Unit 3: Periodic Table Chapter 6
Objectives 21 Understand the historical background of the periodic table including such contributions of Newlands, Mendeleev, and Moseley 22 Use the periodic table to predict properties of certain elements 23 Identify the difference between periods and groups 24 Identify where main group, metals, metalloids, non-metals, alkali, alkaline-earth, halogen, noble gases, transition metals, lanthanides (rare earth metals) and actinides exist on the periodic table along with their characteristics and electron configurations 25 Define and apply the periodic law to the trends on the periodic table including atomic radius, ionization energy, and electron affinity/electronegativity
21 Historical Background on the Periodic Table The Periodic Table was original designed by John Newlands and was published in the 1860s. He also proposed an idea known as the Law of Octaves to help explain his setup for the table. More extensive work was done on the Periodic Table by Dmitri Mendeleev. He arranged the table by atomic mass. He left blanks where he felt elements should belong even though they were not discovered yet.
Mendeleev s Periodic Table
Historical Background Continued While Mendeleev s table was the first that resembled our current model, there were some flaws. None of the noble gases were present. There were some discrepancies Mendeleev could not explain in certain properties. It did not account for isotopes.
Historical Background Continued In 1914, Henry Moseley rearranged the Periodic Table by atomic number. By doing this, he eliminated the discrepancies that Mendeleev could not explain. Moseley s table is the model used today with the elements found after 1914 added to the table.
22 Characteristics of the Table The Periodic Table was carefully designed to provide as much information as possible. The table is ordered into: Periods: horizontal rows Groups: vertical columns Blocks: S, P, D, or F
23,24 Sections of the Table There are certain sections of the Periodic Table that have common properties. Metals Metalloids Non-metals Main Group Elements
Metals Shiny Form positive ions Ductile Malleable High melting and boiling points Good conductors Heat and energy Return
Non-Metals Return Make up the majority of the crust, atmosphere and living organisms Low melting and boiling points Form negative ions Low densities If solid, tend to be dull and brittle Poor conductors Both heat and energy
Metalloids Semi-conductors Contain some metallic properties Contain some nonmetallic properties Return
Main Group Elements Made up of the S and P blocks Consists of some of the most common elements. Return
Groups and Blocks Alkaline Earth Metals Halogens Noble Gases Transition Metals Alkali Metals Lathanide Series Actinide Series
Alkali Metals Highly reactive Rarely found in the elemental form Soft metals Low densities Make +1 ions Last electron is always a s 1 Return
Alkaline Earth Metals Always from +2 ions Last electron is always an s 2 High melting points Reactive but not as violent as the alkali metals Return
Halogens Highly reactive Form -1 ions All but astatine can form a diatomic molecule Common in acids Used as disinfectants and in pesticides Return
Noble Gases Return Odorless, colorless gases Outer (valence) energy level is full Very low reactivity Melting and boiling points are low and very close together Cryogenic refrigerants
Transition Metals Form the D-Block of the Periodic Table Magnetic Properties Return 1 or more unpaired electrons High melting and boiling points Can from +1, +2, +3 ions Generally solid
Lanthanide Series Make up the 4f block Typically used in lasers Sometimes referred to as the rare earth metals Though actually found in high concentrations in the crust Superconductors Batteries and magnets Return
Actinide Series Make up the 5f-block Most are man-made Thorium and uranium are the only two the occur naturally with any abundance. Radioactive Return
25 Periodic Law Certain properties follow periodic law. Periodic law refers to the increasing or decreasing of a trend as one progresses across a period or group on the Periodic Table. Three of the most common trends that are monitored and follow periodic law are atomic radius, ionization energy, and electron affinity.
Atomic Radius Atomic radius refers to the size of the electron cloud surrounding the nucleus. The atomic radius increases as each new energy level is added. While electrons are being added to an energy level, electron shielding allows for the affects the size of the atom.
Electron Shielding As electrons start filling energy levels, the nucleus holds them close. As more are added, the inner ring prevents the nucleus from pulling the outer ring too close (it shields the positive charge). The nucleus will pull the energy level slightly closer though as you progress across the table.
Atomic Radius Trend The atomic radius increases in the direction of the arrow.
Ionization Energy Ionization energy is the energy required to remove an electron from an atom. The larger the atom, the more difficult it is for the nucleus to hold onto its electrons. Smaller atoms can hold onto electrons much easier.
Ionization Energy Trend The ionization energy increases in the direction of the arrow.
Electronegativity Electronegativity refers to how well an atom attracts electrons. Smaller atoms have more nuclear charge to attract electrons. As that large atoms have a difficult time holding onto their electrons, they do not readily attract electrons.
Electronegativity Trend The electronegativity increases in the direction of the arrow.
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