Development of the Periodic Table and Periodic Trends Study Guide

Organization of the Periodic Table

Periods: * These are the horizontal rows on the periodic table. * The period number for an element is equal to the total number of occupied energy levels (shells).
Groups: * These are the vertical columns on the periodic table. * Elements within the same group possess the same number of valence electrons. * Because they have the same number of valence electrons, elements in the same group exhibit similar chemical properties.

Reactivity of Elements

Definition: Reactivity refers to the ability of elements to form chemical bonds.
Determinant: The reactivity of an element is primarily determined by its number of valence electrons.
The Stable Octet: * Nearly all atoms achieve stability when they possess $8$ valence electrons, creating a full valence shell. * Exceptions: Hydrogen ( $H$ ) and Helium ( $He$ ) are exceptions to the octet rule; they are stable with only $2$ valence electrons. * Period 1 Stability: Elements in Period 1 require only $2$ electrons to achieve a full valence shell and reach stability.
Reactivity Levels: The closer an element is to possessing a stable octet, the more reactive that element is.
Comparative Example: Sodium ( $Na$ ) and Potassium ( $K$ ) have similar chemical properties because they are located in the same group and have the same number of valence electrons. Magnesium ( $Mg$ ) does not share these same properties as it has a different number of valence electrons.

Properties of Metals, Nonmetals, and Metalloids

Properties of Metals: * Malleable: Metals can be hammered or rolled into thin sheets. * Ductile: Metals can be drawn into thin wires. * Conductivity: They are excellent conductors of both heat and electricity. * Luster: They possess a shiny appearance. * Ion Formation: Metals tend to lose electrons to form positive ions, known as cations. * Phase at STP: Metals are solid at Standard Temperature and Pressure ( $STP$ ), with the exception of Mercury ( $Hg$ ), which is a liquid. * Structural Mechanics: Metals are malleable and ductile because they consist of layers of atoms that can slide over one another. Their high conductivity is due to a "sea" of mobile, delocalized electrons.
Properties of Nonmetals: * Conductivity: They are poor conductors of heat and electricity. * Brittle: Nonmetals shatter or crumble when struck. * Dull: They lack luster and are not shiny. * Ion Formation: Nonmetals tend to gain electrons to form negative ions, known as anions. * Example: Carbon-graphite (the material used as "lead" in pencils).
Properties of Metalloids: * Semiconductors: They act as good or moderate conductors of electricity. * Hybrid Characteristics: They exhibit properties of both metals (such as luster) and nonmetals (such as being brittle). * Applications: Used extensively in the manufacturing of computer microchips. * Examples: Silicon ( $Si$ ) and Boron ( $B$ ).

Specific Element Groups

Group 1: Alkali Metals: * Possess $1$ valence electron. * Lose $1$ electron to form ions with a charge of $+1$ . * Extremely reactive with water. * Francium ( $Fr$ ) is the most reactive metal.
Group 2: Alkaline Earth Metals: * Possess $2$ valence electrons. * Lose $2$ electrons to form ions with a charge of $+2$ . * Fairly reactive in water.
Groups 3-12: Transition Metals: * Characterized as the least reactive metals. * Distinctively form colored ions when in solution.
Group 17: Halogens: * Possess $7$ valence electrons. * Gain $1$ electron to form ions with a charge of $-1$ . * Fluorine ( $F$ ) is the most reactive nonmetal.
Group 18: Noble Gases: * Considered unreactive or inert. * Possess a stable octet ( $8$ valence electrons). * Exception: Helium ( $He$ ) is stable with only $2$ valence electrons. * They always exist as monoatomic species.
Hydrogen ( $H$ ): * Not officially part of any specific group. * Classified as a nonmetal. * Exists as a gas at $STP$ .

Vocabulary of Periodic Trends

Atomic Radius: The distance measured from the center of the nucleus to the outermost electrons of an atom.
Ionic Radius: The distance measured from the center of the nucleus to the outermost electrons of an ion.
Ionization Energy: The energy required to remove an electron. * First Ionization Energy: The energy needed specifically to remove the first (most loosely bound) electron from an atom. This value is always the lowest compared to subsequent ionization energies.
Electron Affinity: The energy change that occurs when an atom gains an electron.
Electronegativity: The measure of an atom's ability to attract an electron within a chemical bond. * Measured using the Pauling Scale, ranging from $0$ (least ability) to $4$ (greatest ability). * Follows the same trend as electron affinity.
Effective Nuclear Charge ( $Z_{eff}$ ): A measure of the pull of the protons in the nucleus on the valence electrons of an atom or ion. * Formula: $Z_{eff} = \text{atomic number} - \text{shielding electrons}$ . * Shielding Electrons: For the purposes of these notes, these are the inner core electrons (all non-valence shell electrons) that block some of the nuclear pull from reaching the outermost electrons.

Trends in Atomic Radius

Data Source: Measured values are found in Table S.
Trends Across a Period: * Trend: Atomic radius decreases. * Explanation: As the atomic number increases across a period, the nuclei have a greater nuclear pull or higher effective nuclear charge ( $Z_{eff}$ ). Larger positive charges in the nucleus pull the electrons closer. * Summary Mnemonic: (Period, Proton, Pull).
Trends Down a Group: * Trend: Atomic radius increases. * Explanation: There is a greater number of energy shells as you move down a group (indicated by the increment in Period number).

Trends in Ionic Radius

Metal Ion Size: * The ion radius is smaller than the corresponding atomic radius. * Explanation: Metals lose electrons to form positive cations, resulting in a reduction of electron-electron repulsion or the loss of an entire shell.
Nonmetal Ion Size: * The ion radius is larger than the corresponding atomic radius. * Explanation: Nonmetals gain electrons to form negative anions, increasing electron-electron repulsion within the shell.
Isoelectronic Species: * These are two different species that share the same electronic structure (the same electron configuration and number of valence electrons). * Example: $Mg^{+2}$ and $Na^{+1}$ . * Trend for Isoelectronic Species: Atoms or ions with more protons (a higher atomic number) will have smaller radii because the higher effective nuclear charge pulls the electrons in more tightly.

Trends in Ionization Energy

Definition: Energy required to remove the most loosely bound electron in the valence shell.
Data Source: Measured values are found in Table S.
Trends Across a Period: * Trend: Ionization energy increases. * Explanation: A stronger proton (nuclear) pull makes it increasingly difficult to remove electrons from the atom.
Trends Down a Group: * Trend: Ionization energy decreases. * Explanation: Larger atomic radii result in less nuclear pull on the outer electrons due to the electron shielding effect from inner core electrons.

Trends in Electronegativity

Definition: Quantitative measure of the attraction for electrons in a bond.
Data Source: Measured values are found in Table S.
Key Reference: Fluorine is the most electronegative element with a value of $4.0$ on the scale of $0$ to $4$ . The closer an atom is to Fluorine on the table, the higher its electronegativity.
Trends Across a Period: * Trend: Electronegativity increases. * Explanation: There is a greater proton pull ( $Z_{eff}$ ) available to attract incoming electrons.
Trends Down a Group: * Trend: Electronegativity decreases. * Explanation: Larger atomic radii mean there is less nuclear pull available to attract electrons to the valence shell because of the electron shielding effect.