Chemistry I - Periodic Table

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Dmitri Mendeleev


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Dmitri Mendeleev

Russian Chemist, used atomic mass and similar chemical properties to arrange elements into groups (families). Constructed first table in 1869. Left blank spaces and predicted the existence of three elements whose properties ended up being very similar.

Henry Mosely

English scientist, modified the periodic table in 1913. He found that elements relate better when arranged by atomic number (how they are arranged today).

Periodic Law

When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.

Periodic Table

Used to understand and predict properties of elements. Has 18 groups and 7 periods.


How many elements are on the modern day periodic table?


18 of them. Elements have similar physical and chemical properties


7 of them. The properties of elements change as you move across them from element to element. Pattern of properties repeats as you move from one to the next.


Good conductors of heat and electricity, high luster when clean, ductile, malleable, and solid at room temp (except Mercury).


Nonmetallic, non-lustrous, poor conductors of heat and electricity. Some are gas (oxygen, chlorine, fluorine, nitrogen), some are brittle solids (sulfur, carbon, selenium, iodine), and bromine is a liquid at a room temp.


Properties are intermediate between those of metals and nonmetals. Form a step line. They are semiconductors and all solids. Elements: Boron, silicon, germanium, arsenic, antimony, and tellurium (sometimes astatine and polonium).

Alkali Metal

Other name for group 1 elements

Alkaline Earth Metals

Other name for group 2 elements


Other name for group 17 elements

Noble Gases

Other name for group 18 elements

Transition metals

Located in groups 3-12 and period 4-7

Inner Transition Metals

Pulled out rows of the periodic table. Include Lanthanides and Actinides. Also called Rare Earth Metals.

Atomic Radius

1/2 the distance between the nuclei of two like atoms in a diatomic molecule

Trends in Atomic Size

Increases as you move down a group because electrons are added to higher energy levels. Decreases as you move left to right across a period because of an increase of positive charges between the nucleus and electrons (pulls in, makes atom smaller). Francium largest, Helium smallest.


Atom or group of bonded atoms that has a charge.

Ionization Energy

Energy required to remove an electron from atom

Trends in Ionization Energy

Decreases as you move down a group because the size of atoms increase and electrons are farther away from the nucleus making it easier to remove them. Increases from left to right across a period because there is a greater attraction of nucleus for the electron. Helium largest, Francium smallest.

Electron Affinity

Energy charge that occurs when an electron is gained by a neutral atom.

Trends in Electron Affinity

Decreases as you move down a group. Increases as you move left to right across a period because the closer to a full sublevel, the higher it is. Fluorine largest, Francium smallest. Groups 2 & 18 have zero.

Ionic radius

The size of a cation or anion. Cation is smaller and a anion is larger.


Positively charged ion, lost electrons, size decreases (half the radius of neutral)


Negatively charged ion, lost electrons, size decreases (half the radius)

Trends in Ionic Size

Increase as you go down a group. Decreases from left to right until Group 15. Group 15 elements increase (between metals and nonmetals) and then decrease until Group 18.


The measure of the ability of an atom in a chemical compound to attract electrons from another atom.

Trends in Electronegativity

Decreases as you move down a group. Increases as you move left to right across a period. Noble Gases are not included because they are not attractive. Fluorine largest because it attracts electrons, Francium smallest because it easily loses an electron.