Why is the periodic table shaped like it is and how are the elements arranged?. The Periodic Table - History. Two scientists, D i mitri Mendeleev (Russia) and Lothar Meyer (Germany) properties of elements did not change smoothly with increasing atomic mas.
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Why is the periodic table shaped like it is and how are the elements arranged?
Two scientists, Dimitri Mendeleev (Russia) and Lothar Meyer (Germany)
properties of elements did not change smoothly with increasing atomic mas.
Instead the properties of the elements repeated periodically.
Periodic Law: the properties of the elements repeat periodically as the elements are arranged in order of increasing atomic number (# of protons)
This periodic law is used to form the Periodic Table
Vertically into Groups
Horizontally Into Periods
Each atom has the same number of electrons in it’s outermost shell.
for group A elements
Charge on stable monatomic ion
3A 4A 5A 6A 7A
Period number gives shell (n)
Elements are arranged according to atomic #
and e- configuration.
Li: 3 e-’s 1s22s1
Na: 11 e-’s 1s2 2s2 2p63s1
K: 19 e-’s 1s2 2s2 2p6 3s2 3p64s1
Paramagnetic or diamagnetic?
Valence orbitals: outer shell orbitals
beyond the closest noble-gas configuration
Valence electrons: “the ones that can react” (located in the valence orbitals).
The other e-’s are called core electrons and don’t react.
Elements in a vertical row have the same number of valence electrons.
you would see…
The elements in the same horizontal row are called a period
Hydrogen and Helium are in period 1
Lithium through Neon are in period 2
Fe (Iron) Atom
s1 s2p1 p2 p3 p4 p5 p6
3d1 - d10
f1 - f14
Condensed Ground-State Electron Configurations in the First Three Periods
Los Alamos National Laboratory's Chemistry Division
Periodic Table of the Elements
typical charge on ion in binary compound
Discovery dates of elements
Physical States at Room Conditions
Modified from http://www.cem.msu.edu/~djm/cem384/ptable.html
Natural vs. Man-Made Elements
not found on Earth
MolecularNature of Free Elements
All other free elements are atomic in nature.
All Isotopes Radioactive
no stable forms
(binary hydrogen compounds)
salts with H-
molecules with +1
chains with H-to-H bonds
3A 4A 5A 6A 7A
mostly gases, some liquids
What would you expect from Francium?!?!
Potassium (K), in Water (H2O)
Soft, silvery colored metals
Alkali Metal Family
Alkaline Earth Metals
Many are found in rocks in the earth’s crust
-Group VIII A
- Generally unreactive
Chlorine (Cl) Bromine (Br) and Iodine (I)
Most are Poisonous
Chlorine Gas was used as a chemical weapon during World War I.
It was used by the Nazis in World War II.
Most are good
Conductors of heat and electricity
Ductile and malleable
(easily bent/hammered into wires or sheets)
Metals very close to the “staircase” line
They have properties of metals and non-metals.
What are semiconductors used in?
.Do not conduct heat and electricity
Modern Periodic Table
Main Group (Representative Group) - Groups IA - VIIIA
Transition Metals - Groups IB – VIIIB
Rare Earth Elements -
Lanthanides (Ce - Lu) and
Actinides (Th - Lr)
about 75% of all the elements
lustrous, malleable, ductile, conduct heat and electricity
dull, brittle, insulators
also know as semi-metals
some properties of both metals & nonmetals
Li Be B C N O F
Ionization Energy (IE) - The amount of energy needed to remove an electron from an atom or ion. Each electron in any atom or ion has a specific ionization energy.
First Ionization Energy - The amount of energy needed to remove an electron from the outermost shell of a neutral (uncharged) atom.
The ionization energy is the energy that must be supplied to an atom in the gas phase in order to remove an electron. If the electron is the first one to be removed, one then refers to the first ionization energy for that element.
The first ionisation energy of sodium is 494 kJ.mol-1.
The ionisation energy tells us how easy it is to convert an element to a cation. The lower the first ionisation energy, the easier it is to convert an element such as sodium, Na, to its cation Na+.
The energy required to completely remove an e- from an atom in its gaseous state.
Mg(g) Mg1+ + e-
1st ionization energy
Mg1+(g) Mg2+ + e-
2nd ionization energy
Question: Which of the above ionizations would have the highest ionization energy and why?
electron being lost:
1st 2nd 3rd 4th 5th 6th 7th
The first ionisation energy of most elements are known to a good degree of accuracy. It is interesting to see how these values vary with atomic number Z:
The rare gases (He, Ne, Ar, Kr, Xe, Rn) appear at peak values of ionization energy, which reflect their chemical inertness, while the alkali metals (Li, Na, K, Rb, Cs) appear at minimum values of ionization energy, in keeping with their reactivity and ease of cation formation.
Ionization energy decreases as you go down a group
Ionization energy increases as you go from left to right in a period
What is meant by metallic character?
Periodicity of ionic radii:
The size of ions differ markedly from the size of their parent atoms. Take the case of sodium as an example:
The sodium atom (which has a single 3s electron), has an atomic radius of 186 pm. Upon taking up energy (the ionization energy), it loses this electron and is converted to the cation Na+, whose radius is 97 pm.
The radius of cations is always smaller than the radius of the atoms from which they are derived, as shown in the figure on the right, which applies to the Group I elements:
The main reason for this is that whenever metals are converted to their cations, they always do so by losing the electrons in their highest energy level.
Further, since the ion has less electrons than the atom from which they are derived, there is less mutual repulsion between these electrons, and the electron orbitals shrink to some extent.
What happens in the case of anions? Let's take the case of chlorine:
The chlorine atom has a covalent radius of 99 pm. It can gain an extra 3p electron (with release of energy, the electron affinity) in order to form the chloride anion Cl-, with ionic radius 181 pm.
Periodicity of electronegativities:
When a covalent bond is formed between two identical atoms, such as H-H or Cl-Cl, the pair of electrons which joins the atoms is evenly shared between the two atoms. However, when two different atoms are joined together by a covalent bond, the sharing of the electron pair is not even, and the pair of electrons is shifted towards one of the atoms:
The covalent bond is said to have been polarized, and one refers to such a bond as a polar covalent bond.
The ability to polarize a covalent bond differs from one atom to another, and is known as the electronegativity of the atom.
Electronegativities are measured on an arbitrary scale ranging from 0 (He) to 4.1 (F).
The rare gases (He, Ne, Ar, Kr, Xe, Rn) have zero electronegativities, i.e., they only form covalent bonds, and this only in exceptional cases, and have no tendency to attract electrons of that bond.
The halogens ( F, Cl, Br, I, At) appear at peak values of electronegativities. Within a period, they have the highest tendency to polarize covalent bonds. Note that within the group, the electronegativity decreases.
On the whole, electronegativities increase from left to right along a period (see Li to F) and decrease from top to bottom within a group (see F to At).
Periodicity of electron affinity:
The electron affinity is the energy that is released when an atom in the gas phase gains an electron and is thus converted to an anion, also in the gas phase:
The electron affinity of chlorine is 349 kJ.mol-1.
Electron affinities are difficult to measure and there is no reliable data available for most elements. However, the larger the atom, the lower its electron affinity, as shown with Group VII elements:
For reasons outside the scope of this discussion, the electron affinity of fluorine is an exception to this trend.
Common Oxidation states: note the vertical similarities.