Download 6.1 Development of the Modern Periodic Table Objectives: 1

6.1 Development of the Modern Periodic Table
Objectives:
1. Describe the major advancements in development of the periodic table
2. Describe the organization of the elements on the periodic table
3. Classify elements by group, period, block, metallic characteristics, family name,
phase, series name
4. Characterize metals, nonmetals and metalloids.
History of the Periodic Table’s Development
th
 18 Century – 23 elements
th
 19 Century – 70 elements
 John Newlands – English Chemist (1837-1898)
 1864 – Newlands arranged elements by atomic mass and noticed that their
properties repeated every 8 elements (Law of Octaves)
 Meyer, Mendeleev, & Moseley
 Lothar Meyer and Dmitri Mendeleev related chemical properties to the
atomic masses of elements.
 Mendeleev organized his table of elements by atomic mass and arranged
elements with similar properties so that they lined up in columns.
 This was the first periodic table.
Mendeleev’s periodic table left blanks where he predicted elements
were yet to be discovered (Sc, Ga, Ge)
 He also predicted the properties of those elements.
Henry Moseley – English Chemist (1887-1915)
 Working in Rutherford’s lab he correlates the positive charges from
the nucleus with an integer resulting in the atomic number.
 Determines that the order of the elements should be based on this
fundamental property rather than the property of mass



Periodic Law
 When elements are arranged in order based on increasing atomic
number, their physical and chemical properties show a periodic
repetition.
The Modern Periodic Table
 Reading the Periodic Table
 Different periodic tables include different information. Typically, you
find:
 Atomic number (an integer)
 Atomic symbol (case matters)
 Atomic mass (to a certain number of sig figs)

Periodic Nomenclature
 Columns are called groups or families
o 18 columns in standard periodic table
o Traditionally numbered I-VIII, followed by A or B
o Modern number system is 1-18
 Rows area called periods
o 7 periods in standard periodic table
o


The two rows at the bottom of the periodic table are called the
“inner transition elements or metals”.
o The inner transition metals include the lanthanide series (4f) and
actinide series (5f).
Main group or representative elements
o Elements whose number ends with an “A”
o The s and p blocks.

Transition elements
o Elements whose number ends with “B”
o The d block

Inner transition elements
o The elements at the bottom of the periodic table
o The f block
Metals, Non-metals and Metalloids
 The Stair Step
o To the right of the P.T. there is a heavy stair step line that gives us
a good idea if an element is metallic, non-metallic or semimetallic.
Metals
o Elements to the left of the stair step are metals
o Largest group of elements
o Typically solid at room temperature
o Have luster
o Good conductors of heat and electricity
o Malleable and ductile
 Non-Metals
o Elements to the right of the stair step are non-metals
o Most are gas at room temperature
o No metallic luster
o Poor conductors of heat & electricity
o Neither malleable nor ductile
 Metalloids/Semi-metals
o Elements along the stair step line are metalloids
o Share properties of metals and non-metals.
o Have metallic properties
o Include semi-conductors
Special Group Names
 Alkali metals (group IA)
 Alkaline-earth metals (group IIA)
 Halogens (group VIIA)
 Noble Gases (group VIIIA)
 Other families are named based on the element that appears at the top
of the column (e.g. Group 14 (4A) = carbon family)
 Hydrogen is often placed by itself, in Group 1A, Group 17 or both 1
& 17 because of its unique properties.


6.2 Classification of the Elements
Objectives:
1. Determine the valence electrons for elements
2. Describe general properties of elements in each block
Organizing the Elements by Electron Configuration
 Valence Electrons
 Electrons in the outermost energy level of an atom
 Determine the chemical properties of the element
 Elements are grouped on the periodic table based on having the same
number of valence electrons
Li
Na
K
Rb
Cs

1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p63d104s24p65s1
1s22s22p63s23p63d104s24p64d105s25p66s1
There may be no more than 8 valence electrons for any one atom
 Valence Electrons & Period
 The period number of an element indicates the energy level of the valence
electrons for that element
 Valence Electrons and Group Number
 The Roman numerals on the periodic table for the main-group elements
show the number of valence electrons available for bonding
 The number of valence electrons for the transition elements is technically
2, but in reality the electrons from the lower s sublevel mingle with the d
sublevel electrons to create a variety of bonding possibilities.
The s-, p-, d-, & f-Block Elements
 s-block elements
 Groups IA & IIA (plus Helium)
 Will have either 1 or 2 valence electrons
 p-block elements
 Groups IIIA – XIIIA
 Will have 3 – 8 valence electrons
 Includes the noble gases
o Incredibly stable because the have a maximum number of valence
electrons
o Rarely react chemically (larger noble gases may be forced to react)
 d-block elements
 Transition metals
 Will usually have 2 electrons in outermost level (s-sublevel) and a
partially filled set of d orbitals
 Have similar chemical & physical properties
 f-block elements
 Inner transition metals
 Will usually have 2 electrons in outermost level (s-sublevel) and a partially
filled set of f orbitals
 Have similar chemical & physical properties
 d- and f- sublevels
 Have many different possible distributions of electrons as these large
sublevels shift electrons among similar energy sublevels and many
orbitals
6.3 Periodic Trends
Objectives:
1. Describe the nature of periodic trends
2. Define atomic radius, ion radius, ionization energy, electronegativity and electron
3. Relate ionization energy to electron configurations (box orbital really)
4. Compare elements by atomic radius, ion radius, ionization energy,
electronegativity and electron
 Certain properties of elements change in a predictable way as you move through
the periodic table. These predictable changes are called periodic trends.
 Periodic trends relate to the attraction of the nucleus to electrons.
Atomic Radius
 Atomic radius
 conceptually – distance from the center of the nucleus to the outermost
electron
 pragmatically – distance between nuclei of adjacent atoms
 Rules for estimating relative atomic radii
 Atoms get larger down a group.
 The principal quantum number
increases which means that higher
and higher energy levels are being
used to store electrons.
 Atoms get smaller moving left to right
across the periodic table.
 Moving across a period, additional
electrons occupy the same principal energy level (they don’t get
farther from the center), but because additional protons are present in
the nucleus, there is greater pull from the positively charged nucleus
on the negatively charged electrons
Ionic Radius
 Ionic radius – size of an ion of a particular element
 Rules for estimating relative ionic sizes
 When an atom loses electrons, it
becomes smaller.
 Electrons are shed from the
outermost energy level making it
smaller.
 Since the size of the negative charge
decreases, the positively charged
nucleus can pull harder on the
electrons and get them closer to the
nucleus.
 When an atom gains electrons, it
becomes larger.
 Additional electrons are added at the
fringe of the atom.
 Repulsive forces between additional electrons make them spread out.
 Attraction from the positively charged nucleus is spread thinner to
more electrons and is not as strong.
Ionization Energy
 Ionization energy – energy required to strip off electrons (create an ion)
 Basically ionization energy is how well an atom holds its electrons
 This is directly related to atomic radius. The greater the atomic radius, the
less attraction and therefore the less energy required to strip electrons off.
 First ionization energy – energy required to strip off highest energy electron
 Ionization energy is measured in joules/atom or joules/mole

 Rules for estimating relative first ionization
energies
 Ionization energy increases as you move
left to right across the periodic table
 Ionization energy increases as you move
up the periodic table
 Successive ionization energies
 Successive ionization energies are the
energies required to strip off successive electrons (i.e., energy to strip 2 nd,
3rd, etc.)
 These successive ionization energies have patterns that tell us how stable
certain electron configurations are.
 Noble gas configurations are very stable. When there are as many
electrons as in a noble gas, stripping one off will be very difficult. (High
ionization energy.)
Electronegativity
 Electronegativity – ability to attract electrons in a chemical bond.
 Atoms with high electronegativity will tend to “pull” on electrons harder
in a chemical bond.
 When electrons are pulled to one side, that side develops a more negative
charge. The other side develops a more positive charge.
Summary of Periodic Trends
Trend/Property Across a Period (LR)
Atomic Number
Increases
Atomic Mass
Increases
Atomic Radius
Decreases
Ionic Radius
Depends
Ionization Energy
Increases
Electronegativity
Increases
Down a Group
Increases
Increases
Increases
Increases
Decreases
Decreases