The modern periodic table emphasizes the electronic structure of atoms. The original periodic table emphasized the reactivity of the elements. In this chapter we will learn about the connection between the two, and how the electronic structure is related to the macroscopic properties of the elements. 6.1: Early History of the Periodic Table This page reviews methods for organizing library books through the Dewey Decimal and Library of Congress systems. It then shifts to early chemical element organization, highlighting Johann Dobereiner's triad system based on chemical properties and atomic masses, followed by John Newlands' "Law of Octaves," which arranged elements by atomic mass with recurring properties every eighth element. Both methods had their drawbacks but were foundational for the evolution of the periodic table. 6.2: Mendeleev's Periodic Table Dmitri Mendeleev published his periodic table in 1869, organizing elements by atomic mass and highlighting patterns in their chemical properties. He utilized a card system for rearranging elements, leading to significant insights, such as placing tellurium before iodine based on chemical similarity. Mendeleev also predicted elements like eka-aluminum, which was later identified as gallium. He is recognized for establishing the periodic law, and element 101, mendelevium, was named in his honor. 6.3: Periodic Law This page discusses the periodic table's organization of elements by increasing atomic number and its correlation with chemical and physical properties. Initially based on atomic mass by Mendeleev, it was redefined by Moseley's work in 1913 linking atomic number to X-ray wavelengths. 6.4: Modern Periodic Table- Periods and Groups This page covers the evolution of dictionaries and the periodic table. Dictionaries adapt to new words and usages for effective communication, while the periodic table, originally created by Mendeleev and Moseley, has expanded with new elements such as Nihonium and Moscovium. It is structured by atomic number into periods and groups, with standardized group numbering by the International Union of Pure and Applied Chemistry to enhance clarity. 6.5: Metals This page discusses the characteristics of various metal screws, highlighting their differences in size, shape, and type. It covers the classification of metals based on physical properties such as conductivity and malleability. Key points include gold's value in jewelry for its softness, copper's role in electrical conduction, and the toxicity of mercury, the only liquid metal at room temperature, which has fallen out of common use. 6.6: Nonmetals This page discusses nonmetals, which are elements with poor heat and electricity conductivity and can be solids, liquids, or gases. Examples include sulfur (used in gunpowder and rubber), bromine (used as a disinfectant and flame retardant), and helium (used in balloons and cooling superconductors). Nonmetals display varied properties and applications that differ significantly from metals. 6.7: Metalloids This page discusses metalloids, which possess properties between metals and nonmetals. Key examples are silicon, valuable in electronics for its luster and brittleness; boron, used in borosilicate glass and lightweight materials; arsenic, historically associated with poison but now declining in use due to toxicity; and antimony, which strengthens alloys and is crucial in battery production. The page emphasizes the complexity of classifying elements in chemistry. 6.8: Blocks of the Periodic Table This page explains the structure of the periodic table, which comprises seven horizontal rows or periods, each determined by the number of electrons that can fill its sublevels (s, p, d, f). The first period has only the 1s sublevel for two elements, while longer periods, like the fourth and sixth, accommodate more electrons. Understanding the sublevel filling sequence is essential for determining an element's period and block based on its electron configuration. 6.9: Hydrogen and Alkali Metals This page outlines the properties of alkali metals in Group I of the periodic table, including their outer electron configuration with one \(s\) orbital electron, resulting in high reactivity. It stresses the dangers of water reactions, which can lead to violent explosions, and underscores the importance of safety precautions during demonstrations in chemistry classes. The main point is the vigorous reaction of these metals with water and their similar electron configurations. 6.10: Alkaline Earth Metals This page discusses the structural role of calcium compounds in oyster shells and bones, highlights the properties of alkaline earth metals in Group 2, and notes their lower reactivity compared to Group 1 elements. It outlines their uses, such as magnesium in fireworks, calcium in cement, and strontium in fireworks. Additionally, the page mentions radium as a radioactive and unstable alkaline earth metal that shares characteristics with barium. 6.11: Noble Gases This page discusses noble gases, such as helium, xenon, and radon, which are used in neon lights for colorful displays. These gases are chemically inert and exist as monatomic gases at room temperature. Although traditionally thought to be unreactive, noble gases can form compounds, with xenon being the first in 1962. When an electric current passes through them, they emit unique colors, enhancing their use in illumination, though radon is excluded due to its radioactivity. 6.12: Halogens This page discusses halogens, including their high reactivity, electron configuration with seven valence electrons, and physical states at room temperature—fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. It highlights their tendency to exist in nature combined with other substances, such as salts found in the ocean. 6.13: Transition Elements This page compares stock cars and racing cars of the same model, emphasizing the importance of analyzing their internal components. It also explains transition elements in Groups 3-12 of the periodic table, noting their unique electron configurations with partially filled \(d\) sublevels. Transition metals are characterized by typical metallic properties, lower reactivity, and colorful compounds resulting from \(d\) electron transitions absorbing visible light. 6.14: Lanthanides and Actinides This page discusses Russian nesting dolls as a metaphor for the layered structure of lanthanides and actinides in the periodic table. Lanthanides (atomic numbers 58-71) are used in alloys and optical devices, while actinides (atomic numbers 90-103) are radioactive and serve energy applications such as nuclear power. Both groups, known as inner transition elements, exhibit unique properties and have various industrial applications. 6.15: Periodic Trends- Atomic Radius This page explains that the atomic radius measures an atom's size as half the distance between bonded identical atoms. It notes that atomic radii decrease across a period due to increased nuclear charge, while they increase down a group due to additional energy levels. Measurements are given in picometers, highlighting size differences among elements, exemplified by the small radius of hydrogen compared to the larger radius of potassium. 6.16: Ion This page explains the northern lights, which are produced by charged particles (ions) interacting with Earth's magnetic field. Ions form when atoms gain or lose electrons during electron transfer, resulting in positive or negative charges. These reactive ions can combine to create neutral compounds and are influenced by magnetic fields, as illustrated by the northern lights phenomenon. 6.17: Periodic Trends - Ionization Energy This page discusses two topics: sheep behavior, highlighting their tendency to herd influenced by attraction and external factors, and the concept of ionization energy in chemistry, explaining how it varies across periods and groups in the periodic table due to nuclear attraction and atomic size. Understanding these trends is essential for predicting atomic behavior. 6.18: Electron Shielding This page discusses roller derby, where a jammer scores points by passing opponents while blockers try to stop them. It also explains electron shielding in atoms, detailing how inner electrons affect attraction to outer electrons and influence ionization energy, with examples from lithium and aluminum, as well as variations in ionization energy across groups 13-15 and 15-16, highlighting the effects of electron configuration and pairing. 6.19: Periodic Trends - Electron Affinity This page explains electron affinity as the energy released when an atom gains an electron, typically measured in negative values in the gaseous state. It describes trends where electron affinities become more negative across a period and less negative down a group, highlighting halogens as having the highest affinities due to their stable electron configurations. The text also acknowledges some exceptions to these trends. 6.20: Periodic Trends - Ionic Radii This page discusses the sale of shelled peanuts for snacking or cooking and explains ionic radius, detailing how ion sizes vary with electron loss or gain. Cations become smaller when electrons are removed due to reduced energy levels and a higher proton-electron ratio, while anions grow larger with added electrons due to increased repulsion. It notes that ionic radius measurements often utilize crystal lattice structures, with differences stemming from the measurement techniques used. 6.21: Periodic Trends- Electronegativity This page explains electronegativity, defining it as an atom's ability to attract electrons. It notes that electronegativity increases across periods and decreases down groups, highlighting fluorine as the most electronegative element. Metals generally have low electronegativities due to their tendency to lose electrons, while nonmetals gain electrons, leading to higher values. 6.22: Periodic Trends - Metallic and Nonmetallic Character This page examines U.S. eating habits from 1971 to 2000 and their health implications while exploring metallic and non-metallic character in the periodic table. It notes that metallic character increases down a group due to lower ionization energy and larger atomic radius, whereas non-metallic character rises across the table, peaking with fluorine as the most reactive nonmetal.