Periodic Properties of the Elements - Summary

Development of the Periodic Table

• Dmitri Mendeleev and Lothar Meyer independently grouped elements.
• Mendeleev is credited for using chemical properties, predicting missing elements like germanium.

Atomic Number

• Mendeleev's table was based on atomic masses.
• Henry Moseley developed atomic number concept experimentally.
• Number of protons became the basis for periodic properties.

Periodicity

• Periodicity: Repetitive pattern of properties based on atomic number.
• Properties discussed:
  • Sizes of atoms and ions
  • Ionization energy
  • Electron affinity
  • Group chemical property trends

Effective Nuclear Charge

• Effective nuclear charge ($Z_{eff}$) influences many properties.
• $Z_{eff} = Z - S$, where Z is atomic number, S is screening constant.
• Trends:
  • Increases across a period.
  • Increases slightly down a group.

Size of Atoms

• Nonbonding atomic radius (van der Waals radius): Half the shortest distance between two nuclei during a collision.
• Bonding atomic radius (covalent radius): Half the distance between nuclei in a bond.
• Trends in bonding atomic radius:
  • Decreases from left to right across a period (increasing $Z_{eff}$).
  • Increases from top to bottom of a group (increasing n).

Sizes of Ions

• Determined by interatomic distances in ionic compounds.
• Depends on nuclear charge, number of electrons, and orbitals.
• Cations: Smaller than parent atoms (electron removal reduces repulsions).
• Anions: Larger than parent atoms (electron addition increases repulsions).

Size of Ions—Isoelectronic Series

• Isoelectronic series: Ions with the same number of electrons.
• Ionic size decreases with increasing nuclear charge.

Ionization Energy (I)

• Ionization energy: Minimum energy to remove an electron from a gaseous atom or ion.
  • First ionization energy: Energy to remove the first electron.
  • Second ionization energy: Energy to remove the second electron.
• Higher ionization energy means it is more difficult to remove an electron.

Ionization Energy Trends

• Successive ionization energies increase.
• Significant increase after all valence electrons are removed.
• Trends in first ionization energy ($I_1$):
  • Increases across a period.
  • Decreases down a group.
  • s- and p-block elements show larger range of values.

Factors Influencing Ionization Energy

• Smaller atoms have higher I values.
• I values depend on effective nuclear charge and distance from the nucleus.

Irregularities in Ionization Energy Trend

• Trend not followed when:
  • Added electron enters a new, higher energy sublevel.
  • First electron pairs in one orbital (electron repulsions lower energy).

Electron Configurations of Ions

• Cations: Electrons lost from highest energy level (n value).
  • Example: $Li^+$ is $1s^2$.
  • Example: $Fe^{2+}$ is $1s^22s^22p^63s^23p^63d^6$.
• Anions: Configurations filled to $ns^2np^6$.
  • Example: $F^−$ is $1s^22s^22p^6$.

Electron Affinity

• Electron affinity: Energy change upon adding an electron to a gaseous atom: $Cl + e^- \rightarrow Cl^−$.
• Typically exothermic (negative value).

Electron Affinity Trend

• Little change in a group.
• Generally increases across a period with exceptions:
  • Group 2A: s sublevel is full.
  • Group 5A: p sublevel is half-full.
  • Group 8A: p sublevel is full.

Metals, Nonmetals, and Metalloids

• Metals form cations.
• Nonmetals form anions.

Metals

• Shiny luster, conduct heat and electricity, malleable and ductile.
• Solids at room temperature (except mercury).
• Low ionization energies/form cations easily.
• Metal oxides are basic and react with acids
• Metal and nonmetal compounds tend to be ionic

Nonmetals

• Solid, liquid, or gas; solids are dull, brittle, poor conductors.
• Large negative electron affinity, form anions readily.
• Nonmetal oxides are acidic.
• Substances containing only nonmetals are molecular compounds.

Metalloids

• Have properties of both metals and nonmetals.
• Electrical semiconductors.

Group Trends

• Elements in a group have similar properties.
• Groups compared:
  • Group 1A: Alkali metals
  • Group 2A: Alkaline earth metals
  • Group 6A: Oxygen group
  • Group 7A: Halogens
  • Group 8A: Noble gases
  • Hydrogen is a nonmetal

Alkali Metals

• Soft, metallic solids; found in compounds.
• Metallic properties (luster, conductivity).
• Low densities, melting points, and ionization energies.
• React exothermically with water.

Alkali Metal Chemistry

• Lithium reacts with oxygen to make an oxide: $4 Li + O<em>2 \rightarrow 2 Li</em>2O$

• Sodium reacts with oxygen to form a peroxide: $2 Na + O<em>2 \rightarrow Na</em>2O_2$

• K, Rb, and Cs also form superoxides: $M + O<em>2 \rightarrow MO</em>2$

Alkaline Earth Metals

• Higher densities and melting points than alkali metals.
• Low ionization energies.
• Readily form +2 cations.
• Reactivity with water increases down the group.

Group 6A

• Increasing metallic character down the group.
• Oxygen, sulfur, and selenium are nonmetals; tellurium is a metalloid; polonium is a metal.
• Oxygen can exist as dioxygen ($O<em>2$) or ozone ($O</em>3$).

Group 7A (Halogens)

• Typical nonmetals with highly negative electron affinities.
• Exist as anions, react directly with metals to form metal halides.

Group 8A (Noble Gases)

• Large ionization energies, positive electron affinities.
• Relatively unreactive, found as monatomic gases.

Hydrogen

• Can form H+ and H- ions