CHAPTER 6
The Periodic Table
and Periodic Law
6.1 Organizing the Elements
Antoine Lavoisier – 1789
33 elements (23 actual elements)
John Dalton – 1808
36 elements
Jöns Berzelius -1814
47 elements
Used letters for symbols
Development of Periodic Table
Law of Triads
Law of Octaves
Elements could be classified into groups of three, or triads. Trends in physical properties such as density, melting
point, and atomic mass were observed.
Arranged the 62 known elements into groups of seven according to increasing atomic mass.
He proposed that an eighth element would then repeat the properties of the first element in the previous group.
J.W. Döbereiner (1829)
J.A.R. Newlands (1864)
Döbereiner’s Triads
Name Atomic
Mass Name
Atomic
Mass Name
Atomic Mass
Calcium 40 Barium 137
Average 88.5
Strontium 87.6
Chlorine 35.5 Iodine 127
Average 81.3
Bromine 79.9
Sulfur 32 Tellurium 127.5
Average 79.8
Selenium 79.2
Johann Döbereiner ~1817
Döbereiner discovered groups of three related elements which he termed a triad.
Newlands Law of Octaves
1 Li Na K John Newlands ~1863Smoot, Price, Smith, Chemistry A Modern Course 1987, page 161
2 Be Mg 3 B Al 4 C Si 5 N P 6 O S 7 F Cl
•
John Newlands
-
1864
– “Law of Octaves”
– Organized by increasing atomic
mass
– Repeating properties
Mendeleev’s Periodic Table
"...if all the elements be arranged in
order of their atomic weights a periodic repetition of properties is obtained."
- Mendeleyev
Dmitri Mendeleev
Russian
Published the periodic table
Organized elements by
similar chemical properties
Arranged elements by increasing atomic mass Predicted existence of
several unknown elements
63 elements Dmitri Mendeleev
Mendeleev’s Table of 1869
In 1875, a French chemist discovered Gallium
(eka-aluminum) and its properties were very close to what
Mendeleev predicted! Mendeleev predicted the existence of
unknown elements like eka-aluminum
Mendeleev’s Revised Table
1871
Elements Properties are Predicted
Property Mendeleev’s Predictions in 1871 Observed Properties
Molar Mass Oxide formula Density of oxide Solubility of oxide
Scandium (Discovered in 1877)
44 g M2O3 3.5 g / ml Dissolves in acids
43.7 g Sc2O3 3.86 g / ml Dissolves in acids Molar mass
Density of metal Melting temperature Oxide formula
Solubility of oxide
Gallium (Discovered in 1875)
68 g 6.0 g / ml
Low M2O3
Dissolves in ammonia solution
69.4 g 5.96 g / ml
30 0C
Ga2O3
Dissolves in ammonia Molar mass
Density of metal Color of metal
Melting temperature Oxide formula
Density of oxide Chloride formula Density of chloride Boiling temperature of chloride
Germanium (Discovered in 1886)
72 g 5.5 g / ml Dark gray
High MO2 4.7 g / ml
MCl4 1.9 g / ml Below 100 oC
71.9 g 5.47 g / ml Grayish, white
900 0C
GeO2 4.70 g / ml
GeCl4 1.89 g / ml
86 0C
O’Connor Davis, MacNab, McClellan, CHEMISTRY Experiments and Principles 1982, page 119,
6.1 Organizing the Elements
Henry Moseley – 1913
Determined the atomic numbers of
elements from their
X
-ray spectra
Rearranged elements by increasing
atomic number
Current periodic table is arranged by
atomic number
6.1 Organizing the Elements
History Review
Lavoisier – list of elements
Berzelius – symbols as letters
Mendeleev – table with similar chemical
properties and increasing atomic mass
Moseley – increasing atomic number
6.1 Organizing the Elements
Modern Periodic Table
Arranged in order of increasing atomic
number
Group (or family) – column
Period – row
6.1 Organizing the Elements
Periodicity is the tendency to recur at
regular intervals.
The pattern of properties within a period
repeats as you move from one period to the
next
Periodic law – when organized by
increasing atomic number, there is a
periodic repetition of chemical and
physical properties.
Why?
What is similar for all elements in
a group or family?
They have the same number of outer level
electrons!
These have 1
These have 3
6.1 Organizing the Elements
Electrons, particularly the
valence electrons, control
many of the chemical and
physical properties of atoms!
Valence Electrons
H Li Na K Be Mg Ca B Al Ga C Si Ge N P As O S Se F Cl Br Ne Ar Kr He Group1A 2A 3A 4A 5A 6A 7A 8A
= valence electron
s1 s2 s2p1 s2p2 s2p3 s2p4 s2p5 s2p6
1 2 13 14 15 16 17 18
Metals, Nonmetals, &
Metalloids
1
2
3
4
5
6
7
Metals
Metalloids
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 349
6.2 Classification of the
Elements
•
Classifying by properties
1. Metals – left of stairs
– Lustrous (shiny)
– Malleable (not brittle)
– Ductile (drawn into a wire)
– Good conductor of heat and electricity
– Vast majority of known elements are metals
2. Nonmetals – right of stairs
– Dull-looking – Brittle
– Poor conductors
3. Metalloids – on the stairs (minus aluminum)
– Properties of both metals and nonmetals
6.2 Classification of the
Elements
Elements can be classified into four
categories based upon their electron
configurations
1.
Noble gases
He, Ne, Ar, Kr, Xe, and Rn Full outer s and p sublevels
2.
Representative elements
Groups 1A – 8A (this includes Noble Gases) Partially filled s and p sublevels
6.2 Classification of the
Elements
3.
Transition metals
d block
Outer s and nearby d sublevel contain
electrons
4.
Inner transition metals
f block
Outer s and nearby f sublevel contain
electrons
Noble Gases
Inner Transition Metals Transition Metals
Groups of Elements
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 367
N 7 P 15 As 33 Sb 51 Bi 83 O 8 S 16 Se 34 Te 52 Po 84 F 9 Cl 17 Br 35 I 53 At 85 He 2 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87 Be 4 Ca 20 Sr 38 Ba 56 Ra 88 Mg 12 1 2 1 2 15 16 17 18 Alkali metals
Alkaline earth metals Nitrogen family
Oxygen family Halogens
Noble gases
13 14 15 16 17
18
6.2 Classification of the
Elements
Families Review
Alkali metals – group 1A (minus hydrogen),
reacts vigorously with water
Alkaline Earth metals – Group 2A
Halogens – group 7A, 17, very reactive gases Noble Gases – group 8A, 0, or 18, unreactive
gases
Other families are just referred to as the
element at the top of the column
Alkali Metals, Group 1
H
N O F
Cl
Br
I
Li
Na
K
Fr
Be
Mg
Ca
Ra Sc
Ac
He
Ne
Ar
Kr
Rn Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Al Si P S
B C
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
The Alkali Metals (Group 1)
- The alkali metals are lithium (Li), sodium (Na),
potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr)
-
Hydrogen is placed in Group 1 but is not a metal.- The alkali metals react readily with
nonmetals to give ions with a +1 charge
- Compounds of alkali metals are common in
nature and daily life.
Chemistry of the Groups
Copyright 2007 Pearson Benjamin Cummings. All rights reserved.
H
1
Li
3
Na
11
K
19
Rb
37
Cs
55
Fr
87 1
Alkaline Earth Metals, Group
2
H
N O F
Cl Br I Li Na K Fr Be Mg Ca Ra Sc Ac He Ne Ar Kr Rn Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Al Si P S
B C
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
The Alkaline Earths (Group 2)
- The alkaline earths are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra).
All are metals that react readily with
nonmetals
Form ions with a +2 charge.
Chemistry of the Groups
Copyright 2007 Pearson Benjamin Cummings. All rights reserved.
Be
4
Ca
20
Sr
38
Ba
56
Ra
88
Mg
12 2
Halogens, Group 17
H
N O F
Cl
Br
I
Li
Na
K
Fr
Be
Mg
Ca
Ra Sc
Ac
He
Ne
Ar
Kr
Rn Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At
Al Si P S B C
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
The Halogens (Group 17, 7A)
- The halogens are fluorine (F), chlorine (Cl), bromine (Br) iodine (), and astatine (At).
- They react readily with metals to form ions with a
1 charge.
Chemistry of the Groups
Copyright 2007 Pearson Benjamin Cummings. All rights reserved.
F 9 Cl 17 Br 35 I 53 At 85 17
Fluorine F 19.0 F2 pale-yellow gas -187.0 Chlorine Cl 35.5 Cl2 greenish-yellow gas -34.5 Bromine Br 79.9 Br2 red-brown liquid 58.0 Iodine I 126.9 I2 black solid (m.p.113oC) 184.0
Element
Element At. Mass Normal Form at STP b.p., At. Mass Normal Form at STP b.p., ooCC
Astatine At (210)
Noble Gases, Group 18
H
N O F
Cl
Br
I Li
Na
K
Fr
Be
Mg
Ca
Ra Sc
Ac
He
Ne
Ar
Kr
Rn
Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te Xe
Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Al Si P S
B C
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
The Noble Gases (Group 18, 8A, 0)
- are helium (He), neon (Ne), argon, (Ar), krypton (Kr), xenon (Xe), and radon (Rn);
- are monatomic;
- are unreactive gases at room temperature and pressure;
- are called inert gases.
Chemistry of the Groups
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
He
2
Ne
10
Ar
18
Kr
36
Xe
54
Rn
86 18
6.3 Periodic Trends
Because the arrangement of elements
on the periodic table is closely linked to
the electron configuration, there are a
number of trends that can be used to
predict chemical and physical behavior.
In order to understand and use these
trends, you must first understand how
NUCLEAR CHARGE and SHELLS AND
SHIELDING influence electron behavior.
Periodic Trends
All trends are explained by two
trends:
•
NUCLEAR CHARGE
•
SHELLS AND SHIELDING
6.3 Periodic Trends
• Nuclear Charge
Period Trend
•As you move across a period or down a group, the atomic
number increases.
•This means the number of protons in the nucleus is
increasing.
•With more protons, the positive pulling strength (nuclear
charge) of the nucleus is increasing Group Trend
•As you move down a group, the atomic number increases. •This means the number of protons in the nucleus is
increasing.
•With more protons, the positive pulling strength (nuclear
charge) of the nucleus is increasing, but you also have to account for shells and shielding
•Therefore, The group trend for Nuclear Charge will not matter!
Nuclear Strength
Moving across the periodic table, nuclear charge increases
The number of core, or innermost, electrons remains the same while the number of protons increases.
Li : [He]
Li : [He] 22ss11
Be:
Be: [He] [He] 22ss22
B :
B : [He] [He]
2
2ss2222pp11
C :
C : [He] [He]
2
2ss2222pp22
N :
N : [He] [He]
2
2ss2222pp33
O :
O : [He] [He]
2
2ss2222pp44
F :
F : [He] [He]
2
2ss2222pp55
[He] represents the core electrons
+++
e
-e- e
- ++++ e
-e
-e
-e
-Li Be B
+++ ++
e
-e- e
-e
-e
Nuclear Strength
Increases
1A
2A 3A 5A 6A 7A
0
*G
e
n
e
ra
lly
In
cr
e
a
se
s
* Does not matter because of shells and shielding
Electron Shells
n: 1 2 3 4
+++e
-e- e
-Li : 1s22s1
Li : [He] 2s1 Na : [Ne] 3s1 K : [Ar] 4s1 Rb : [Kr] 5s1
•
Shells
• Energy levels
• The higher the level, the farther from the nucleus • Across – highest energy does not change
• Down – energy levels increase
+ Shielding
• Inner level electrons interfere or shield the valence electrons from the
nucleus
• Across – shielding is constant
• Down – shielding increases
Li : [He]
Li : [He] 22ss11
Na : [Ne]
Na : [Ne] 33ss11
K : [Ar]
K : [Ar] 44ss11
Rb : [Kr]
Rb : [Kr] 55ss11
Si ze in cr ea se
Electron Shells and Atom Size
n: 1 2 3 4
+++e
-e
-e
-Li :1s22s1
e
-e
-e
Shells and Shielding
Constant
In
cr
e
a
se
s
1A
2A 3A 5A 6A 7A
0
Nuclear Charge verses
Shells and Shielding
Period Trend
Nuclear Strength Wins!!!
Because Shells and Shielding are constant
across a period they don’t affect period trends
Therefore, ALL PERIOD TRENDS are caused
Nuclear Charge verses Shells
and Shielding
Group Trend
Shells and Shielding Win
Even though nuclear charge is increasing,
more Shells (farther from the nucleus) and Shielding (inner level electron interference) decreases the effective nuclear
strength.
Therefore, ALL GROUP TRENDS are caused
by shells and shielding or their effect on the nuclear charge.
Nuclear Charge verses Shells
and Shielding
Nuclear Strength Increases
S /S In cr e a se s m in im iz e s e ff e ct iv e N u cle a r S tr e n g th 1A
2A 3A 5A 6A 7A
0
6.3 Atomic Radius
Atomic radius – estimated as ½ the distance
between the nuclei of 2 like atoms of the
same element when the atoms are joined.
•
As atomic radius increases, the element increases
in size.
+++
+
++++
Decreasing Atomic Size
Across a Period
+++
+
Li
Be
B
1s22s1 1s22s2 1s22s22p1
Li
Be
B
Atomic Radius
6.3 Atomic Size
Group Trend
Increases down a group
Caused by an increase in Shells and Shielding
Period Trend
Decreases across a period
As nuclear strength increases the nucleus pulls
the outer electrons closer
Atomic Radius of Atoms
Na K Rb Cs Cl S P Si Al Br Se As Ge Ga I Te Sb Sn InTl Pb Bi Mg
Ca
Sr
Ba
Be B C N O F
6.3 Ionization Energy
Ionization energy – energy needed to
remove an electron from an atom
Na
(g)
Na
+(g)
+ e
6.3 Ionization Energy
To remove an electron you have to overcome
the nucleus’ hold (nuclear charge) on the electron.
1st Ionization Energy – Energy needed to remove first
electron.
2nd Ionization Energy – Energy needed to remove a
second electron.
This is always higher than the 1st ionization energy.
When an electron is removed, the nucleus has a
stronger hold on the remaining electrons.
When you have a noble gas electron configuration
it becomes very difficult to remove an electron.
Multiple Ionization Energies
Al
57Al
+Al
2+Al3+
8 kJ /mol
e
-1817 kJ/
mol e
-2745 kJ/
mol e
-The second, third, and fourth ionization energies of aluminum are higher
than the first because the inner electrons are more tightly held by the nucleus.
1st Ionization
energy
2nd Ionization
energy
3rd Ionization
6.3 Ionization Energy
6.3 Ionization Energy
Group Trend
Decreases as you go down a group
1. Shells – the farther the outer electrons are from the nucleus, the weaker the pull.
2. Shielding – the inner level electrons block the nucleus’ ability to attract the valence electrons.
Period Trend
Increase as you go across a period
– Greater nuclear strength makes it harder to remove electrons
6.3 Ionization Energy
1st Ionization Energy
increases
1
st
Io
n
iz
a
tio
n
E
n
e
rg
y
d
e
cr
e
a
se
s
1A
2A 3A 5A 6A 7A
0
6.3 Ion Formation
Octet Rule – atoms will gain or lose electrons (sometimes even sharing like in molecules) to acquire a full set of 8 valence electrons.
Metals will lose electrons Nonmetals gain electrons
Two atoms are walking down the
street.
Hey, I think I just lost an
electron! Are you
sure!
Yeah! I’m
POSITIVE
!
6.3 Ion Formation
Cation – formed when electrons are
removed from a neutral atom
Nucleus has a stronger pull on the
remaining electrons decreasing the size
Usually involves a decrease in number
shells
The Cation is
smaller than
the neutral
atom!
152
Li
60
Li+
e
152
60
Li+
e
e
e e
Li
Lithium atom Lithium ion
+
152
Li
Lithium atom
Li
Energy
6.3 Ion Formation
Anion – formed when electrons are
added to a neutral atom
Nucleus has a weaker hold on the
increased number of valence electrons and
the ion size increases
The Anion is
larger than the
neutral atom!
6.3 Ion Size
Period Trend
Cations and Anions both decrease in size
across a period
Increased nuclear charge pulls in the valence
electrons
Group Trend
Cations and Anions both increase in size
down a group
Increased shells and shielding mean greater
ion size
Trends in Atomic and Ionic
Size
152 186 227 Li Na K 60 Li+ 95 Na+ 133 K+ e e e F -136 Cl -181 Br -195 F Cl Br 64 99 114 e e e Metals Nonmetals Group 1 Al 143 50 e e e Group 13 Group 17Cations are smaller than parent atoms Anions are larger than parent atoms
Al3+
6.3 Ionic Size
Anions decrease
Cations decrease
B
o
th
in
cr
e
a
se
6.3 Electronegativity
Electronegativity – is the ability of an atom
of an element to attract electrons when
the atom is in a compound.
O
H
H
Water
Molecule
Measure of how
strong the
oxygen nucleus
attracts the
hydrogen’s
electron
6.3 Electronegativity
6.3 Electronegativity
Period Trend
Increases across a period
Nuclear strength increases Greater hold on electrons
Group Trend
Decreases down a group
Shells and shielding mean the effective nuclear
strength decreases down the group
Weaker hold on the electrons
6.3 Electronegativity
Electronegativity increases
E
le
ct
ro
n
e
g
a
tiv
ity
d
e
cr
e
a
se
s
1A
2A 3A 5A 6A 7A
0
Summary of Periodic Trends
Ionic size (cations) Ionic size (anions) decreases decreases
Shielding is constant Atomic radius decreases Ionization energy increases Electronegativity increases Nuclear charge increases
N u c le ar c h ar g e *i n cr ea se s , b u t is m in im iz ed b y S /S S h ie ld in g i n c re as es A to m ic r ad iu s in cr ea se s Io n ic s iz e in c re as es Io n iz at io n e n er g y d ec re as es E le ct ro n eg at iv it y d ec re as
es 1A 2A 3A 4A 5A 6A 7A 0
