Chemistry: Atoms First Second Edition Julia Burdge & Jason Overby Copyright (c) The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 4 Periodic Trends of the Elements M. Stacey Thomson Pasco-Hernando State College

Development of the Periodic Table In 1864, John Newlands noted that when the elements were arranged in order of atomic number that every eighth element had similar properties.  He referred to this as the law of octaves. In 1869, Dmitri Mendeleev and Lothar Meyer independently proposed the idea of periodicity. Mendeleev grouped elements (66) according to properties. Mendeleev predicted properties for elements not yet discovered, such as Ga. 4.1

Classification of Elements The main group elements (also called the representative elements) are the elements in Groups 1A through 7A. 4.2

Classification of Elements The noble gases are found in Group 8A and have completely filled p subshells.

The Modern Periodic Table The transition metals are found in Group 1B and 3B through 8B.  Group 2B have filled d subshells and are not transition metals.

The Modern Periodic Table The lanthanides and actinides make up the f-block transition elements.

The Modern Periodic Table There is a distinct pattern to the electron configurations of the elements in a particular group.

The Modern Periodic Table The electrons in the outermost occupied PRINCIPLE ENERY LEVEL of an atom are called the valence electrons. Valence electrons are involved in the formation of chemical bonds. For Group 1A: [noble gas]ns 1 valence core For Group 2A: [noble gas]ns 2 valencecore For Group 7A: [noble gas]ns 2 np 5 valencecore Example: [He]2s 1 valence core Example: [Ar]4s 2 valence core Example: [Ne]3s 2 3p 5 valence core

Without using a periodic table, give the ground-state electron configuration and block designation (s-, p-, d-, or f-block) of an atom with (a) 17 electrons, (b) 37 electrons, and (c) 22 electrons. Classify each atom as a main group element or transition metal. Worked Example 4.2

Effective Nuclear Charge Effective nuclear charge (Z eff ) is the actual magnitude of positive charge that is “experienced” by an electron in the atom. In a multi-electron atom, electrons are simultaneously attracted to the nucleus and repelled by one another.  This results in shielding, where an electron is partially shielded from the positive charge of the nucleus by the other electrons.  Although all electrons shield one another to some extent, the core electrons shield the most.  As a result, the value of Z eff increases steadily from left to right because the core electrons remain the same but Z increases. LiBeBCNOF Z Z eff

Atomic Radius The Atomic radius: the distance from the atom’s nucleus and its valence shell. Atomic radius increases from top to bottom because outermost shell lies farther from the nucleus Atomic radius decreases from left to right because of increasing Z eff which draws the valence shell closer to the nucleus Atomic radii (in picometers) 4.4

Referring only to a periodic table, arrange the elements P, S, and O in order of increasing atomic radius. Worked Example 4.3 Strategy Use effective nuclear charge to compare the atomic radii of two of the three elements at a time. Text Practice:

Ionization Energy Ionization energy (IE) is the minimum energy required to remove an electron from an atom in the gas phase. The result is an ion, a chemical species with a net charge. Sodium has an ionization energy of kJ/mol. Specifically, kJ/mol is the first ionization energy of sodium, IE 1 (Na), which corresponds to the removal of the most loosely held electron. Na(g) → Na + (g) + e −

Ionization Energy

In general, as Z eff increases, ionization energy also increases.  Thus, IE 1 increases from left to right across a period.

Ionization Energy Within a given shell, electrons with a higher value of l are higher in energy and thus, easier to remove.

Ionization Energy Removing a paired electron is easier because of the repulsive forces between two electrons in the same orbital.

Ionization Energy It is possible to remove additional electrons in subsequent ionizations, giving IE 1, IE 2, and so on. IE 1 (Na) = 496 kJ/mol IE 2 (Na) = 4562 kJ/mol Na(g) → Na + (g) + e − Na + (g) → Na 2+ (g) + e −

Ionization Energy It takes more energy to remove the 2nd, 3rd, 4th, etc. electrons because it is harder to remove an electron from a cation than an atom. It takes much more energy to remove core electrons than valence.  Core electrons are closer to nucleus.  Core electrons experience greater Z eff because of fewer filled shells shielding them from the nucleus.

Would you expect Na or Mg to have the greater first ionization energy (IE 1 )? Which should have the greater second ionization energy (IE 2 )? Worked Example 4.4 Strategy Consider effective nuclear charge and electron configuration to compare the ionization energies. Na has one valence electron and Mg has two. Text Practice: 4.50

Electron Affinity Electron affinity (EA) is the energy released when an atom in the gas phase accepts an electron. Cl(g) + e − → Cl − (g)

Electron Affinity Like ionization energy, electron affinity increases from left to right across a period as Z eff increases.  Easier to add an electron as the positive charge of the nucleus increases.

Electron Affinity It is easier to add an electron to an s orbital than to add one to a p orbital with the same principal quantum number.

Electron Affinity Within a p subshell, it is easier to add an electron to an empty orbital than to add one to an orbital that already contains an electron.

For each pair of elements, indicate which one you would expect to have the greater first electron affinity, EA 1 : (a) Al or Si. Worked Example 4.5 Strategy Consider effective nuclear charge and electron configuration to compare the ionization energies. (a) Al is in Group 3A and Si is in Group 4A. Al has three valence electrons ([Ne]3s 2 3p 1 ), and Si has four valence electrons ([Ne]3s 2 3p 2 ). Text Practice: 4.58

Study Guide for Sections DAY 9, Terms to know: Sections valence electrons, core electrons, effective nuclear charge (Z eff ), shielding, atomic radii, ionization energy, electron affinity DAY 9, Specific outcomes and skills that may be tested on exam 1: Sections Be able to give a complete or abbreviated electron configuration for an atom in either its ground state or one possible excited state Be able to give a complete or abbreviated orbital diagram for an atom either its ground state or one possible excited state Be able to describe what effective nuclear charge is and how it is calculated Be able to rank relative atomic radii, electron affinity, ionization energy ionic radii, and explain WHY they are ranked based on attractions and repulsions within the atom Be able to rank atoms in order of greatest ionization energies including IE 1, IE 2, IE 3, etc. and explain WHY they should be ranked in that order

Extra Practice Problems for Sections Complete these problems outside of class until you are confident you have learned the SKILLS in this section outlined on the study guide and we will review some of them next class period

Prep for Day 10 Must Watch videos: (Tyler DeWitt: ionic bonds part 1) (Tyler DeWitt: ionic bonds part 2) (Tyler DeWitt: ionic bonds part 3) (Tyler DeWitt: writing ionic formulas) Other helpful videos: k32U&list=PLqOZ6FD_RQ7kTjN4O2MNzf5YfeiIx7SGIhttps:// k32U&list=PLqOZ6FD_RQ7kTjN4O2MNzf5YfeiIx7SGI (UC-Irvine watch first 15 minutes) shhttps:// sh (UC-Irvine ions and molecules) Read Sections ,