Chapter 3 Electronic Structure and the Periodic Law

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Chapter 3 Electronic Structure and the Periodic Law

Learning Objectives Locate elements in the periodic table on the basis of group and period designations Determine the number of electrons in designated atomic orbitals, subshells, or shells Determine the number of valence shell electrons and the electronic structure for atoms, and relate this information to the location of elements in the periodic table Determine the following for elements: the electronic configuration of atoms, the number of unpaired electrons in atoms, and the identity of atoms based on provided electronic configurations

Learning Objectives (continued) Determine the shell and subshell locations of the distinguishing electrons in elements, and based on their location in the periodic table, classify elements into the categories given in Figures 3.10 and 3.12

Periodic Law Statement about the behavior of the elements when they are arranged in a specific order Present form - Elements with similar chemical properties occur at regular (periodic) intervals when the elements are arranged in order of increasing atomic numbers

Periodic Table Arrangement of the elements in a table based on the periodic law In a modern periodic table, elements with similar chemical properties are found in vertical columns called groups or families 18 groups/families 7 periods

Group or Family of the Periodic Table Designations The U.S. system uses a Roman numeral and a letter (either A or B) at the top of the column The IUPAC (but not universally used) system uses only a number from 1 to 18

Period of the Periodic Table Horizontal row of elements arranged according to increasing atomic numbers Numbered from top to bottom

Modern Periodic Table Elements 58–71 and 90–103 are not placed in their correct periods but are located below the main table

Group and Period Identification Elements and the periodic table Each element belongs to a group and period of the periodic table Examples of group and period location for elements Calcium, Ca, element 20: group IIA(2), period 4 Silver, Ag, element 47: group IB(11), period 5 Sulfur, S, element 16: group VIA(16), period 3

Example 3.1 - Groups and Periods of the Periodic Table Identify the group and period to which each of the following belongs P Cr Element number 30 Element number 53

Example 3.1 - Solution Phosphorus (P) is in group VA(15) and period 3 Chromium (Cr) is in group VIB(6) and period 4 Element with atomic number 30 is zinc (Zn), which is found in groups IIB(12) and period 4 Element number 53 is iodine (I), found in group VIIA(17) and period 5

Bohr's Theory Bohr proposed that the electron in a hydrogen atom moved in any one of a series of circular orbits around the nucleus Electron could change orbits only by absorbing or releasing energy Absorption moves electrons to higher-energy orbits Release moves electrons to lower-energy orbits

Quantum Mechanical Model Precise paths of electrons moving around the nucleus cannot be determined accurately Location and energy of electrons around a nucleus is specified using shell, subshell, and orbital

Shell Location and energy of electrons around a nucleus that is designated by a value for n Value of n can be 1, 2, 3, etc. Higher n values for a shell correspond to higher energies and greater distances from the nucleus for the electrons of the shell

Subshell Each shell is made up of one or more subshells Designated by a letter from the group s, p, d, and f Combination of shell number and subshell letter of the subshell is used to identify subshells For example, a p subshell located in the fourth shell (n = 4) would be referred to as a 4p subshell

Subshell (continued) Number of subshells located in a shell is the same as the number of the shell Shell number 3 (n = 3) contains three subshells, designated 3s, 3p, and 3d Electrons located in a subshell are identified by using the same designation as the subshell Electrons in a 3d subshell are called 3d electrons

Atomic Orbitals Volume of space around atomic nuclei in which electrons of the same energy move Groups of orbitals with the same n value form subshells Have different shapes, depending on the energy of the electrons they contain All s subshells consist of a single s orbital All p subshells consist of three p orbitals All d subshells consist of five d orbitals All f subshells consist of seven f orbitals

Figure 3.3 - Shapes of Typical s, p, and d Orbitals

Atomic Orbitals (continued) According to the quantum mechanical model, all types of atomic orbitals can contain a maximum of two electrons Single d orbital can contain a maximum of two electrons, and a d subshell that contains five d orbitals can contain a maximum of 10 electrons

Energy of Electrons Increases with increasing n value Electron in the third shell (n = 3) has more energy than an electron in the first shell (n = 1) For equal n values but different subshells, the energy of electrons in orbitals increases in the order s, p, d, and f 4p electron has more energy than a 4s electron

Table 3.1 - Relationship between Shells, Subshells, Orbitals, and Electrons

Example 3.3 - Shells and Subshells of Atoms Determine the following for the third shell of an atom: Number of subshells Designation for each subshell Number of orbitals in each subshell Maximum number of electrons that can be contained in each subshell Maximum number of electrons that can be contained in the shell

Example 3.3 - Solution Number of subshells is the same as the number used to designate the shell Therefore, the third shell contains three subshells Subshells increase in energy according to the order s, p, d, f Subshells in the shell are therefore designated 3s, 3p, and 3d Number of orbitals in the subshells is 1, 3, and 5 because s subshells always contain a single orbital, p subshells always contain three orbitals, and d subshells always contain five orbitals

Example 3.3 - Solution (continued) Each atomic orbital can contain a maximum of electrons, independent of the type of orbital under discussion Therefore, the 3s subshell (one orbital) can hold a maximum of two electrons, the 3p subshell (three orbitals) a maximum of 6 electrons, and the 3d subshell (five orbitals) a maximum of 10 electrons Maximum number of electrons that can be contained in the shell is simply the sum of the maximum number for each subshell, 2 + 6 + 10 = 18

Join In (2) Correct Answer 1. shell

Join In (3) Which of the following can contain 10 electrons? 4s orbital 5p subshell shell 2 3d subshell Correct Answer 4. 3d subshell

Valence Shell Outermost (highest-energy) shell of an element that contains electrons Atoms with the same number of electrons in the valence shell have similar elemental chemical properties Members of Group IIA(2)

Table 3.2 - Electron Occupancy of Shells

Electron Shell Occupancy What do magnesium and calcium have in common? 2 electrons in valence shell What predictions can be made about the number of electrons in strontium's valence shell? Sr has similar chemical properties to Mg and Ca, so it is likely to have 2 electrons in its valence shell What other element on this chart has similar properties to Mg, Ca, and Sr? Beryllium

Example 3.4 - Number of Electrons in Valence Shells for Atoms Referring to Table 3.2, indicate the number of electrons in the valence shell of elements in groups IA(1), IIA(2), IIIA(13), and IVA(14)

Example 3.4 - Solution According to Table 3.2, the elements in group IA(1) are hydrogen, lithium, sodium, and potassium Each element has one electron in the valence shell Hydrogen belongs in group IA(1) on the basis of its electronic structure, but its properties differ significantly from other group members Group IIA(2) elements are beryllium, magnesium, and calcium Each has two electrons in the valence shell Group IIIA(13) elements are boron and aluminum Both have three electrons in the valence shell Group IVA(14) elements are carbon and silicon Each has four valence-shell electrons

Electronic Configurations Detailed arrangement of electrons in atoms indicated by a specific notation, 1s22s22p4, etc. Occupied subshells are indicated by their identifying number and letter such as 2s and the number of electrons in the subshell is indicated by the superscript on the letter Thus, in 1s22s22p4, the 2s2 notation indicates that the 2s subshell contains two electrons

Subshell-Filling Order Electrons will fill subshells in the order of increasing energy of the subshells A 1s subshell will fill before a 2s subshell Order must obey Hund's rule and the Pauli exclusion principle

Hund's Rule Electrons will not join other electrons in an orbital if an empty orbital of the same energy is available for occupancy Thus, the second electron entering a p subshell will go into an empty p orbital of the subshell rather than into the orbital that already contains an electron

Pauli Exclusion Principle Electrons behave as if they spin on an axis Only electrons spinning in opposite directions (indicated by ↑ and ↓) can occupy the same orbital within a subshell

Filling Order When it is remembered that each orbital of a subshell can hold a maximum of two electrons and that Hund's rule and the Pauli exclusion principle are followed, it results in the following filling order for the first 10 electrons: H He Li Be B C N Ne

Figure 3.7 - Relative Energies and Electron-Filling Order for Shells and Subshells

Electron-Filling Order for Shells and Subshells Filling order for any number of electrons is obtained by following the arrows in the diagram Shells are represented by large rectangles Subshells are represented by colored rectangles Orbitals within the subshells are represented by circles

Filling Order Memory Aid Diagram provides a compact way to remember the subshell filling order Correct order is obtained by following the arrows from top to bottom of the diagram, going from the arrow tail to the head, and then from the next tail to the head, etc.

Filling Order Memory Aid (continued) Maximum number of electrons each subshell can hold must be remembered s subshells can hold 2 p subshells can hold 6 d subshells can hold 10 f subshells can hold 14

Filling Order and Periodic Table Order of subshell filling matches the order of subshell blocks on the periodic table if the fill occurs in the order of increasing atomic numbers

Electron Configuration Example The following electronic configurations result from the correct use of any of the diagrams given earlier Magnesium, Mg, 12 electrons: 1s22s22p63s2 Silicon, Si, 14 electrons: 1s22s22p63s23p2 Iron, Fe, 26 electrons: 1s22s22p63s23p64s23d6 Gallium, Ga, 31 electrons: 1s22s22p63s23p64s23d104p1

Noble Gas Configurations With the exception of helium, all noble gases (group VIIIA) have electronic configurations that end with completely filled s and p subshells of the highest occupied shell Can be used to write electronic configurations in an abbreviated form in which the noble gas symbol enclosed in brackets is used to represent all electrons found in the noble gas configuration

Noble Gas Configurations: Example Magnesium: [Ne]3s2 Symbol [Ne] - Represents the electronic configuration of neon, 1s22s22p6 Iron: [Ar]4s23d6 Symbol [Ar] - Represents the electronic configuration of argon, 1s22s22p63s23p6 Gallium: [Ar]4s23d104p1

Example 3.6 - Electronic Configurations and Unpaired Electrons Write the electronic configuration for an atom that contains 7 electrons, and indicate the number of unpaired electrons

Example 3.6 - Solution Correct filling order of subshells is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p Even though s subshells can hold 2 electrons, p subshells 6, d subshells 10, and f subshells 14, only enough subshells to hold 7 electrons will be used Therefore, the configuration is written as shown, starting on the left, with circles representing orbitals and arrows representing electrons It is apparent that no subshells beyond 2p are needed because that subshell contains only 3 electrons

Example 3.6 - Solution (continued) Note that both the 1s and 2s subshells are full; the 2p subshell is half-full, with one electron in each of the three orbitals (Hund's rule) These three electrons are unpaired 1s2 2s2 2p3

Example 3.7 - Abbreviated Electronic Configurations Write abbreviated electronic configuration for an atom that contains 7 electrons

Example 3.7 - Solution From Example 3.6, the configuration is 1s22s22p3 Two electrons in the 1s subshell represent the noble gas configuration of helium (He), so the configuration can be written as [He]2s22p3

Join In (7) Which of the following electron configurations matches aluminum? [Ne]3s23p1 [Ne]3p3 [Ar]3p5 [Ar]3s22p1 Correct Answer 1. [Ne] 3s2 3p1

Classification of Elements Representative element: Element in which the distinguishing electron is found in an s or a p subshell Transition element: Element in which the distinguishing electron is found in a d subshell Inner-transition element: Element in which the distinguishing electron is found in an f subshell

Figure 3.12 - Locations of Metals, Nonmetals, and Metalloids in the Periodic Table

Metals and Nonmetals Metals Elements found in the left two-thirds of the periodic table Properties - High thermal and electrical conductivities, high ductility and malleability, and metallic luster Nonmetals Elements found in the right one-third of the periodic table Properties - Opposite those of metals and occur as brittle, powdery solids or as gases

Metalloids Elements that form a diagonal separation zone between metals and nonmetals in the periodic table Have properties between those of metals and nonmetals and often exhibit some characteristic properties of each type

Example 3.8 - Distinguishing Electrons and Element Classifications Use the periodic table and Figures 3.9 and 3.10 to determine the following for Ca, Fe, S, and Kr Type of distinguishing electron Classification based on Figure 3.10

Example 3.8 - Solution On the basis of the location of each element in Figure 3.9, the distinguishing electrons are of the following types: Ca: s Fe: d S: p Kr: p Classifications based on figure 3.10 are: Ca - Representative element Fe - Transition element S - Representative element Kr - Noble gas