Chemistry: Structure and Properties - Chapter 3 Notes
Periodic Law and The Periodic Table
Dmitri Mendeleev (1834-1907) - Russian chemistry professor.
Listed known elements (around 60 at the time) in order of increasing atomic mass.
Grouped elements with similar chemical properties, noticing a periodic recurrence of these properties.
His tabulation is considered the precursor to the modern periodic table.
He notably left gaps in his table for undiscovered elements and accurately predicted their properties (e.g., eka-silicon, later germanium, and eka-aluminum, later gallium), which were confirmed by subsequent discoveries.
Periodic Law - Originally observed as a repeating pattern of chemical and physical properties when elements are arranged by increasing atomic mass. The modern periodic law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.
These repeating properties are primarily due to the recurring patterns in the electron configurations of the elements, specifically the number of valence electrons.
These repeating properties lead to the formation of groups (columns) in the periodic table, where elements within a group exhibit similar chemical behavior.
Development of the Periodic Table - Mendeleev's arrangement (1869) accounted for all known elements but contained some inconsistencies where elements had to be placed out of order of increasing atomic mass to fit their chemical properties (e.g., tellurium and iodine).
Henry Moseley (1887-1915) - British physicist, approximately 40 years after Mendeleev's work.
Used X-ray spectroscopy to experimentally determine the atomic number of each element.
Discovered that elements should be arranged by increasing atomic number rather than atomic mass to resolve the inconsistencies in Mendeleev's table.
Moseley's work led to the modern statement of the Periodic Law and the refined structure of the modern periodic table.
The Modern Periodic Table
Elements are arranged in order of increasing atomic number.
Periods (Rows): There are seven horizontal rows (periods).
The period number corresponds to the principal energy level (n) of the valence electrons.
Elements in the same period have their valence electrons in the same principal energy shell.
Groups (Columns): There are 18 vertical columns (groups or families).
Elements within the same group generally have similar chemical properties because they have the same number of valence electrons and similar outer electron configurations.
Groups are often labeled with numbers (1-18) or historically with Roman numerals and letters (e.g., IA, IIA, IIIB).
Blocks: The periodic table can be divided into blocks corresponding to the type of orbital (s, p, d, f) being filled with electrons.
s-block: Groups 1 and 2 (alkali metals and alkaline earth metals).
p-block: Groups 13 to 18 (e.g., halogens, noble gases).
d-block: Groups 3 to 12 (transition metals).
f-block: Lanthanides and Actinides (inner transition metals), typically placed below the main table.
Electron Configuration Principles
Aufbau Principle (German: "building up"): States that electrons fill atomic orbitals of the lowest available energy levels before occupying higher energy levels. This systematic filling explains the structure of the periodic table.
Pauli Exclusion Principle: Proposed by Wolfgang Pauli, it states that no two electrons in the same atom can have identical values for all four of their quantum numbers (n, l, ml, and ms). This implies that an atomic orbital can hold a maximum of two electrons, and these two electrons must have opposite spins.
Hund's Rule (of maximum multiplicity): Proposed by Friedrich Hund, it states that for a set of orbitals of equal energy (degenerate orbitals) within a subshell (e.g., p, d, or f orbitals), electrons will first occupy separate orbitals with parallel spins before pairing up in any one orbital.
Key Concepts Influencing Periodic Trends
Effective Nuclear Charge (Z{eff}): This is the net positive charge experienced by an electron in a multi-electron atom. It is less than the actual nuclear charge (Z) because of the shielding effect of inner electrons. Z{eff} generally increases across a period and remains relatively constant or slightly increases down a group for valence electrons.
Shielding Effect: Inner-shell electrons "shield" (or screen) the outer-shell electrons from the full attractive force of the nucleus. The more inner electron shells an electron has beneath it, the greater the shielding, and thus the lower the Z_{eff} experienced by that electron. This effect causes valence electrons to be less attracted to the nucleus than they would be if there were no inner electrons.
Penetration Effect: The ability of an electron to get close to the nucleus. Atomic orbitals with different angular momentum quantum numbers (l) at the same principal energy level (n) have different degrees of penetration. For a given n, the order of penetration is s > p > d > f. Electrons in orbitals that penetrate more effectively experience a higher Z_{eff} and are more strongly attracted to the nucleus. This explains why, for instance, a 2s electron is generally lower in energy and harder to remove than a 2p electron in the same atom, as the 2s electron spends more time closer to the nucleus.
Trends in First Ionization Energy (Energy to remove the first electron)
Definition: The minimum energy required to remove one electron from a gaseous atom in its ground state (X(g) \to X^+(g) + e^-, where IE_1 is the first ionization energy).
Down a column (group): Requires less energy to remove the first electron.
Although the actual nuclear charge (Z) increases down a group, the shielding effect from increasing numbers of inner electron shells largely offsets this increase.
The valence electrons are in higher principal energy levels (n), meaning they are farther from the nucleus and experience a weaker electrostatic attraction, thus requiring less energy to remove. Consequently, Z_{eff} for valence electrons in the same group is relatively similar, but the increasing distance weakens the attraction.
Across a row (left to right, period): Generally requires more energy to remove an electron.
As one moves across a period, the number of protons in the nucleus increases, leading to a higher actual nuclear charge (Z).
The shielding provided by inner electrons remains relatively constant, and the added electrons are in the same valence shell. This results in an increasing effective nuclear charge (Z_{eff}) attracting the valence electrons more strongly, making them harder to remove.
Discontinuities: These are deviations from the general trend.
Between Groups IIA ($ns^2$) and IIIA ($ns^2 np^1$): It is easier to remove an electron in Group IIIA.
The electron removed from Group IIIA elements is from a p-orbital (e.g., np^1), which is at a slightly higher energy level and penetrates less towards the nucleus compared to an s-orbital (e.g., ns^2).
This p-electron also experiences additional shielding/repulsion from the ns^2 electrons, making it slightly easier to remove despite the overall increasing Z_{eff} across the period.
Between Groups VA ($ns^2 np^3$) and VIA ($ns^2 np^4$): It is easier to remove an electron in Group VIA.
In Group VA elements, all three p-orbitals are half-filled (e.g., px^1 py^1 p_z^1), offering a relatively stable electron configuration (as per Hund's rule).
In Group VIA elements, the added electron must go into an already occupied p-orbital (e.g., px^2 py^1 p_z^1), causing electron-electron repulsion within that orbital. This repulsion aids in the removal of one of these paired electrons, requiring less energy than expected.
Successive Ionization Energies: The energy required to remove subsequent electrons (IE2, IE3, etc.).
Generally, IE1 < IE2 < IE_3 < \dots because each subsequent electron is removed from an ion with a higher positive charge, meaning stronger electrostatic attraction to the remaining electrons.
There is a particularly large jump in ionization energy when an electron is removed from a filled noble gas core (inner shell) rather than from the valence shell, demonstrating the high stability of noble gas electron configurations.
Electron Affinity (EA): Energy change accompanying the addition of an electron to a gaseous atom (e.g., Cl(g) + e^- \to Cl^-(g), EA = -348 \text{ kJ/mol}). Typically, a more negative EA value indicates greater affinity.
Across a row (left to right): Generally becomes more exothermic (more negative EA values). Atoms with nearly full valence shells tend to have a greater attraction for an additional electron due to increasing Z_{eff} and approaching a stable noble gas configuration.
Discontinuities: These are deviations from the general trend.
Between Groups 1A ($ns^1$) and 2A ($ns^2$): The added electron in Group 2A must go into a p-orbital (e.g., np^1) instead of the s-orbital. This p-orbital is at a higher energy level and penetrates less, making the electron farther from the nucleus and experiencing repulsion from the filled ns^2 electrons. Thus, Group 2A elements often have very low or even positive electron affinities.
Between Groups 4A ($ns^2 np^2$) and 5A ($ns^2 np^3$): The added electron in Group 5A must go into an already occupied orbital (as Group 5A elements have no empty p-orbitals in their valence shell that would maintain Hund's rule configuration, e.g., to form np^4). This leads to electron-electron repulsion, making the addition of an electron less favorable (less exothermic or even endothermic).
Electronegativity
Definition: A measure of the tendency of an atom in a molecule to attract a shared pair of electrons towards itself in a chemical bond. Unlike electron affinity, which is an energy change for an isolated atom gaining an electron, electronegativity describes the pull on electrons within a bond. It is typically reported on a relative scale (e.g., Pauling scale).
Across a row (left to right): Electronegativity generally increases.
This is due to the increasing effective nuclear charge (Z_{eff}) and decreasing atomic radius. As the nucleus exerts a stronger pull on its own valence electrons, it also exerts a stronger pull on shared bonding electrons, making atoms more likely to attract electrons in a bond.
Down a column (group): Electronegativity generally decreases.
This is primarily due to the increasing atomic size and shielding effect. As the valence electrons are in higher principal energy levels and farther from the nucleus, the attractive force of the nucleus on both its own valence electrons and shared bonding electrons is weaker.
Highest Electronegativity: Fluorine (F) is the most electronegative element.
Lowest Electronegativity: Francium (Fr) and Cesium (Cs) are among the least electronegative elements.
Trends in Metallic Characteristics: Differences between metals and nonmetals are largely defined by these periodic properties (ionization energy, electron affinity, electronegativity, and atomic size).