CHAPTER 3: Chemical Periodicity and the Formation of Simple Compounds

•Groups of Elements •Multiple bonds •The •Formal Charges • • •Core and •Octet Rule •VSEPR Theory •Lewis dot structures •Elements forming more •Ionic and covalent than one ion bonds •Names of Ions

CHEM 1310 A/B Fall 2006 Groups of Elements

• Most obvious groupings: - (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/ , , and • Eight “main” groups of elements (neglects transition elements)

CHEM 1310 A/B Fall 2006 Eight Main Groups

• I. Alkali metals. (Li), (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 across membranes in biochemistry • II. Alkaline earth metals. (Be), (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. . 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. . 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

Cl 2 Br2 I2

CHEM 1310 A/B Fall 2006 Main Groups, Cont’d

• VIII. Noble gases. (He), (Ne), (Ar), krypton (Kr), xenon (Xe), (Rn) • Very unreactive • All are gases • “Less distinct” groups: III, IV, V. Contain mixtures of metals, semimetals, .

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 ’s tendency to gain electrons.

• Electronegativity usually increases going left to right across a , 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:

2- 2- •O2 peroxide •SO3 sulfite - 2- •O2 superoxide •CO3 - 3- • HCO3 hydrogen carbonate •PO4 phosphate (bicarbonate) 2- •HPO4 hydrogen phosphate •HSO- hydrogen 4- 4 •H2PO dihydrogen phosphate (bisulfate) 4- •SiO4 silicate - •OHhydroxide • CNO- cyanate - •CNcyanide •SCN- thiocyanate - •NO3 nitrate - •ClO4 perchlorate - •NO2 nitrite 2- •CrO4 chromate •SO2- sulfate 4 • … and more! 2- •Cr2O7 dichromate 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 with similar . 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 made of 2 shared electrons is depicted by a short line –

CHEM 1310 A/B Fall 2006 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 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 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 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., )

• 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 repulsion theory – predict 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 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 bipyramidal and octahedral complexes • Replace the “equatorial” atoms first

CHEM 1310 A/B Fall 2006 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 ______

CHEM 1310 A/B Fall 2006