Chapter 7: Periodic Properties of the Elements – Study Notes

Periodic Properties of the Elements

Introduction

This chapter explores the periodic trends and properties of elements as organized in the periodic table. Understanding these trends is essential for predicting chemical behavior and reactivity.

Development of the Periodic Table

Historical Background

• Dmitri Mendeleev and Lothar Meyer independently proposed organizing elements by recurring chemical and physical properties.

• Mendeleev's table was based on atomic mass, the most fundamental property known at the time.

• Mendeleev predicted the existence and properties of undiscovered elements, such as eka-silicon (later identified as germanium).

Table: Comparison of Predicted and Observed Properties (Eka-Silicon vs. Germanium)

Property	Mendeleev's Prediction (Eka-Silicon, 1871)	Observed Germanium (1886)
Atomic Weight	72	72.59
Density (g/cm3)	5.5	5.35
Specific Heat (J/g·K)	0.305	0.309
Melting Point (°C)	High	947
Appearance	Dark gray	Grayish white
Type of Oxide	XO2	GeO2
Density of Oxide (g/cm3)	4.7	4.70
Formula of Chloride	XCl4	GeCl4
Boiling Point of Chloride (°C)	A little under 100	84

Additional info: This table demonstrates the accuracy of Mendeleev's predictions and the utility of periodic trends in forecasting element properties.

Atomic Number and Modern Periodic Law

Discovery and Importance

• Ernest Rutherford discovered the nuclear atom, leading to a deeper understanding of atomic structure.

• Henry Moseley established the concept of atomic number (number of protons) as the basis for periodicity, using X-ray experiments.

• Periodic properties are now organized by atomic number, not atomic mass.

Periodicity

Definition and Key Properties

• Periodicity is the recurring pattern of element properties as a function of atomic number.

• Main periodic properties discussed:

  • Sizes of atoms and ions

  • Ionization energy

  • Electron affinity

  • Chemical property trends within groups

Additional info: These trends are explained by changes in effective nuclear charge, electron configuration, and atomic structure.

Effective Nuclear Charge ($Z_{eff}$)

Concept and Calculation

Many atomic properties depend on the interactions between valence electrons and the nucleus, which are influenced by both attraction to the nucleus and repulsion from other electrons.

• Effective nuclear charge ($Z_{eff}$) is the net positive charge experienced by an electron in a multi-electron atom.

• Calculated as: $Z_{eff} = Z - S$ where $Z$ is the atomic number and $S$ is the screening constant (approximate number of core electrons).

• $Z_{eff}$ increases left to right across a period and slightly increases down a group.

Sizes of Atoms and Ions

Atomic Radius

• Nonbonding atomic radius (van der Waals radius): Half the shortest distance between nuclei during atomic collisions.

• Bonding atomic radius (covalent radius): Half the distance between nuclei in a bond.

• Bonding atomic radius decreases left to right across a period (due to increasing $Z_{eff}$) and increases down a group (due to increasing principal quantum number $n$).

Ionic Radius

• Cations are smaller than their parent atoms (loss of electrons reduces electron-electron repulsion).

• Anions are larger than their parent atoms (gain of electrons increases electron-electron repulsion).

• Ionic size depends on nuclear charge, number of electrons, and the orbitals occupied by valence electrons.

Isoelectronic Series

• An isoelectronic series consists of ions with the same number of electrons.

• Within an isoelectronic series, ionic size decreases as nuclear charge increases.

Ionization Energy ($I$)

Definition and Trends

• Ionization energy is the minimum energy required to remove an electron from a gaseous atom in its ground state.

• First ionization energy ($I_1$): Energy to remove the first electron.

• Second ionization energy ($I_2$): Energy to remove the second electron, and so on.

• Successive ionization energies increase, especially after all valence electrons are removed.

• General trends:

  • $I$ increases across a period (left to right).

  • $I$ decreases down a group.

  • Exceptions occur when electrons enter new sub levels or pair within orbitals, due to electron repulsion effects.

Equation: $M(g) \rightarrow M^+(g) + e^-$

Electron Affinity (EA)

Definition and Trends

• Electron affinity is the energy change when an electron is added to a gaseous atom.

• Usually exothermic (negative EA), indicating energy is released.

• General trend: EA becomes more negative across a period, with notable exceptions (full or half-full sub levels).

Equation: $X(g) + e^- \rightarrow X^-(g)$

Example: $Cl(g) + e^- \rightarrow Cl^-(g)$, $EA = -349\,kJ$

Metals, Nonmetals, and Metalloids

Classification and Properties

• Metals: Shiny, conductive, malleable, ductile, mostly solids at room temperature, low ionization energies, form cations.

• Nonmetals: Found on the right side of the periodic table, can be solid, liquid, or gas, dull, brittle, poor conductors, large negative electron affinity, form anions.

• Metalloids: Exhibit properties intermediate between metals and nonmetals; several are semiconductors (e.g., silicon).

Table: Comparison of Metals and Nonmetals

Property	Metals	Nonmetals
Luster	Shiny, silvery	No luster, various colors
Malleability/Ductility	Malleable, ductile	Brittle (solids), some hard/soft
Conductivity	Good conductors	Poor conductors
Oxide Type	Ionic, basic	Molecular, acidic
Ion Formation	Form cations	Form anions/oxyanions

Group Trends

Overview

• Elements in a group share similar properties due to similar valence electron configurations.

• Key groups:

  • Group 1A: Alkali metals – soft, metallic solids, low densities and melting points, highly reactive, form +1 cations.

  • Group 2A: Alkaline earth metals – higher densities and melting points than alkali metals, form +2 cations, reactivity increases down the group.

  • Group 6A: Oxygen group – includes nonmetals (oxygen, sulfur, selenium), metalloid (tellurium), and metal (polonium); metallic character increases down the group.

  • Group 7A: Halogens – typical nonmetals, highly negative electron affinities, form anions, react with metals to form halides.

  • Group 8A: Noble gases – monatomic gases, very high ionization energies, low electron affinities, chemically inert.

Special Cases

• Hydrogen: Electron configuration suggests metallic character, but its properties are nonmetallic; forms H+ and H- ions.

• Oxygen: Exists as O2 (colorless, odorless) and O3 (ozone, pale blue, sharp odor); forms oxides, peroxides, and super-oxides.

Summary Table: Group Properties

Group	Key Properties
1A (Alkali Metals)	Soft, low density, low melting point, highly reactive, form +1 cations
2A (Alkaline Earth Metals)	Higher density/melting point than 1A, form +2 cations, reactivity increases down group
6A (Oxygen Group)	Nonmetals, metalloid, metal; metallic character increases down group
7A (Halogens)	Nonmetals, high electron affinity, form anions, react with metals
8A (Noble Gases)	Monatomic gases, inert, high ionization energy, low electron affinity

Additional info: Understanding periodic trends is essential for predicting element reactivity, bonding, and chemical behavior in various contexts.