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Lecture Presentation - Chapter 7: Periodic Properties of the Elements

• The periodic table organizes elements based on recurring patterns in their properties.
• Dmitri Mendeleev and Lothar Meyer independently developed similar periodic tables, based on atomic mass.
• Mendeleev is credited for using chemical properties to both organize the table and predict missing elements accurately. He predicted the properties of germanium based on the properties of neighboring elements, which were later validated when discovered.
• Mendeleev's table was based on atomic mass, the most fundamental property known at the time.
• Henry Moseley later determined the atomic number is the basis for periodic properties.
• The concept of atomic number was developed experimentally and determined that the number of protons defined the periodic property.

Periodicity

• Periodicity is the repeating pattern of properties of elements based on atomic number.
• Key properties discussed include atomic/ionic size, ionization energy, electron affinity, and group trends.
• The effective nuclear charge, zeff, is a major factor affecting periodic trends.

Effective Nuclear Charge

• Atoms' properties depend on attractions between valence electrons and the nucleus.
• Electrons are attracted to the nucleus and repelled by other electrons.
• Effective nuclear charge (Z_eff) is calculated as atomic number (Z) minus the screening constant (S).
• Z_eff increases across a period and decreases down a group.

Sizes of Atoms and Ions

• The bonding atomic radius is half the internuclear distance between bonded atoms.
• Atomic radius tends to decrease across a period and increase down a group. This is due to increasing screening effects and increasing principal quantum numbers.
• Cations are smaller than their parent atoms—valence electrons are lost which causes a reduction in electronic shielding.
• Anions are larger than their parent atoms—valence electrons are gained, increasing electron repulsions.
• Isoelectronic series have ions with the same number of electrons, and ionic size decreases with increasing nuclear charge.

Ionization Energy

• Ionization energy is the energy required to remove an electron from a gaseous atom/ion.
• First ionization energy is needed to remove the first electron, second ionization energy for the second, and so on.
• Ionization energy generally increases across a period and decreases down a group. This is affected by the increasing effective nuclear charge and decreasing distance from the nucleus, affecting the amount of energy needed.
• Irregularities occur due to sublevel changes and electron pairing in orbitals.

Electron Affinity

• Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom.
• Electron affinity is typically exothermic (negative) for most elements.
• Electron affinity generally increases across a period and fluctuates depending on electron configurations and the stability of the resulting anion.

Metal, Nonmetal, and Metalloids

• Metals tend to form cations, are solid at room temperature (except mercury), shiny, and are good conductors of heat and electricity. Most metal oxides are basic.
• Nonmetals tend to form anions, exist in various states (solid, liquid, gas), tend to be dull and brittle, and are poor conductors. Most nonmetal oxides are acidic.
• Metalloids have properties intermediate between metals and nonmetals.

Group Trends

• Elements in a group tend to have similar properties due to valence electron configurations.
• There are specific group trends (such as alkali metals, alkaline earth metals, oxygen group, halogens, and noble gases) and their characteristics.

Alkali Metals

• Alkali metals are soft, metallic solids.
• Found primarily in compounds in nature—not in elemental form.
• Exhibit typical metallic properties (luster and high conductivity).
• Have low densities and melting points.
• Low ionization energies—so they readily form cations and react with water to create an exothermic reaction.
• Their reactions with water also vary, with lithium reacting to create an oxide, sodium a peroxide, and other alkaline earth metals superoxides.

Alkaline Earth Metals

• Have higher densities and higher melting points than alkali metals.
• Ionization energies are lower than halogens and nonmetals, but higher than those of alkali metals.
• Reactivity of alkaline earth metals increases down the group.
• Beryllium does not readily react with water, increasing in reactivity with the other alkaline earth metals moving down the group.

Other Group Trends

• Group 6A elements exhibit increasing metallic character moving down the group.
• Group 7A halogens are typical nonmetals with high negative electron affinities—so they exist primarily as anions, and readily react with metals to produce metal halides.
• Group 8A noble gases have very high ionization energies and positive electron affinities. They are relatively unreactive and exist as monatomic gases.