Periodic Properties of the Elements
Development of the Periodic Table
Dmitri Mendeleev and Lothar Meyer independently concluded on the systematic grouping of elements. Mendeleev is primarily credited because he utilized chemical properties for organization and successfully predicted missing elements and their characteristics, including germanium.
Mendeleev and the Periodic Table
Mendeleev's periodic table was originally arranged according to atomic masses, the most fundamental property known during that era. This structure was later refined following the discovery of the nuclear atom by Ernest Rutherford. Henry Moseley contributed further by defining atomic number as a primary basis for classifying elements within the periodic properties.
Periodicity
Periodicity refers to the repetitive patterns observable in element properties as a function of atomic number. This chapter will cover various properties including:
Sizes of atoms and ions
Ionization energy
Electron affinity
Trends in group chemical properties
The concept of effective nuclear charge (ZeffZeff) is foundational for many trends discussed in this chapter.
Effective Nuclear Charge
The effective nuclear charge, denoted as ZeffZeff, is defined as: Zeff=Z−SZeff=Z−S
Where:
Z is the atomic number (total number of protons).
S is the screening constant, typically approximated to the number of inner electrons.
ZeffZeff serves as a periodic property:
It generally increases across a period.
It decreases down a group.
Trends in Effective Nuclear Charge
Graphically depicted, the effective nuclear charge for core and valence electrons changes based on periodic trends. Notably, as we move across a period from B (Boron) to Ne (Neon), Zeff increases, signifying stronger attraction of valence electrons by the nucleus as more protons are added.
Atomic Size
The size of an atom is measured by two primary radii:
Nonbonding atomic radius (van der Waals radius): This is defined as half the distance between two nuclei during a collision of unbonded atoms.
Bonding atomic radius: This radius is half the internuclear distance when two atoms bond.
Bonding atomic radius exhibits trends:
It decreases from left to right across a period due to increasing Zeff.
It increases down a group due to increased n values (principal quantum number).
Example of Bond Lengths
For given bonds:
C—S bond length is estimated using radii:
extC—extS=0.76 A˚ (C radius)+1.05 A˚ (S radius)=1.81 A˚extC—extS=0.76 A˚ (C radius)+1.05 A˚ (S radius)=1.81 A˚
Experimental values confirm deviations typical with hydrogen bonds owing to their unique characteristics.
Calculating Atomic Radius Trends
Considering elements B, C, Al, Si, we assess relative atomic sizes:
C and B are on the same period, with C being smaller than B as radii decrease across a period.
Continuing downward, Si is smaller than Al (group trend and period trend considerations apply).
In summary, we expect the order of radii to be:
extC<?<?<extAlextC<?<?<extAl
Ion Size
Ion sizes are determined by interatomic distances in ionic compounds and are influenced by:
Nuclear charge
Number of electrons
Electron orbitals occupied
Cations (positively charged) tend to be smaller than their parent atoms due to the removal of electrons reducing electron-electron repulsion. Conversely, anions (negatively charged) are larger due to added electrons increasing repulsion.
Isoelectronic Series
In an isoelectronic series, ions share the same number of electrons. The size of ions diminishes with increasing nuclear charge:
An example of an isoelectronic series is:
O2−(1.26 A˚)>F−(1.19 A˚)>Na+(1.16 A˚)>Mg2+(0.86 A˚)>Al3+(0.68 A˚)O2−(1.26 A˚)>F−(1.19 A˚)>Na+(1.16 A˚)>Mg2+(0.86 A˚)>Al3+(0.68 A˚)
Ionization Energy (I)
Ionization energy is described as the minimum energy required to remove an electron from a gaseous atom or ion. The first ionization energy (I1I1) pertains to the initial electron removal while the second (I2I2) pertains to the subsequent removal:
Typically, ionization energy increases for every sequential electron removed.
The challenges escalate significantly once all valence electrons have been eliminated.
Trends in Ionization Energy
General trends indicate:
I1I1 tends to increase across a period.
I1I1 usually decreases down a group.
S- and p-block elements show a broader range compared to d-block, which rises slowly, and f-block elements that exhibit minimal variation.
Factors influencing ionization energy entail effective nuclear charge as well as the average distance of the electron from the nucleus.
Irregularities in Ionization Energy
Irregularities may occur, particularly when adding a valence electron to a new sublevel or forming pairs, thus causing electron repulsions leading to lower energy.
Electron Configurations of Ions
Electron configurations in cations involve the loss of electrons from the highest energy levels. For example:
The lithium ion Li+Li+ is represented as 1s21s2 (removal of the 2s electron).
The cobalt ion Co3+Co3+ is illustrated as [Ar]3d6[Ar]3d6 after losing 2 electrons from 4s and a 3d.
Conversely, anions typically gain electrons to complete their octet, such as F−F− gaining one electron to become [Ne]3s23p6[Ne]3s23p6.
Electron Affinity
Electron affinity quantifies the energy change accompanying an electron's addition to a gaseous atom, generally leading to an exothermic reaction. As elements are assessed within groups or across periods, periodic trends arise:
Little variation occurs down a group
Electron affinity typically augments across a period with notable exceptions identified along Groups 2A, 5A, and 8A, which are characterized by filled or half-filled sublevels, rendering them unstable upon an additional electron.
Metals, Nonmetals, and Metalloids
Metal Properties
Metals generally form cations and exhibit:
Shiny luster
High electrical and thermal conductivity
Malleability and ductility
Solid state at room temperature (exception: mercury)
Nonmetal Characteristics
Nonmetals form anions and comprise various states, showcasing:
Dull and brittle solids when in solid state
Low conductivity, large negative electronegativity
Nonmetal Chemistry
Compounds formed solely from nonmetals tend to be molecular, with nonmetal oxides frequently presenting acidic properties.
Group Trends
Each group of the periodic table—such as alkali metals (Group 1A), alkaline earth metals (Group 2A), oxygen group (Group 6A), halogens (Group 7A), and noble gases (Group 8A)—exhibits distinctive and similar chemical properties.
Alkali Metals
Alkali metals, characterized by low density and low ionization energy, predominantly react with water, yielding exothermic reactions.
Notably, lithium forms an oxide 4Li+O2→2Li2O4Li+O2→2Li2O, while sodium forms peroxides and larger alkali metals form superoxides.
Flame Tests
Alkali metals can be qualitatively identified by their distinct flame colors.
Alkaline Earth Metals
Alkaline earth metals present higher densities and thermal/melting points than alkali metals, although beryllium does not react with water, magnesium does so with steam, and other alkaline earth metals react readily with water. Reactivity increases down the group.
Group Properties
Group 6A displays a reduction in metallic character down the group while Group 7A (halogens) features highly negative electron affinities.
Noble gases (Group 8A) possess very high ionization energies and typically exist in monatomic gaseous forms, much more inert than other elements due to their strong electron configurations.
Formulas and Equations from the Notes
Effective Nuclear Charge: Z_{eff}=Z-S
Where Z is the atomic number and S is the screening constant.
Estimated C—S Bond Length: \text{C—S}=0.76 \text{ \AA (C radius)}+1.05 \text{ \AA (S radius)}=1.81 \text{ \AA}
Atomic Radii Order (Example): \text{C}<\text{Si}<\text{B}<\text{Al}
Isoelectronic Series Example: \text{O}^{2-}(1.26 \text{ \AA})>\text{F}^{-}(1.19 \text{ \AA})>\text{Na}^{+}(1.16 \text{ \AA})>\text{Mg}^{2+}(0.86 \text{ \AA})>\text{Al}^{3+}(0.68 \text{ \AA})
Reaction of Lithium with Oxygen: 4Li+O2→2Li2O