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Home , Electron configuration

Ch. 3: Periodic Properties of the Elements

Dr. Namphol Sinkaset

Chem 200: General Chemistry I I. Chapter Outline

I. Introduction II. The III. in the IV. V. Sublevel Energy Splitting VI. Using the Periodic Table VII. Periodic Properties and Trends I. Organizing Chemical Info

• When information of the elements was organized, chemistry began to advance quickly. • Element “triads” and “octaves” • Mendeleev’s periodic table in 1869 • mechanics explains why the periodic table appears as it does. II. Periodic Law

• Initially, Mendeleev ordered elements by increasing atomic mass. • Later work by Moseley showed that they should be ordered by . II. The Modern Periodic Table II. Major Divisions of the Table

• Main- elements have properties that are largely predictable based on their location. • Transition and inner-transition elements have properties that are less predictable based on their location. • Each column within the main group region is known as a family or group. III. Electrons Occupying Orbitals

• From Chapter 3, we know how orbitals are ordered for the atom • Since hydrogen has only one e-, the can be written as an : III. Many e-

• The Schrödinger equation can’t solve multi-e- atoms; we only get approximate solutions. • We use quantum #’s from H atom solution to describe orbitals of other atoms. III. New Considerations

• An atom with more than 1 e- is more complicated. • Two more concepts are needed to understand these larger atoms: 1) Electron spin 2) Sublevel energy splitting IV. H Atoms in a IV. e- Spin

• e- generate a small magnetic field as if they were spinning. • There are two possible directions e- can spin, so there are two possible states.

• spin (ms) can be either +1/2 or –1/2. IV. Representing e- Spin

• Orbital diagrams are used to show electron occupation and spin. IV. Pauli Exclusion Principle

• No two e- in the same atom can have the same 4 quantum #’s!!

• H: n=1, l=0, ml=0, ms=1/2 • He has two p+, so it needs two e-: st . 1 e-: n=1, l=0, ml=0, ms=1/2 nd . 2 e-: n=1, l=0, ml=0, ms=-1/2 • The orbital is filled and the e- have paired spins. IV. Electrons in V. H vs. He Energy Levels

• One additional e- complicates the He spectrum greater than expected. Why? V. Removal of Degeneracy

• In H atom, energy of an orbital depends only on n. . e.g. Energies of 3s, 3p, 3d are degenerate. • In every other atom, this is not true. . E (s orbital) < E (p orbital) < E (d orbital) < E (f orbital), etc. • What removes the degeneracy? V. Sublevel Energy Splitting

• Three factors contribute to differing sublevel energies: 1) Coulomb’s Law (Z) 2) shielding 3) penetration V. Coulomb’s Law

1 q q E  1 2 4 0 r

• The PE of like charges is positive (unstable), but decreases as they move apart. • The PE of unlike charges is negative (stable) and increases as they get closer. • The magnitude of the interaction increases as charges on particles increases. V. Nuclear Charge

• p+ in nucleus constantly pull all e-. • Higher charges attract more strongly. • More p+ lowers orbital E by increasing e-/nucleus attraction. V. Shielding

• Electrons shield each other from the full charge of the nucleus. • The effective nuclear

charge, Zeff, is the actual positive charge an e- feels. V. Penetration

• The movement of an outer e- into the region occupied by inner e- is called penetration. • Penetrating e- experience higher nuclear charge, lowering its PE. V. 2s and 2p Radial Distribution V. 3s, 3p, 3d Penetration

• This is the reason why energetically, s < p < d. V. Order of Sublevels V. The Aufbau Principle

• Since e- are “lazy,” they want to “occupy” the lowest possible. • Thus, if we know the energy order of sublevels, then we can “build up” the e- configurations of each atom. V. Writing e- “in” Orbitals

• Two ways to represent how e- are situated in atoms: 1) e- configuration, nl# 2) orbital diagram, which uses arrows indicating e-’s and their spin V. Hund’s Rule

• In the orbital diagram of C, there was a choice as to where to place the 2nd p orbital. • We follow Hund’s rule. . When filling degenerate orbitals, electrons fill singly first with parallel spins. • Hund’s rule to lower energy. V. Examples VI. The Periodic Table

• As you go left to right on the periodic table, you are using the . VI. The Periodic Table

• Each region of the periodic table indicates what orbitals are being “filled.” VI. Using the Periodic Table • You can use an element’s location to write its full or condensed electron configuration/orbital diagram. VI. Using the Periodic Table

• Therefore, Cl is: [Ne] 3s2 3p5. • From the orbital diagram, we can write specific quantum numbers for each e-. • Which e-’s are identified with the

following quantum #’s {n, l, ml, ms}? . {3, 0, 0, -1/2} . {3, 1, 1, 1/2} VI. Some Caveats

• Because energy differences between s and d are small, some exceptions to how e-’s fill exist. . Same for d and f. • Remember that d principal quantum # lags by one. • Remember that f principal quantum # lags by two. VI. Sample Problem 3.1

• Write condensed electron configurations and orbital diagrams for the following elements. . Mn . Sb . Nd VI. The Periodic Table VI. Important Parts of the Periodic Table

1) Each element placed in box w/ atomic #, atomic mass, and atomic symbol. 2) Atomic # increases as go L to R. 3) Each horizontal row is . 4) Each vertical column is a group or family. 5) Main group elements are in groups 1,2 and 13-18 (s and p blocks). VI. Important Parts of the Periodic Table

6) Transition elements are in groups 3-12 (d ). 7) Inner-transition elements at the bottom ( and , f block). 8) Staircase line separates metals on L from on R. or semimetals lie adjacent to the line. 9) Some groups have special names: alkali metals, alkali earth metals, halogens, noble gases. VI. Types of Elements VI. Core vs. e-’s VI. Valence Electrons

• valence electrons: the outermost e- in an atom • Valence e- determine an atom’s chemistry; thus, atoms in the same vertical column have similar chemical properties. • Valence e- can be determined from the Group number. VI. Formation of

• Metals tend to lose e-’s and nonmetals tend to gain e-’s. • Main-group ions can be predicted. VI. Cations

• When forming transition metal cations, remove e-’s from highest n-value orbital first! . V: [Ar] 4s2 3d3 . V2+: [Ar] 4s0 3d3 VI. Magnetic Properties

• Some metals exhibit . paramagnetic: atom or that has unpaired e-’s . diamagnetic: atom or ion in which all e-’s are paired VI. Sample Problem 3.2

• Draw condensed orbital diagrams for the following and determine whether they are diamagnetic or paramagnetic. . Sc3+ . Ir2+ . Mn4+ VII. Atomic Radii VII. Trend in Atomic Radii Trend in Atomic Radii VII. Trend in Ion Size

• Why? VII. Trend in

• ionization energy: energy in kJ needed to remove an e- from gaseous atoms/ions • Why? • What about 1st, 2nd, 3rd, ionization energies? VII. Successive IE’s VII. Electron Affinity

• electron affinity: energy change in kJ when e- added to a gaseous atom/ion (generally negative) • Why? VII. Trend in Metallic Character