Chapter 3

Aluminum: Low-Density Atoms Result in Low-Density Metals

A periodic property is one that is generally predictable based on an elements position within the periodic table.

The Modern Periodic Table: Its Format

- The elements are listed in order of increasing atomic number rather than increasing relative mass as they were in Mendeleev's periodic table. •Rows of the table are referred to as periods. • Columns in the table are referred to as groups or a family. -Elements in a group or family have similar properties. -NOTE: Mendeleev's periodic law predicts pattern but does NOT explain why the patterns or similarity in properties occurs. •Quantum theory explains the why. •Elements in the periodic table are classified as the following: -Metals -Nonmetals -Metalloids •The periodic table can also be divided into -main-group elements, whose properties tend to be largely predictable based on their position in the periodic table. •In the periodic table this area is labeled by a number and the letter A. -transition elements and inner transition metals or transition metals, whose properties tend to be less predictable based simply on their position in the periodic table. •In the periodic table this area is labeled by a number and the letter B.

Electron Configuration for Multi-electron Atoms

Aufbau Principle: •Energy levels and sublevels fill from lowest energy to highest: -s → p → d → f •Orbitals that are in the same sublevel have the same energy. •There can be no more than two electrons per orbital. -Pauli exclusion principle Hund's Rule: •When filling orbitals that have the same energy (degenerate), place one electron in each orbital before completing pairs.

Periodic Trend: Electron Affinities (EA) for Main Group Elements

Electron Affinity: •It is the energy associated with the addition of an electron to the valence shell of an atom that is in the gas phase. M(g) + 1 e− → M1−(g) + EA •It is defined as exothermic (−) release of energy, but may actually be endothermic (+) intake of energy. -Some alkali earth metals' and all noble gases' electron affinities are endothermic. •The more energy that is released, the larger the electron affinity. The more negative the number, the larger the EA. -General Trend for Main-Group Elements: •EA increases across a period. -EA becomes more negative from left to right. Halogens have the highest EA for any period -Summarizing Electron Affinity for Main-Group Elements •Most groups (columns) of the periodic table do not exhibit any definite trend in electron affinity. •Among the group 1A metals, however, electron affinity becomes more positive as we move down the column (adding an electron becomes less exothermic). -Alkali metals (group 1A) decrease electron affinity down the column. •Generally irregular increase in EA from second period to third period -Group 5A generally has lower EA than expected because extra electron must pair. -Groups 2A and 8A generally have very low EA because added electron goes into higher sublevel or energy level. •Electron affinity generally becomes more negative (adding an electron becomes more exothermic) as we move to the right across a period (row) in the periodic table.

General Energy Ordering of Orbitals for Multi-electron Atoms

Electrons fill from the bottom up, but for metals electrons are removed from nS-orbital before the (n-1)D-orbitals Extra stability is produced by having a filled- or half-filled orbital

Electron Configuration and Elemental Properties: The Halogens

Halogens: •They are nonmetals. •They have one fewer electron than the next noble gas. •In their reactions with metals, the halogens tend to gain an electron and attain the electron configuration of the next noble gas, forming an anion with charge 1−. •In their reactions with nonmetals, they tend to share electrons with the other nonmetal so that each attains the electron configuration of a noble gas.

Periodic Trend: Ionization Energy (Potential)

Ionization Energy (IE): •It is the minimum energy needed to remove an electron from an atom or ion in the gas phase. •It is an endothermic process (requires the input of energy to remove the electron). -Valence electron easiest to remove, lowest IE •First ionization energy = energy to remove electron from neutral atom -All atoms have first ionization energy. M(g) + IE1 → M1+(g) + 1 e- •Second IE = energy to remove from 1+ ion, etc. M+1(g) + IE2 → M2+(g) + 1 e- Ionization Energy (IE): •The larger the effective nuclear charge on the electron to be removed, the more energy it takes to remove it. •The farther the electron most probably is from the nucleus, the less energy it takes to remove it. •Trend: -First IE decreases down the group. •Valence electron is farther from nucleus. -First IE generally increases across the period. •Effective nuclear charge increases. •GENERAL trend for first ionization energy of main-group elements is that as you go across a period, ionization energy increases. -Exceptions: 2A to 3A and 5A to 6A •Exceptions are usually a result of -the type of orbital (s, p, d, or f) and its shielding ability. -repulsion factors associated with electrons occupying degenerate orbitals (i.e., p orbitals). -B has smaller first ionization energy than Be due to electron position: 2p for B and 2s for Be. -The electron in 2p orbitals has less shielding (i.e., less effective nuclear charge factor) and therefore requires less energy for its removal than an electron in a 2s orbital.

Mendeleev's Periodic Table

Mendeleev's periodic table -organized known elements of the time in a table format. •He arranged the rows so that elements with similar properties would fall in the same vertical columns. -contained some gaps, which allowed him to predict the existence (and even the properties) of yet undiscovered elements. •Mendeleev predicted the existence of an element he called eka-silicon.

Periodic Trend: Metallic Character

Metallic Elements: •Ionization energy decreases down the column. -Very low ionization energies •Good reducing agents; easy to oxidize •Very reactive; not found uncombined in nature •React with nonmetals to form salts •Compounds generally soluble in water; therefore, metal ions are found in seawater. •Electron affinity decreases down the column. Except for the noble gases, metals generally have smaller first ionization energies and nonmetals generally have larger electron affinities. Metallic Elements: •Quantum mechanics predicts that the atom's metallic character should -increase down a column because the valence electrons are not held as strongly. -decrease across a period because the valence electrons are held more strongly and the electron affinity increases. •Atomic radius increases down the column. •Melting point decreases down the column (groups 1A and 2A). -All very low MP for metals •Density increases down the column. -Except K, Mg, Ca

Behavior and Electron Configuration of Metalloids and Nonmetals

Metalloids: •They are located in the d-block area of the periodic table between the metal and nonmetal elements. -Sitting on the "steps" of the zigzag diagonal line indicated on the periodic table •Metalloids in chemical reactions can exhibit metallic or nonmetallic behaviors. •Metalloids can either lose electron(s) from s and then p orbitals to form cations or gain electrons into their p orbitals to form anions. Nonmetals: •They are located in the upper right-hand side of the periodic table. - p-block area •In chemical reactions nonmetal elements will gain electrons into the p orbitals, resulting in their ions having the same electron configuration as a noble gas at the end of their period (row). •Nonmetals form anions.

Orbital Blocks and Their Position in the Periodic Table

Summarizing Periodic Table Organization •The periodic table is divisible into four blocks corresponding to the filling of the four quantum sublevels (s, p, d, and f ). •The group number of a main-group element is equal to the number of valence electrons for that element. The row number of a main-group element is equal to the highest principal quantum number of that element

Transition and Inner Transition Metals

Transition metals (d block) and inner transition metals (f block) exhibit trends differing from those of main-group elements (s block and p block). -Because of sublevel splitting, the 4s sublevel is lower in energy than the 3d sublevel; therefore, the 4s orbital fills before the 3d orbital. -The difference in energy is not large. •Some of the transition metals have irregular electron configurations in which the ns only partially fills before the (n − 1)d or doesn't fill at all. -Therefore, their electron configuration must be found experimentally.

Radii of Atoms and Their Ions: Cations

•Cation radius is smaller than its corresponding atom radius. -The loss of electrons results in the remaining electrons in the atom experiencing a larger effective nuclear charge than the neutral atom. •Cation radius is smaller than its corresponding atom radius. -The loss of electrons results in the remaining electrons in the atom experiencing a larger effective nuclear charge than the neutral atom. •Traversing down a group increases the (n − 1) level, causing the cations to get larger. •Traversing to the right across a period increases the effective nuclear charge for isoelectronic cations, causing the cations to get smaller.

Coulomb's Law

•Coulomb's law describes the attractions and repulsions between charged particles. -For like charges, the potential energy (E) is positive and decreases as the particles get farther apart as r increases. -For opposite charges, the potential energy is negative and becomes more negative as the particles get closer together. -The strength of the interaction increases as the size of the charges increases. •Electrons are more strongly attracted to a nucleus with a 2+ charge than to a nucleus with a 1+ charge.

Shielding and Effective Nuclear Charge

•Each electron in a multi-electron atom experiences both the attraction to the nucleus and repulsion by other electrons in the atom. •These repulsions cause the electron to have a net reduced attraction to the nucleus; it is shielded from the nucleus. •The total amount of attraction that an electron feels for the nucleus is called the effective nuclear charge of the electron.

Ions: Magnetic Properties

•Electron configurations that result in unpaired electrons mean that the atom or ion will have a net magnetic field; this is called paramagnetism. -Will be attracted to a magnetic field •Electron configurations that result in all paired electrons mean that the atom or ion will have no magnetic field; this is called diamagnetism. -Slightly repelled by a magnetic field

Summarizing the Filling of Electrons in Atomic Orbitals

•Electrons occupy orbitals so as to minimize the energy of the atom; therefore, lower energy orbitals fill before higher energy orbitals. -Orbitals fill in the following order: •1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s. •Orbitals can hold no more than two electrons each. When two electrons occupy the same orbital, their spins are opposite. -This is another way of expressing the Pauli exclusion principle (no two electrons in one atom can have the same four quantum numbers). •When orbitals of identical energy are available, electrons first occupy these orbitals singly with parallel spins rather than in pairs (Hund's rule). -Once the orbitals of equal energy are half-full, the electrons start to pair.

First Ionization Potential: Exceptions to the Trend

•First ionization energy's GENERAL trend is that as you go across a period, ionization energy increases. -To ionize N, you must break up a half-full sublevel, which costs extra energy. -When you ionize O, you get a half-full sublevel, which costs less energy.

Finding Patterns: The Periodic Law and the Periodic Table

•In 1869, Mendeleev noticed that certain groups of elements had similar properties. •He found that when elements were listed in order of increasing mass, these similar properties recurred in a periodic pattern. -To be periodic means to exhibit a repeating pattern.

Electron Configuration and Ion Formation

•Ion formation can be predicted by an element's location in the periodic table. •These atoms form ions that will result in an electron configuration that is the same as that of the nearest noble gas. •Metals form cations (positively charged atoms). -Alkali metals (group 1A) form only +1 cations. -Alkaline earth metals (group 2A) form only +2 cations. -Transition, inner transition, and p-block metals form a variety of charged cations. •Nonmetals form anions (negatively charged atoms). -Halogens (group 7A) usually gain one electron to form -1 anions. -Other nonmetals can form a variety of charged anions. •Ion formation can be predicted by an element's location in the periodic table. •These atoms form ions that will result in an electron configuration that is the same as that of the nearest noble gas. •Metals form cations (positively charged atoms). -Alkali metals (group 1A) form only +1 cations. -Alkaline earth metals (group 2A) form only +2 cations. -Transition, inner transition, and p-block metals form a variety of charged cations. •Nonmetals form anions (negatively charged atoms). -Halogens (group 7A) usually gain one electron to form -1 anions. -Other nonmetals can form a variety of charged anions.

Ions: Ionic Radii Summary

•Ions in the same group have the same charge. •Ion size increases down the column. -Higher valence shell, larger •Cations are smaller than neutral atoms; anions are larger than neutral atoms. •Cations are smaller than anions. -Except Rb+ and Cs+, which are bigger than or the same size as F− and O2−. •Larger positive charge = smaller cation -For isoelectronic species -Isoelectronic = same electron configuration •Larger negative charge = larger anion -For isoelectronic species

The Periodic Law

•Mendeleev summarized these observations in the periodic law: -When the elements are arranged in order of increasing mass, certain sets of properties recur periodically.

Electron Configuration and Elemental Properties: The Metals

•Metallic elements make up the majority of the elements in the periodic table. - Alkali Metals: •They have one more electron than the previous noble gas and occupy the first column. •In their reactions, the alkali metals lose one electron, and the resulting electron configuration is the same as that of a noble gas. -Forming a cation with a 1+ charge -Alkaline Earth Metals: •They have two more electrons than the previous noble gas and occupy the second column. •In their reactions, the alkaline earth metals lose two electrons, and the resulting electron configuration is the same as that of a noble gas. -Forming a cation with a 2+ charge •Metallic elements make up the majority of the elements in the periodic table. -Transition and Inner Transition Metals: •They are located in the d-block area of the periodic table. •In chemical reactions, they will lose electron(s) from s and then d orbitals to form cations. -p-block metals: •They are located in the p-block area (left-hand side of the metalloids) of the periodic table. •In chemical reactions, they will lose electrons from the s and p orbitals to form cations.

Characteristics of Metals versus Nonmetals

•Metals -Malleable and ductile -Shiny, lustrous, reflect light -Conduct heat and electricity -Most oxides basic and ionic -Form cations in solution -Lose electrons in reactions—oxidized Nonmetals -Brittle in solid state -Dull, nonreflective, solid surface -Electrical and thermal insulators -Most oxides acidic and molecular -Form anions and polyatomic anions -Gain electrons in reactions—reduced •Metallic character is how closely an element's properties match the ideal properties of a metal. •More malleable and ductile, better conductor, and easier to ionize •Metallic character decreases left to right across a period. •Metals found at the left of the period and nonmetals to the right •Metallic character increases down the column. •Nonmetals found at the top of the middle main-group elements and metals found at the bottom

Electron Configuration: How an Atom's Electrons Occupy Orbitals

•Quantum-mechanical theory describes the behavior of electrons in atoms. •The electrons in atoms exist in orbitals. •A description of the orbitals occupied by electrons is called an electron configuration.

Electron Configuration and Quantum Theory Connection

•Schrödinger's equation showed that hydrogen's one electron occupies the lowest energy orbital in the atom. •For multi-electron atoms, the equation cannot be exactly solved because of the for electron-electron interactions that happen between two electrons. However, approximate solutions showed the orbitals to be hydrogen-like. -Two additional concepts affect multi-electron atoms: electron spin and energy splitting of sublevels. •To understand electron arrangement around an atom's nucleus, we need to account for the effects of electron spin (ms quantum number).

Electron Spin and the Pauli Exclusion Principle

•Spin is a fundamental property of all electrons. •All electrons have the same amount of spin. •The orientation of the electron spin is quantized—it can be only in one direction or its opposite. -Spin up or spin down •The electron's spin adds a fourth quantum number to the description of electrons in an atom, called the spin quantum number, ms. -The spin quantum number is not part of the Schrödinger equation. •ms can have values of +½ or −½. •Orbital diagrams use a square to represent each orbital and a half-arrow to represent each electron in the orbital. •By convention, -a half-arrow pointing up is used to represent an electron in an orbital with spin up; -a half-arrow pointing down is used to represent an electron in an orbital with spin down. •Spins must cancel in an orbital. -Paired meaning the two electrons in an orbital must have opposite spins (i.e. one with the magnetic field the other against the magnetic field). The Pauli Exclusion Principle: •No two electrons in an atom may have the same set of four quantum numbers. •Therefore, no orbital may have more than two electrons, and they must have opposite spins.

Shielding and Penetration

•The closer an electron is to the nucleus, the more attraction it experiences. •The better an outer electron is at penetrating through the electron cloud of inner electrons, the more attraction it will have for the nucleus. •The degree of penetration is related to the orbital's radial distribution function. -In particular, the distance the maxima of the function are from the nucleus •Penetration causes the energies of sublevels in the same principal level to not be degenerate. •In the fourth and fifth principal levels, the effects of penetration become so important that the s orbital lies lower in energy than the d orbitals of the previous principal level. •The energy separations between one set of orbitals and the next become smaller beyond the 4s orbital. -The ordering can therefore vary among elements, causing variations in the electron configurations of the transition metals and their ions.

Electron Configuration, Valence Electrons, and the Periodic Table

•The electrons in all the sublevels with the highest principal energy shell are called the valence electrons. •Electrons in lower energy shells are called core electrons. •One of the most important factors in the way an atom behaves, both chemically and physically, is the number of valence electrons

Core Electrons, Valence Electrons, and the Periodic Table

•The group number corresponds to the number of valence electrons. •The length of each "block" is the maximum number of electrons the sublevel can hold. •The period number corresponds to the principal energy level of the valence electrons.

Electron Configuration and Elemental Properties: Noble Gases

•The noble gases have eight valence electrons. -Except for He, which has only two electrons •They are especially nonreactive. -He and Ne are practically inert. •The reason the noble gases are so nonreactive is that the electron configuration of the noble gases is especially stable.

Electron Configuration and Elemental Properties

•The properties of the elements follow a periodic pattern. -Elements in the same column have similar properties. -The elements in a period show a pattern that repeats. •The quantum-mechanical model explains this because the number of valence electrons and the types of orbitals they occupy are also periodic.

Electron Spatial Distribution and Sublevel Splitting

•The radial distribution function shows that the 2s orbital penetrates more deeply into the 1s orbital than does the 2p. •The weaker penetration of the 2p sublevel means that electrons in the 2p sublevel experience more repulsive force; they are more shielded from the attractive force of the nucleus. •The deeper penetration of the 2s electrons means electrons in the 2s sublevel experience a greater attractive force to the nucleus and are not shielded as effectively.

Sublevel Energy Splitting in Multi-electron Atoms

•The sublevels in each principal energy shell of hydrogen all have the same energy or other single electron systems. •Orbitals with the same energy (E) are said to be degenerate. •For multi-electron atoms, the energies of the sublevels are split. -Caused by charge interaction, shielding, and penetration •The lower the value of the l quantum number (orbital quantum number), the less energy the sublevel has. E (s orbital (l = 0)) < E (p orbital (l = 1)) < E (d orbital (l = 2)) < E (f orbital (l = 3))

Trends in Second and Successive Ionization Potentials

•They depend on the number of valence electrons an element has. •Removal of each successive electron costs more energy. -Shrinkage in size due to having more protons than electrons -Outer electrons closer to the nucleus; therefore harder to remove •Regular increase in energy for each successive valence electron •Large increase in energy when core electrons are removed

Radii of Atoms and Their Ions: Anions

•When atoms form anions, electrons are added to the valence shell. •These "new valence electrons" experience a smaller effective nuclear charge than the "old valence electrons." •When atoms form anions, electrons are added to the valence shell. •These "new valence electrons" experience a smaller effective nuclear charge than the "old valence electrons." •The result is that the anion is larger than the atom. Traversing down a group increases the n level, causing the anions to get larger. •Traversing to the right across a period increases the effective nuclear charge for isoelectronic anions, causing the anions to get smaller.