Chapter 6 Sections 6.1 – 6.5, 6.7, 6.8
The Periodic Table
 Elements were originally grouped into the table into families based on their properties and
reactivities.
 So, elements in Group IA have similar properties.
 We now know that elements in a group have similar electron configurations.
 Vertical columns: groups or families.
 Horizontal rows: periods or series.
Periodic Table Group Names
 Group IA (Li, Na, K...)
 Alkali Metals
 Group IIA (Be, Mg, Ca...)
 Alkali Earth Metals
 Group VIIA (F, Cl, Br..)
 Halogens
 Group VIIIA (He, Ne, Ar...)
 Noble Gases
 The “d” block elements
 Transition Metals
Valence Electrons
 The electrons in highest energy level of an atom are also the farthest away from the nucleus.
 These electrons are called valence electrons.
 They are said to be in the valence shell.
 Valence electrons are involved in making chemical bonds to other atoms.
 Much more later!
Valence and Electron Config.
 The electrons with the highest energy level numbers of an atom make up its valence electrons.
 The electron configuration tells us how many electrons are in the valence shell.
 Be, 4 e –, 1s22s2, has 2 valence e –.
 O, 1s22s22p4, has 6 valence e –.
 Zn, 1s22s22p63s23p64s23d10 has 2.
Periodic Table
 Remember, elements in the same group (column) on the table have similar electron configurations.
 Example: All the elements in Group IA have configurations that end in s1.
 Question: how many valence electrons do each of the elements in Group 1A have?
Valence Electrons and the Periodic Table
 Elements in a column on the periodic table have the same number of valence electrons.
 Group IIA elements have 2 valence electrons.
 Group VIIIA elements have 8 valence electrons
 Except He with 2 electrons total.
 Group VIIA elements have 7 valence electrons.
 etc.
 Most transition metals have 2 valence electrons.
Dot Structures
 We can represent an element’s valence electrons using a dot structure.
 Write the symbol for the element.
 Represent the number of valence electrons as dots around the symbol.
 Pair up dots as needed.
Valence Electrons and Bonding
 Atoms may lose, gain or share electrons in order to get a more stable valence electron configuration.
 CHEMISTRY!!!!
 A configuration with the s and p sublevels filled is especially stable.
 This usually means atoms try to have 8 valence electrons. (Question: why 8?)
 This is the octet rule.
 Atom with small numbers of total electrons (H, He, Li, Be) are “happy” with 2 valence electrons.
 The first energy level only holds 2 electrons.
Ions
 Remember: atoms can lose or gain electrons to form ions.
 The number of electrons lost or gained depends on the number of valence electrons an atom has.
 An atom with 7 valence electrons will gain one electron to have 8.
 An atom with 1 valence electron will lose that outermost electron. The next filled shell below
then becomes the filled valence shell.
Ions and the Periodic Table
 Group IA, 1 (alkali metals):
 (H), Li, Na, K...
 1 valence electron.
 Will tend to lose that electron to have a filled outermost (valence) shell.
 Will have an ionic charge of +1.
 Group IIA, 2 (alkali earth metals):
 Be, Mg, Ca, Sr....
 2 valence electrons.
 Will tend to lose those electrons to have a filled outermost (valence) shell.
 Will have an ionic charge of +2.
More Ions
 Group VIIIA, 18 (Noble gases):
 He, Ne, Ar, Kr...
 8 valence electrons (except He with 2).
 Will not tend to lose or gain electrons; valence shell is already full.
 Will not tend to form ions.
 Group VIIA, 17 (Halogens):
 F, Cl, Br, I...
 7 valence electrons.
 Will tend to gain 1 electron to have a filled valence shell..
 Will have an ionic charge of -1.
Transition Metals
 The electron distributions of the transition metal elements are more complicated.
 Most have configurations of the type s2dx.
 The s sublevel is in the valence shell.
 The d is one shell lower.
 Most transitions metals have 2 valence e-.
 Transition metals have more than one possible ionic charge.
 Fe can be Fe2+ or Fe3+.
Question
 What is the most likely charge of an oxygen ion?
 Why?
Electron Configuration (EC) of Ions
 When writing the EC of an ion, first do the EC of the neutral atom.
 If its a positive ion, subtract the correct number of valence e’s.
 If its a negative ion, add the e’s into the valence shell.
Examples:
Na+
Na: 1s22s22p63s1 or [Ne]3s1
Na+: 1s22s22p6 or [Ne]
O2O: 1s22s22p4 or [He] 2s22p4
O2-: 1s22s22p6 or [Ne]
 Ions or atoms with the same EC are isoelectronic.

Why are many ions isoelectronic with Noble gases?
Example: Transition Metal Ion
Fe2+ and Fe3+
Fe: [Ar]4s23d6
Fe2+: [Ar] 3d6 (loses its valence e’s first)
Fe3+: [Ar] 3d5 (loses its val e’s then a 3d e)
Summary: Ions
 Atoms tend to form ions based on their number of valence electrons.
 Metals tend to form + ions (cations).
 Non-metals tend to form –ions (anions).
 Noble gases have filled valence shells and tend not to form ions.
 We can determine the number of valence electrons in an atom from its group number.
Atomic Radius
 Atoms get smaller as you move across a period on the table from left to right.
 Electrons are being added to the same energy level.
 The number of protons in the nucleus is also increasing.
 The added + charge pulls the electrons in closer.
 Atoms get larger as you move down a group on the table.
 Another electron shell is added as you go down.
 The number of protons are also increased, but their effect on the outermost electrons is shielded
by the inner shells.
Ionization Energy
 The energy required to completely remove an electron from an atom is its ionization energy (I.E.).
 1st I.E.: energy to remove 1st electron
 Makes a +1 charge.
 2nd I.E.: energy to remove 2nd electron
 Makes a +2 charge.
Trends in Ionization Energy
 Metals have lower I.E. than non-metals.
 Metals tend to lose electrons and form cations.
 Non-metals tend to gain electrons and form anions.
 Noble gases have very high ionization energies.
 They have a very stable electron configuration.
 As you go down a group on the table, the I.E. decreases.
 The larger the atom, the easier it is to remove its electron.
 Why?
Homework
 Exercises p. 159 – 162
 9, 10, 17, 29 – 36, 57 – 64, 67 – 70, 75 - 80