Ch. 4: Periodic Properties of the Elements Dr. Namphol Sinkaset Chem 200: General Chemistry I

I. Chapter Outline I.Introduction II.The Periodic Table III.Electrons in the Atom IV.Electron Spin 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 Quantum 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 atomic number.

II. The Modern Periodic Table

II. Major Divisions of the Table Main-group 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 hydrogen atom Since hydrogen has only one e-, the ground state can be written as an electron configuration:

III. Many e- Atoms 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 Magnetic Field

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 quantum number (m s ) 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, m l =0, m s =1/2 He has two p+, so it needs two e-:  1 st e-: n=1, l=0, m l =0, m s =1/2  2 nd e-: n=1, l=0, m l =0, m s =-1/2 The orbital is filled and the e- have paired spins.

IV. Electrons in Helium

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 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, Z eff, 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 energy level 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 2 nd p orbital. We follow Hund’s rule.  When filling degenerate orbitals, electrons fill singly first with parallel spins. Hund’s rule leads to lower energy.

V. Examples

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

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] 3s 2 3p 5. 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, m l, m s }?  {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 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 period. 4)Each vertical column is a group or family. 5)Main group elements are in groups 1,2 and (s and p blocks).

VI. Important Parts of the Periodic Table 6)Transition elements are in groups 3-12 (d block). 7)Inner-transition elements at the bottom (lanthanides and actinides, f block). 8)Staircase line separates metals on L from nonmetals on R. Metalloids 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. Valence 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 Ions Metals tend to lose e-’s and nonmetals tend to gain e-’s. Main-group ions can be predicted.

VI. Transition Metal Cations When forming transition metal cations, remove e-’s from highest n-value orbital first!  V: [Ar] 4s 2 3d 3  V 2+ : [Ar] 4s 0 3d 3

VI. Magnetic Properties Some metals exhibit magnetism  paramagnetic: atom or ion that has unpaired e-’s  diamagnetic: atom or ion in which all e-’s are paired

VI. Sample Problem Draw condensed orbital diagrams for the following and determine whether they are diamagnetic or paramagnetic.  Sc 3+  Ir 2+  Mn 4+

VII. Atomic Radii

VII. Trend in Atomic Radii

Trend in Atomic Radii

VII. Trend in Ion Size Why?

VII. Trend in Ionization Energy ionization energy: energy in kJ needed to remove an e- from gaseous atoms/ions Why? What about 1 st, 2 nd, 3 rd, 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