Periodic Table history, blocks, groups, periods, periodic table trends, metal, non metals and metalloids

1. STAGE 3
PERIODIC
TABLE.
M. C. Laura Verástegui

2. HISTORY OF THE PERIODIC
TABLE
 Law of Triads: In 1817 Johann Döbereiner proposed that
nature contained triads of elements the middle element
properties that were an average of the other two members
when ordered by the atomic weight.
 Law of Octaves: John Newlands, he classified elements
based on similar physical properties, noting that many
of similar elements existed which differed by some multiple
of eight in atomic weight.

3.  Lothar Meyer: In 1868 he developed his own version of the
periodic table, unfortunately for Meyer, Mendeleev's table
became available to the scientific community via publication
two years before his work.
 Dmitri Mendeleev: He is considered the father of the periodic
table, he developed a periodic table of the known elements
ordered according to their atomic mass and letting spaces for
the elements not discovered yet, predicting their properties
atomic masses.

4. MODERN PERIODIC TABLE
Henry Moseley :In 1913, published the results of his
measurements of the positive charge of the nucleus
the atoms, he concluded that the elements coincided
with the ordering of the elements by ascending
atomic number.
It became apparent that atomic weight was not the
significant player in the periodic law as Mendeleev,
Meyers and others had proposed, but rather, the
properties of the elements varied periodically with
atomic number: Periodic Law

5. PERIODIC TABLE
ORGANIZATION
• The periodic table provide a lot of
information and is organized in a special
arrangement, we may found:
• Metal, metalloid and non-metals
• Periods
• Groups or families
• Blocks

6. PERIOD: is a horizontal row in the periodic table.
Indicate the outermost energy level of the
elements.
BLOCK: Named according to the subshell in
which the "last" electron resides. The s-block
comprises the first two groups (alkali metals and
alkaline earth metals) as well as hydrogen and
 helium. The p-block comprises the last six
groups (groups 13 through 18) and contains,
among others, all of the semimetals. The d-
block comprises groups 3 through 12 and
contains all of the transition metals. The f-block,
usually offset below the rest of the periodic
table, comprises the rare earth metals.

8. GROUPS
• Or family is a vertical column in the periodic
table. Groups are considered the most
important method of classifying the elements.
In some groups, the elements have very
similar properties and exhibit a clear trend in
properties down the group. These groups
tend to be given trivial (unsystematic) names,
e.g., the alkali metals, alkaline earth metals,
halogens,, chalcogens, and noble gases.

10. • Representative elements are those label as A
groups: IA – VIIIA
–IA: Alkali metals
–IIA: Alkali Earth metals
–IIIA:Boron family
–IVA: Carbon family
–VA: Nitrogen family
–VIA: Oxygen family or Chalcogens.
–VIIA: Halogens
–VIIIA: Noble gases.
• Group B: All are called transition metals
• Lanthanide and actinide series: Rare Earth
metals.

11. Metalloids or Semimetals: Have similar properties to both
metals and non-metals. These elements are located along the
zigzag that divide the metals from non-metals: B, Si, Ge, As,
Sb, Te, Po, At.

13. VALENCE ELECTRONS
• Are the outermost electrons of an atom, which are
important in determining how the atom reacts
chemically with other atoms. They do show a
repeating or periodic pattern. The valence electrons
increase in number as you go across a period.
• Valence electrons are those in the electronic shell of
highest principal quantum number. For example the
electronic configuration of phosphorus (P) is 1s2 2s2
2p6 3s2 3p3 so that are 5 valence electrons (3s2 3p3),

14. VALENCE ELECTRONS

15. OXIDATION NUMBER
• Called oxidation state, The apparent charge assigned to an
atom within a compound.

16. PERIODIC TRENDS
• Atomic radii: is a measure of the size of its atoms, usually
the distance from the nucleus to the boundary of the
surrounding cloud of electrons.

17. • Ionization energy: The energy required to
remove an electron from an atom or
molecule in gas state.
• Electronic affinity: The energy absorbed or
released when a gaseous atom captures
electron.
• Electronegativity: a chemical property that
describes the ability of an atom to attract
electrons from a covalent bond.

18. ATOMIC RADII INCREASES
IONIZATION E, ELECTRONIC AFF. AND ELECTRONEGATIVITY
INCREASES