decreasing order
of first ionization energy. beryllium, magnesium, calcium, strontium
What generally happens to the first ionization energy of the elements within a group
as the atomic number of the elements increases? The first ionization energy of the
elements within a group generally decreases as the atomic number of the elements
increase.
What accounts for the general trend in the first ionization energy of the elements
within a group?
Summary of Trends in
First Ionization Energy
Trend across a period: increases
Trend down a group: decreases
3. Electronegativity
• How much one atom pulls on another
atom’s electrons in a bond
• Only refers to atoms involved in a bond
(molecule or compound).
• Trend across a period: Increases
• Trend down a group: Decreases
Electronegativity
Electronegativity
Increases
Decreases
History of the Periodic Table
Twelve elements have been known since
ancient times.
What do you think they are?
(Name them, use your periodic table to
help you.)
History of the Periodic Table
Why do you think these particular
elements have been known for so long,
while most elements were not discovered
until the 1800s and 1900s?
Overview of the Periodic Table
Metals
Metalloids
Nonmetals
Noble
gases
Overview of the Periodic Table
Metals
Metalloids
1. excellent heat
conductor
2. excellent
electrical conductor
3. lustrous (shiny)
4. malleable, ductile
5. silvery-gray,
except Cu and Au
6. solids at room T,
except Hg
Some properties of
metals, some
properties of
nonmetals
1. moderate
electrical
conductivity
2. appearance –
more like metals –
lustrous, silvery-gray
3. brittle like
nonmetals
4. solids at room T
Nonmetals
1.
2.
3.
4.
5.
6.
poor heat
conductors
poor electrical
conductors
not lustrous
brittle
variety of
colors
gases or brittle
solids at room
T
Noble
gases
1.
extremely
unreactive –
“inert”
2.
rarely form
compounds
with other
elements
3.
colorless,
odorless
gases at room
T
The Periodic Table
The Periodic Table
1.
How many elements are listed in the periodic table? (the one Dr. Hart gave you…)
__________
2.
What is the atomic number of selenium? _________
3.
What is the symbol for palladium? _________
4.
What is the atomic mass of strontium? ________
5.
How are elements that are gases at room temperature designated in the periodic
table? _________________
6.
How many columns of elements does the periodic table contain? ______
7.
What is another name for a column of elements? __________
8.
What two group numbers can be used to designate elements in the second
column of the periodic table? _________
The Periodic Table
1.
How many elements are listed in the periodic table? (the one Dr. Hart gave you…)
___118_______
2.
What is the atomic number of selenium? __34____
3.
What is the symbol for palladium? ___Pd______
4.
What is the atomic mass of strontium? ___87.62 amu or g_____
5.
How are elements that are gases at room temperature designated in the periodic
table? ___their boxes contain a red balloon______________
6.
How many columns of elements does the periodic table contain? ___18___
7.
What is another name for a column of elements? ___group or family_______
8.
What two group numbers can be used to designate elements in the second
column of the periodic table? __group 2A or group 2_______
The Periodic Table
9.
How many rows of elements does the periodic table contain? ___
10.
What is another name for a row of elements? _____________
11.
Which period contains the least number of elements? ______
12.
What element is found in period 4, group 7B? __________
13.
How are metals designated in this periodic table?
__________________________________
How are metalloids designated in this periodic table?
_______________________________
How are nonmetals designated in this periodic table?
_______________________________
What can be said about the electron configurations of all the
elements in a group? _________
14.
15.
16.
The Periodic Table
9.
How many rows of elements does the periodic table contain? _7_
10.
What is another name for a row of elements? period
11.
Which period contains the least number of elements? Period 1
12.
What element is found in period 4, group 7B? manganese
13.
How are metals designated in this periodic table?
Boxes are tinted blue
How are metalloids designated in this periodic table?
Boxes are tinted green
How are nonmetals designated in this periodic table?
Boxes are tinted yellow
What can be said about the electron configurations of all the elements in
a group? Their valence electron configurations are identical
14.
15.
16.
The s-, p-, d-, and f-Block Elements
The s-, p-, d-, and f-Block Elements
1.
2.
3.
4.
5.
6.
7.
8.
What are the four sections, or blocks, of the periodic table?
_____________
What does each block represent?
_________________________________
What do elements in the s-block have in common?
________________
What is the valence electron configuration of each element in
group 1A? ______
What is the valence electron configuration of each element in
group 2A? ______
Why does the s-block span two groups of elements?
______________________
Why does the p-block span six groups of elements?
_______________________
Why are there no p-block elements in period 1?
__________________________
The s-, p-, d-, and f-Block Elements
1.
2.
3.
4.
5.
6.
7.
8.
What are the four sections, or blocks, of the periodic table?
s-, p-, d- and f-blocks
What does each block represent?
The energy sublevel being filled by valence electrons
What do elements in the s-block have in common?
Valence electrons only in the s orbitals
What is the valence electron configuration of each element in group 1A? s1
What is the valence electron configuration of each element in group 2A? s2
Why does the s-block span two groups of elements?
The single s orbital can hold a maximum of two valence electrons
Why does the p-block span six groups of elements?
The three p orbitals can each hold a maximum of two electrons, resulting in
a maximum of six valence electrons, which corresponds to the six columns
spanned by the p-block.
Why are there no p-block elements in period 1?
The p sublevel does not exist for the first principal energy level.
The s-, p-, d-, and f-Block Elements
9.
10.
11.
12.
13.
14.
15.
16.
What is the ending of the electron configuration of each element
in group 4A? _____
What is the electron configuration of neon? __________
In what period does the first d-energy sublevel appear?
__________
Why does the d-block span ten groups of elements?
_________________________
What is the ending of the electron configuration of each element
in group 3B? _____
What is the electron configuration of titanium? _______________
In what period does the first f-energy sublevel appear?
___________
Determine the group, period, and block for the element having the
electron configuration [Xe]4f145d106s26p3.
a. group_____ b. period ______ c. block _____
The s-, p-, d-, and f-Block Elements
9.
10.
11.
12.
13.
14.
15.
16.
What is the ending of the electron configuration of each element
in group 4A? p2
What is the noble gas electron configuration of neon? [He]2s22p6
In what period does the first d-energy sublevel appear? Period 4
Why does the d-block span ten groups of elements?
The five d orbitals can each hold a maximum of two electrons,
resulting in a total of ten possible valence electrons.
What is the ending of the electron configuration of each element
in group 3B? d1
What is the noble gas electron configuration of titanium?
[Ar]4s23d2
In what period does the first f-energy sublevel appear? Period 6
Determine the group, period, and block for the element having the
electron configuration [Xe]4f145d106s26p3.
a. group__5A or 15___ b. period __6____ c. block __p___
Warmup
• Name the four scientists and the scientific
meeting we talked about Wednesday
• Write them down in chronological order, clearly
indicating who came before and who came after
the scientific meeting
• Use a couple of words or a phrase to remind
yourself of their contribution to the history of the
periodic table, to make a connection you will
remember
Electron Configuration
Compare the charges on the ion list with the position
of the element in the periodic table
Electron Configuration
• Noble gas configuration = [core] e-’s
• ‘Outer’ electrons = valence e-’s
• Elements of groups 1A-8A have valence e’s in s and p orbitals
Isoelectronic Series
= a group of ions and atoms that have the same electron configuration
1. Draw the electron configuration of each of the following elements.
2. What ions will they form?
3. When ions, how many electrons does each have?
How many protons?
4. Predict the relative diameters of the members of this isoelectronic
series.
Isoelectronic Series
Element
Electron
config
Ion
O
F
Ne
Na
Mg
Prediction: smallest to largest:
Ion
# e-’s
Ion
# p+
Isoelectronic Series
Element
O
F
Ne
Na
Mg
Electron
config
Ion
Ion
# e-’s
Ion
# p+
1s22s22p4
 1s22s22p6
O2-
10 e-
8 p+
1s22s22p5
 1s22s22p6
F-
10 e-
9 p+
1s22s22p6
 1s22s22p6
Ne
10 e-
10 p+
1s22s22p63s1
 1s22s22p6
Na+
10 e-
11 p+
1s22s22p63s2
 1s22s22p6
Mg2+
10 e-
12 p+
Prediction: smallest to largest: Mg2+ < Na+ < Ne < F-< O2-
Atomic Radius
½ the distance between nuclei in a diatomic molecule
Atomic Radius
Trend down a group: larger
– Valence e-’s farther from nucleus
– Shielding effect (#e-’s between
nucleus and valence electrons)
decreases pull of nucleus on
valence electrons
Ionic Radius
• Cations (+) smaller than original
atom
– remove e-’s  greater pull from
nucleus
• Anions (-) larger than original atom
– Increased repulsion swells the shell
Ionic Radius
Ionic Radius
1. In this table of ionic radii, how is the charge of the ions of elements in
groups 1A-4A related to the group number?
2. a)
b)
c)
d)
e)
Divide the radius of Cs with the radius of its ion:
Divide the radius of Li with the radius of its ion:
Divide the radius of Be with the radius of its ion:
Divide the radius of B with the radius of its ion:
Summarize your findings about a)-d) here:
3. a)
b)
c)
d)
e)
Divide the radius of the F ion with the radius of the neutral F atom:
Divide the radius of the O ion with the radius of the neutral O atom:
Divide the radius of the N ion with the radius of the neutral N atom:
Summarize your findings about a)-c) here:
Compare and contrast 2 e) and 3 d)