Slide 1

Chapters 15/6 Ionic Bonding

•15.1 Objectives

–Use the periodic table to infer the numberof valence electrons in an atom and draw itselectron dot (Lewis dot) structure.

–Describe formation of cations from metalsand anions from nonmetals

California StandardsCalifornia Standards

 1d. Students know how to use the periodic table todetermine the number of electrons available forbonding. 1d. Students know how to use the periodic table todetermine the number of electrons available forbonding.

 2e. Students know how to draw Lewis dot structures. 2e. Students know how to draw Lewis dot structures.

California StandardsCalifornia Standards

 2e. Students know how to draw Lewis dot structures. 2e. Students know how to draw Lewis dot structures.

ValenceElectrons:ValenceElectrons:

ELECTRONSAVAILABLEFORBONDINGELECTRONSAVAILABLEFORBONDING

(the red ones)(the red ones)

Valence ElectronsValence Electrons

•Valence electrons are electrons in theoutmost shell (energy level). They are theelectrons available for bonding.

•The number of valence electrons largelydetermines the chemical properties of thatelement.

•For Groups 1A-7A, the number 1-7 is thenumber of valence electrons for that atom.

•Group 0 is an exception – you can think of itas group 8A because all the noble gases(except He) have 8 valence electrons.

Group 1 (alkali metals) have 1valence electron

Group 2 (alkaline earth metals)have 2 valence electrons

Group 13 elements have 3valence electrons

Group 14 elements have 4valence electrons

Group 15 elements have 5valence electrons

Group 16 elements have 6valence electrons

Group 17 (halogens) have 7valence electrons

Group 18 (Noble gases) have 8valence electrons, excepthelium, which has only 2

Transition metals (“d” block)have 1 or 2 valence e- . Why?

Lanthanides and actinides

(“f” block) have 1 or 2 valenceelectrons

Lanthanides and actinides

(“f” block) have 1 or 2 valenceelectrons

Valence Electrons

•Valence electrons are usually the only e-used to bond to other atoms.

–Therefore you usually only show the valencee- in electron dot structures.

–Electron dot structures are diagrams thatshow valence e- as dots.

Generic Dot NotationGeneric Dot Notation

An atom’s valence electrons can be representedby electron dot (AKA Lewis dot) notations.

1 valence e-1 valence e-

2 valence e-2 valence e-

3 valence e-3 valence e-

4 valence e-4 valence e-

5 valence e-5 valence e-

6 valence e-6 valence e-

7 valence e-7 valence e-

8 valence e-8 valence e-

Dot Notations – Period 2Dot Notations – Period 2

Lewis dot notations for the valence electrons ofthe elements of Period 2.

lithiumlithium

berylliumberyllium

boronboron

carboncarbon

nitrogennitrogen

oxygenoxygen

fluorinefluorine

neonneon

Electron Dot Structures

•Note how you draw two dots per side x four sides = 8dots maximum.

•Note how each side gets one before any side gets two.

•See how the number of dots is the same for eachelement within a group (column).

Octet Rule

•The Octet Rule was created by Gilbert Lewis in1916.

•That’s why these diagrams are sometimes calledLewis dot structures.

•In forming compounds, atoms tend to achieve thee- configuration of a noble gas, 8 valence e-.

•An octet is a set of 8.

•Each noble gas (except He) has 8 valenceelectrons in their highest principle energy level,and the general configuration is ns2np6

(like 2s22p6 or 3s23p6)

Metallic vs. Nonmetallic Elements

•Atoms of the metallic elements (includingcolumn 1A and 2A) tend to lose their outershell valence e- so they can have a completeoctet at the next energy level down.

•Atoms of nonmetallic elements tend to gain e-(steal e-) or share e- with another nonmetallicelement to achieve their complete octet.

•There are exceptions but the octet ruleusually applies to most atoms in compounds.

Cations and Anions

•If an atom loses a valence e- = cation

•If an atom gains a valence e- = anion

•Metals create cations because they start with 1 to 3e- and usually get all of the valence e- stolen so theycan get down to a full lower level octet.

•Example: Sodium loses 1 e-

•Before Na 1s2 2s2 2p6 3s1

•After Na+ 1s2 2s2 2p6

–(note the 8 e- in the n=2 shell)

•Like Ne 1s2 2s2 2p6

– (Neon has 8 e- in the n=2 shell)

•The change is written as follows:

–Na· Na+ + e-

Cations

•Cations of group 1A alkali metals +1

•Cations of group 2A alkali metals +2

·Mg· Mg2+ + 2e-

•For transition metals, the charges on thecations may vary. Note the Roman Numeral.

•Example: Fe has two: iron(II) or Fe2+

iron(III) or Fe3+

•Some atoms formed by transition metals donot have noble-gas electron configurationsand are therefore exceptions to the octetrule.

Exceptions

•Example: Ag Silver

•1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s1 4d10 (oddball)

•Silver would have to lose 11 electrons to get downto noble gas Krypton’s configuration. To gainenough e- to get to Xenon’s configuration, itwould have to gain 7 electrons. Neither one islikely.

•But if Ag loses its one 5s1 electron, then it has anouter shell with 18 e- (the 4 shell), which is a fullshell, and relatively favorable.

•Therefore Ag always forms the Ag+ cation.

Anions

•Anions are atoms or groups with a negativecharge (extra electrons).

•Atoms of nonmetallic elements have relativelyfull valence shells and are looking to steal e-to make their shells full.

•Cl 1s2 2s2 2p6 3s2 3p5 neutral atom

•Cl- 1s2 2s2 2p6 3s2 3p6 anion

•Ar 1s2 2s2 2p6 3s2 3p6 now Cl- has Ar config.

Section 15.2 Ionic Bonding

California Standards

Students know atoms combine to formmolecules by sharing electrons to formcovalent or metallic bonds or byexchanging electrons to form ionic bonds.

 Students know salt crystals, such asNaCl, are repeating patterns of positiveand negative ions held together byelectrostatic attraction.

Ionic vs. Covalent Bonds

•Bonds: Forces that hold groups of atoms

together and make them function

as a unit.

 Ionic bonds – transfer of electrons Ionic bonds – transfer of electrons

 Covalent bonds – sharing of electrons Covalent bonds – sharing of electrons

(this will be Ch. 16) (this will be Ch. 16)

Ionic Bonding

Na: 1s22s22p63s1 now Na+ 1s22s22p6

Cl: 1s22s22p63s23p5 now Cl- 1s22s22p63s23p6

Aluminum has three valencee- to steal, and the Bromineatoms would each like tosteal one e-.

So the Aluminum atom givesup three electrons and theBromine atoms each receiveone.

Examples of Ionic Compounds

•Mg2+Cl21-

•Magnesium chloride: Magnesium losestwo electrons and each chlorine gainsone electron

•Al23+S32-

•Aluminum sulfide: Each aluminum losesthree electrons (six total) and eachsulfur gains two electrons (six total)

Recall that anions end in –ide.

Sodium Chloride crystal lattice

•Ionic compounds formsolid crystals at ordinarytemperatures.

•Ionic compounds organizein a characteristic crystallattice of alternatingpositive and negative ions.

All salts are ionic compounds and form crystals.

Properties of Ionic Compounds

Two K atoms lose 1 e- each => One O atom gains 2 e-

3 Mg atoms lose x 2 e- each => 2 N atoms gain 3 e- each

Ch. 6 – Ionic Naming

The Laws of Definite and Multiple Proportions

•The law of Definite Proportions states thatin samples of any chemical compound, themasses of the elements are always in thesame proportions.

•The law of Multiple Proportions states thatwhenever two elements form more than onecompound, the different masses of oneelement that combine with the same mass ofthe other element are in the ratio of smallwhole numbers.