Chemical Bonding

 Attractive forces that hold atoms together in compounds. The electrons involved in bonding are usually those in the outermost (valence) shell.  Most elements in compounds want to gain noble gas configuration. They will do so by either losing or gaining electrons (ionic compounds) or by sharing electrons (covalent compounds)

Chemical bonds are classified into two types: ¬ Ionic bonding results from electrostatic attractions among ions; which are formed by the transfer of one or more electrons from one atom to another. (metals low χ with nonmetals high χ) ¬ Covalent bonding results from sharing one or more electron pairs between two atoms. (nonmetals only similar χ )

IonicCovalent  Melting Pt  Solubility ◦ (polar solvents)  Solubility ◦ (nonpolar solvents)  Conductivity ◦ (molten & aqueous solutions) High Low Soluble Insoluble Insoluble Soluble High Low

2 extremes in bonding pure covalent bonds ◦ electrons equally shared by the atoms pure ionic bonds ◦ electrons are completely lost or gained by one of the atoms most compounds fall somewhere between these two extremes

 # of atoms in the molecule ◦ Monatomic = 1 atomEx. He ◦ Diatomic = 2 atomsEx. O 2 ◦ Triatomic = 3 atomsEx. O 3 ◦ Polyatomic = manyEx. H 2 SO 4 or S 8  Homonuclear: the mlcl is composed of only 1 kind of atom: O 2, H 2, P 4  Heteronuclear: the mlcl is made up of more than 1 kind of atom: H 2 O

or Lewis dot formulas, a convenient bookkeeping method for valence electrons (electrons that are transferred or involved in chemical bonding) Only the electrons in the outermost s and p orbitals are shown as dots.

elements in the same group have same Lewis dot structures For groups IA – VIIIA, the group number equals the # of valence electrons Valence electrons determine the chemical and physical properties of the elements as well as the kinds of bonds they form.

metals react with nonmetals to form ionic compounds cations or positive (+) ions (metals) ◦ atoms have lost 1 or more electrons anions or negative (-) ions (nonmetals) ◦ atoms have gained 1 or more electrons

We can use Lewis formulas to represent the neutral atoms and the ions they form.

underlying reasons for LiF formation 1s 2s 2p Li  F  becomes Li +  [He] F -  [Ne]

Li + ions contain two electrons ◦ same number as helium F - ions contain ten electrons ◦ same number as neon Li + ions are isoelectronic with helium F - ions are isoelectronic with neon Isoelectronic species contain the same number of electrons. cations become isoelectronic with preceding noble gas anions become isoelectronic with following noble gas

IIA metals with VIIA nonmetals, mostly ionic compounds ~ exceptions - BeCl 2, BeBr 2, BeI 2 these are covalent compounds Be(s) + F 2 (g)  BeF 2 (s) electronically this is happening similarly for all of the IIA & VIIA M(s) + X 2  M 2+ X 2 -

IA + VIIA MX IIA + VIIA MX 2 IIIA + VIIA MX 3 IA + VIA M 2 X IIA + VIA MX IIIA + VIA M 2 X 3  NaF  BaCl 2  AlF 3  Na 2 O  BaO  Al 2 S 3

IA + VA M 3 X IIA + VA M 3 X 2 IIIA + VA MX  Na 3 N  Mg 3 P 2  AlN H forms ionic compounds with IA and IIA metals LiH, KH, CaH 2, BaH 2,, etc. other H compounds are covalent

extended three dimensional arrays of oppositely charged ions high melting points because coulomb force is strong

~ ions with high charges F is large ~ ions with small charges F is small arrange these compounds in order of increasing attractions among ions KCl, Al 2 O 3, CaO K + Cl - < Ca 2+ O 2- <Al 2 3+ O 3 2-

covalent bonds formed when atoms share electrons share 2 electrons - single covalent bond share 4 electrons - double covalent bond share 6 electrons - triple covalent bond attraction is electrostatic in nature ◦ lower potential energy when bonded

 Covalent bonding may be explained by 2 different theories ◦ Valence bond (VB) theory: each atom has electrons in atomic orbitals which overlap to form bonds (Ch. 8) ◦ Molecular orbital (MO) theory: the electrons belong to the molecule as a whole and are in molecular orbitals instead of belonging to each atom (Ch. 9)

 The element needing the most electrons to fill its octet is usually the central atom  The most symmetrical skeleton is usually correct  Halogens and H always share one electron to complete outer shell  In ternary acids, H are bonded to O (ternary acids are oxy-acids: they contain H, O, and another nonmetal)

 Carbon always obeys the octet rule  Carbon rarely has lone pairs of electrons. Exception: If it’s at the end of a molecule or ion. Ex. CN -, CO, CNO  When forming multiple bonds between atoms, both atoms donate the same number of electrons

 Oxygen atoms normally bond to other nonmetals, not to each other  Oxygen can do several things depending on the mlcl. ◦ Single bond by sharing an electron ◦ Single bond by accepting 2 electrons from another atom and not sharing at all ◦ Double bonds by sharing 2 of its electrons

homonuclear diatomic molecules ◦ hydrogen, H 2 ◦ fluorine, F 2 ◦ nitrogen, N 2 nonpolar covalent bonds - electrons are shared equally symmetrical charge distribution - must be the same element to share exactly equally

Lewis dot representation H 2 molecule formation

heteronuclear diatomic molecules hydrogen halides ◦ hydrogen fluoride, HF ◦ hydrogen chloride, HCl ◦ hydrogen bromide, HBr

polar covalent bonds - unequally shared electrons assymmetrical charge distribution different electronegativities Some bonds are very polar, Ex. HF

Electron density map of HF ◦ blue areas - low electron density ◦ red areas - high electron density polar molecules have separation of centers of negative and positive charge

Some bonds are only slightly polar, ex. HI

Electron density map of HI ◦ blue areas - low electron density ◦ red areas - high electron density notice that the charge separation is not as big as for HF ◦ HI is only slightly polar

Representative elements achieve noble gas configurations in most of their compounds. Lewis dot formulas are based on the octet rule.  H needs two electrons to have Helium's noble gas configuration, everything else wants 8

water, H 2 O ammonia molecule, NH 3 ammonium ion, NH 4 + hydrogen cyanide, HCN sulfite ion, SO 3 2-

 Two or more Lewis dot diagrams are needed to describe the bonding in a molecule or ion.  LDD for sulfur trioxide, SO 3

three possible structures for SO 3 invoke resonance ◦ Double-headed arrows are used to indicate resonance formulas.

flaw in our representations of molecules no single or double bonds in SO 3 all bonds are the same best picture

 The concept of formal charges helps us choose the correct Lewis structure for a molecule. If a resonance structure has a high formal charge it’s not a very good one.  Formal charge = group # - e - you can assign to that atom Or F.C.= (valence e - ) – (# of bonds + # of unshared e - ) pg 289

Sigma bonds (σ) : result of head-on (end to end overlap, there is a free rotation around σ bonds. Pi bonds (π) : result of side-on overlap of p orbitals. There is no free rotation around a π bond. The side –on overlap locks the molecule into place. All single bonds are sigma bonds: 1σ bond All double bonds: 1 σ bond, 1 π bond All triple bonds: 1 σ bond, 2 π bonds

¹ species in which the central element must have a share of more or less than 8 valence electrons to accommodate all substituents º compounds of the d- and f-transition metals In cases where the octet rule does not apply, the elements attached to the central atom nearly always attain noble gas configurations. ◦ The central atom does not

 Write LDD for BBr 3  Write LDD for AsF 5  Write LDD for XeF 4

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