CHEM 1310 A/B Fall 2006

CHAPTER 3: Chemical Periodicity and the

Formation of Simple Compounds•Groups of Elements•The Periodic Table•Electronegativity•Core and valence electrons•Lewis dot structures•Ionic and covalent bonds•Names of Ions

•Multiple bonds•Formal Charges•Resonance•Octet Rule•VSEPR Theory•Elements forming more than one ion

CHEM 1310 A/B Fall 2006

Groups of Elements• Most obvious groupings:

- metals (shiny, look metallic, conduct heat and electricity)- non-metals (don’t have above properties)- semi-metals (some metallic, some non-metallic properties)

• Less obvious groupings:Base grouping on chemical properties, esp. the empirical formulas of their binary compounds w/ chlorine, oxygen, and hydrogen

• Eight “main” groups of elements (neglects transition elements)

CHEM 1310 A/B Fall 2006

Eight Main Groups• I. Alkali metals. Lithium (Li), sodium (Na),

potassium (K), rubidium (Rb), cesium (Cs), francium (Fr)– Form 1:1 binary compounds with chlorine– React with H2O to give off H2– Dissolved Na & Cl important in transport of molecules

across membranes in biochemistry• II. Alkaline earth metals. Beryllium (Be),

magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), radium (Ra)– Form 1:2 compounds with chlorine– Form 1:1 compounds with oxygen

CHEM 1310 A/B Fall 2006

Main Groups• VI. Chalcogens. Oxygen (O), sulfur (S),

selenium (Se), tellurium (Te)– Form 1:1 compounds with alkaline earth metals, e.g.,

CaO– Form 2:1 compounds with alkali metals, e.g., Li2O

• VII. Halogens. Fluorine (F), chlorine (Cl), bromine (Br), iodine (I)– Form 1:1 binary compounds with alkali metals– F has somewhat different properties than the others– Different physical properties (F, Cl are gases F2 and

Cl2; Br2 is liquid, I2 is solid at room temp)

CHEM 1310 A/B Fall 2006

Different physical properties of the chemically similar halogens

Cl2 Br2 I2

CHEM 1310 A/B Fall 2006

Main Groups, Cont’d

• VIII. Noble gases. Helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), radon (Rn)

• Very unreactive• All are gases• “Less distinct” groups: III, IV, V. Contain

mixtures of metals, semimetals, nonmetals.

CHEM 1310 A/B Fall 2006

Allotropes

CHEM 1310 A/B Fall 2006

Periodic Table

• Surprisingly, if we line up the elements according to their masses, the “columns” of the table all have similar chemical properties. Why?

CHEM 1310 A/B Fall 2006

C. Seife, Science 293, 777 (2001)

CHEM 1310 A/B Fall 2006

Electronegativity• Metals give up electrons easily.

“Electropositive.”

• Non-metals prefer to gain electrons. Electronegativity: a measure of the atom’s tendency to gain electrons.

• Electronegativity usually increases going left to right across a period, and decreases going down a column. Exception: noble gases not electronegative.

CHEM 1310 A/B Fall 2006

Core & valence electrons

• Periodicity depends on “valence” electrons• Core electrons

• Valence electrons

CHEM 1310 A/B Fall 2006

Lewis dot symbols• Consider only valence electrons• Electrons represented by dots• G.N. Lewis, 1916 (before QM!)• Put one e- on each side until 4, then go back around and start

adding a second e- on each side until run out of electrons• 8 electrons are “closed shell” – all e- paired up, no empty spaces

CHEM 1310 A/B Fall 2006

Lewis dot symbols

• Number of unpaired electrons tells something about reactivity and bonding

• Noble gases are filled with paired electrons; don’t want to react

• For ions, just add/subtract the appropriate number of electrons. E.g., F- gives …

CHEM 1310 A/B Fall 2006

Octets

• Octet: 4 pairs of electrons.• Having an octet (and no unpaired

electrons) makes noble gases “happy”• Ions of elements tend to form octets

CHEM 1310 A/B Fall 2006

Ionic Bonds

• Ionic bonds form due to the Coulomb attraction between cations and anions

• Ionic bonds can be fairly strong, and the attraction between two opposite charges persists to long distances

CHEM 1310 A/B Fall 2006

Octet and ionic compounds

• Use the octet rule to predict empirical formulas for these ionic compounds

– K and Br

– Mg and Cl

– Na and O

CHEM 1310 A/B Fall 2006

Names of monoatomic ions

• Cations: Named same as the element, plus “ion”

• Anions: Add “-ide” suffix

CHEM 1310 A/B Fall 2006

Polyatomic ions• See Table 3-5. Some common ones:• O2

2- peroxide• O2

- superoxide• HCO3

- hydrogen carbonate (bicarbonate)

• HSO4- hydrogen sulfate

(bisulfate)• OH- hydroxide• CN- cyanide• NO3

- nitrate• NO2

- nitrite• SO4

2- sulfate• Cr2O7

2- dichromate

• SO32- sulfite

• CO32- carbonate

• PO43- phosphate

• HPO42- hydrogen phosphate

• H2PO4- dihydrogen phosphate• SiO4

4- silicate• CNO- cyanate• SCN- thiocyanate• ClO4

- perchlorate• CrO4

2- chromate• … and more!

CHEM 1310 A/B Fall 2006

Polyatomic example

• Just as for monoatomic ions, compounds including polyatomic ions should be electrically neutral.

• What’s the empirical formula of potassium sulfate?

CHEM 1310 A/B Fall 2006

Covalent bonds

• In ionic compounds, there is a transfer of one or more electrons from one unit to another

• In covalent bonds, electrons are shared. Covalent bonds more likely between atoms with similar electronegativities. Groups III-V more likely for covalent bonds (hard to get ions with |charge| > 2)

CHEM 1310 A/B Fall 2006

Lewis structures and covalent compounds

• As for ionic compounds, try to make atoms “happy” by giving them 8 electrons (octet)

• In covalent compounds, achieve octet by sharing electrons

• H is a special case, it likes 2 e-

CHEM 1310 A/B Fall 2006

Depicting bonds

• A covalent bond made of 2 shared electrons is depicted by a short line –

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Multiple bonds

CHEM 1310 A/B Fall 2006

Multiple bonds

• Bond strength: single < double < triple• Bond length: single > double > triple

CHEM 1310 A/B Fall 2006

Formal Charges

• Formal charge = (# of valence e-) - (# of e- in lone pairs)- ½ (# of e- in bonds)

• Help distinguish between “better” or “worse” Lewis structures (smaller or no formal charges is better)

• (Formal charges don’t really mean that the atoms have that charge; might have a partial charge)

CHEM 1310 A/B Fall 2006

Tips on Lewis structures• See book for a more complete list• Remember to add/subtract e- if molecule is charged• You have to know the molecular “skeleton” or

connectivity to get started• Connect all bonded atoms by at least 1 bond• Share additional electrons as necessary (forming double

or triple bonds) to try to achieve octets• Unshared electron pairs are “lone pairs”• Compute formal charges to see how “good” structure

seems (minimize formal charges)• Read section 3-5 carefully if you have problems

CHEM 1310 A/B Fall 2006

Lewis structure examples

• H2O

• H2CSO

CHEM 1310 A/B Fall 2006

Resonance

• Most famous example is benzene…

Two equally good Lewis structures. True molecule is best represented by both structures simultaneously

CHEM 1310 A/B Fall 2006

Exceptions to the rules…

• Odd-electron molecules (radicals)

• “Octet deficient” molecules (e.g., boron)

• Valence-shell expansion. When a central atom is S, P, I, Xe (some others), can have > 8 e-

CHEM 1310 A/B Fall 2006

VSEPR Theory: Shapes of Molecules

• Valence-shell electron pair repulsion theory – predict molecular geometry from Lewis structure.

• Guess geometry around each atom based on # of bonds and lone pairs

• Basic idea: electron pairs repel each other, stay as far apart as geometrically possible. Steric # : # of bonded atoms + # of lone pairs

CHEM 1310 A/B Fall 2006

VSEPR examples

CHEM 1310 A/B Fall 2006

More VSEPR

• It’s just a question of how many atoms are around each other atom

• Only complication: sometimes a lone pair instead of a bonded atom. Tweaks geometry a bit.

• Repulsion forces: lone pair – lone pair > lone pair – bonding pair > bonding pair – bonding pair

CHEM 1310 A/B Fall 2006

Bond angles in CH4 vs NH3

CHEM 1310 A/B Fall 2006

Lone pairs in trigonal bipyramidaland octahedral complexes

• Replace the “equatorial” atoms first

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Two centers

• C2H6 and C2H4

CHEM 1310 A/B Fall 2006

Dipole moments• A vector quantity --- gives response of molecule

to electric field• Each bond contributes a bond dipole vector

pointing in the direction of the less electronegative atom

• Total dipole moment is a vector sum of bond dipoles

CHEM 1310 A/B Fall 2006

Elements forming more than one ion

• See section 3-8• Metals late in groups III, IV, V and transition

metals often form more than one stable ion– Cu+ copper (I) “cuprous” ion– Cu2+ copper (II) “cupric” ion– Fe2+ iron (II) “ferrous” ion– Fe3+ iron (III) “ferric” ion

• Iron (III) sulfate is ____________• Iron (II) sulfate is ____________