Bonding : General Concepts Chapter 8

Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that participate in chemical bonding. 1A 1ns 1 2A 2ns 2 3A 3ns 2 np 1 4A 4ns 2 np 2 5A 5ns 2 np 3 6A 6ns 2 np 4 7A 7ns 2 np 5 Group# of valence e - e - configuration

Lewis Electron-Dot Symbols A Lewis electron-dot symbol is a symbol in which the electrons in the valence shell of an atom or ion are represented by dots placed around the letter symbol of the element. Note that the group number indicates the number of valence electrons. Na..... Si... : P :. :. S Mg... Al.. : : Cl. : Ar : : : : Group IGroup IIGroup VIIGroup VIII Group VI Group IVGroup VGroup III

Lewis symbol (Lewis structure) - a way to represent atoms (and their bonds) using the element symbol and valence electrons as dots.

Chemical Bonding Chemical Bond - the force of attraction between any two atoms in a compound. Interactions involving valence electrons are responsible for the chemical bond.

6 Li + F Li + F - The Ionic Bond 1s22s11s22s1 1s22s22p51s22s22p5 1s21s2 1s22s22p61s22s22p6 [He][Ne] Li Li + + e - e - + FF - F - Li + + Li + F - LiF Ionic bond: the electrostatic force that holds ions together in an ionic compound.

Essential Features of Ionic Bonding Atoms with low I.E. and low E.A. tend to form positive ions. Atoms with high I.E. and high E.A. tend to form negative ions. Ion formation takes place by electron transfer. The ions are held together by the electrostatic force of the opposite charges. Reactions between metals and nonmetals (representative) tend to be ionic.

Describing Ionic Bonds Noble gas configurations are stable. The atoms of metals loses electron(s) to become a cation (positive). eg., Na →Na + [Ne]3s 1 → [Ne] + 1 e - (Na has low I.E, easy to loose an e - ) The atom that gains the electron becomes an anion (negative), eg., Cl, Cl + e→Cl - (Cl has high E.A) 3s 2 3p 5 + e- →3s 2 3p 6 or [Ar]

Another example Consequently, magnesium can accommodate two fluorine atoms. : : F. : : : F. : Mg.. [ F ] : : : : -2+ [ F ] : : : : - The magnesium has two electrons to give, whereas the fluorines have only one “vacancy” each.

9.1 Use Lewis dot symbols to show the formation of aluminum oxide (Al 2 O 3 ). The mineral corundum (Al 2 O 3 ).

9.1 Thus, the simplest neutralizing ratio of Al 3+ to O 2− is 2:3; two Al 3+ ions have a total charge of +6, and three O 2− ions have a total charge of −6. So the empirical formula of aluminum oxide is Al 2 O 3, and the reaction is

Ionic bonds MgO – magnesium oxide 2Mg(s)+O 2 (g)→ 2MgO

13 Lattice energy increases as Q increases and/or as r decreases. CompoundLattice Energy (kJ/mol) MgF 2 MgO LiF LiCl Q: +2,-1 Q: +2,-2 r F - < r Cl - Electrostatic (Lattice) Energy E = k Q+Q-Q+Q- r Q + is the charge on the cation Q - is the charge on the anion r is the distance between the ions Lattice energy (U) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions. E is negative for ionic bond formation, but lattice energy is the reverse, energy to break a bond, which is +ve.

Born-Haber Cycle for Determining Lattice Energy  H overall =  H 1 +  H 2 +  H 3 +  H 4 +  H 5 oooooo Lattice energy = +1017

16 A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Why should two atoms share electrons? FF + 7e - FF 8e - F F F F Lewis structure of F 2 lone pairs single covalent bond Covalent Bond Noble gas config.

Features of Covalent Bonds The diatomic elements have totally covalent bonds (totally equal sharing.) H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 Covalent compounds are usually formed from nonmetals. Molecules - compounds characterized by covalent bonding. not a part of a massive three dimensional crystal structure.

8e - H H O ++ O HH O HHor 2e - Lewis structure of water single covalent bonds Hydrogen molecule, diatomic

19 Double bond – two atoms share two pairs of electrons O C O or O C O 8e-8e- 8e-8e- 8e-8e- double bonds Triple bond – two atoms share three pairs of electrons N N 8e-8e- 8e-8e- N N triple bond or Carbon dioxide and nitrogen molecule

Lengths of Covalent Bonds Bond Lengths Triple bond < Double Bond < Single Bond

Polar Covalent Bonds A polar covalent bond is one in which the bonding electrons spend more time near one of the two atoms involved. When the atoms are alike, as in the H-H bond of H 2, the bonding electrons are shared equally (a nonpolar covalent bond). When the two atoms are of different elements, the bonding electrons need not be shared equally, resulting in a “polar” bond.

H F F H Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms electron rich region electron poor region e - riche - poor ++ --

24 Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Electron Affinity - measurable, Cl is highest Electronegativity - relative, F is highest X (g) + e - X - (g)

Polar Covalent Bonds Electronegativity is a measure of the ability of an atom in a molecule to draw bonding electrons to itself. In general, electronegativity increases from the lower-left corner to the upper- right corner of the periodic table. The current electronegativity scale, developed by Linus Pauling, assigns a value of 4.0 to fluorine and a value of 0.7 to cesium.

The Electronegativities of Common Elements

Variation of Electronegativity with Atomic Number

Covalent share e - Polar Covalent partial transfer of e - Ionic transfer e - Increasing difference in electronegativity Classification of bonds by difference in electronegativity DifferenceBond Type 0Covalent  2 Ionic 0 < and <2 Polar Covalent

The greater the difference in electronegativity between two atoms, the greater the polarity of a bond. Which would be more polar, a H-F bond or a H- Cl bond? H-F … = 1.9 H-Cl… = 0.9 HF bond is more polar than the HCl bond.

Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H 2 S; and the NN bond in H 2 NNH 2. Cs – 0.7Cl – – 0.7 = 2.3Ionic H – 2.1S – – 2.1 = 0.4Polar Covalent N – – 3.0 = 0Covalent

Chemical bonding (cont.) Classify the following bonds as ionic, polar covalent, or covalent: E.N. diff (a)the bond in HCl3.0 – 2.1 = 0.9 polar covalent (a)The bond in KF4.0 – 0.8 = 3.2 Ionic (b)the CC bond in H 3 CCH – 2.5 = 0 covalent

32 1.Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. 2.Count total number of valence e -. 3.Complete an octet for all atoms except hydrogen. 4.If structure contains too many electrons, form double and triple bonds on central atom as needed. Writing Lewis Structures

9.3 Write the Lewis structure for nitrogen trifluoride (NF3) in which all three F atoms are bonded to the N atom. NF 3 is a colorless, odorless, unreactive gas.

9.3 Solution We follow the preceding procedure for writing Lewis structures. Step 1: The N atom is less electronegative than F, so the skeletal structure of NF 3 is Step 2: The outer-shell electron configurations of N and F are 2s 2 2p 3 and 2s 2 2p 5, respectively. Thus, there are 5 + (3 × 7), or 26, valence electrons to account for in NF 3.

9.3 Step 3: We draw a single covalent bond between N and each F, and complete the octets for the F atoms. We place the remaining two electrons on N: lone pair of electrons bonding pair Because this structure satisfies the octet rule for all the atoms, step 4 is not required. Check Count the valence electrons in NF 3 (in bonds and in lone pairs). The result is 26, the same as the total number of valence electrons on three F atoms (3 × 7 = 21) and one N atom (5). How about NH 3 ? Similar to NF 3, Difference for H ?

Multiple Bonds In the molecules described so far, each of the bonds has been a single bond, that is, a covalent bond in which a single pair of electrons is shared. It is possible to share more than one pair. A double bond involves the sharing of two pairs between atoms. orC : C H H H H :: : : :

Triple bonds are covalent bonds in which three pairs of electrons are shared between atoms. Multiple Bonds CCor HH :: :::

38 Formal Charge of an atom An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. formal charge on an atom in a Lewis structure = 1 2 total number of bonding electrons () total number of valence electrons in the free atom - total number of nonbonding electrons - The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.

39 C – 4 e - O – 6 e - 2H – 2x1 e - 12 e - 2 single bonds (2x2) = 4 1 double bond = 4 2 lone pairs (2x2) = 4 Total = 12 H CO H formal charge on C = ½ x 8 = 0 formal charge on O = ½ x 4 = 0 formal charge on an atom in a Lewis structure = 1 2 total number of bonding electrons () total number of valence electrons in the free atom - total number of nonbonding electrons - 00 Formaldehyde, CH 2 O

Lewis structure and formal charge of carbonate ion The Lewis structure for the carbonate ion is, formal charge on an atom in a Lewis structure = 1 2 total number of bonding electrons () total number of valence electrons in the free atom - total number of nonbonding electrons -

The Lewis structure for the carbonate ion is, C: =0 O(a): =0 O (b):6-6-1 = -1 O ©: = -1, Total of -2 for carbonate ion. Lewis structure and formal charge of carbonate ion formal charge on an atom in a Lewis structure = 1 2 total number of bonding electrons () total number of valence electrons in the free atom - total number of nonbonding electrons -

Resonanace A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. (+ and – are the formal charges) O-O bond in O 3 should be longer than the double bond, but both oxygen bonds are equal in length because of resonance as shown. OOO + - OOO + -

Delocalized Bonding: Resonance According to theory, one pair of bonding electrons is spread over the region of all three atoms. This is called delocalized bonding, in which a bonding pair of electrons is spread over a number of atoms. OO O

. OCO O -- OCO O - - OCO O - - What are the resonance structures of the Benzene? carbonate (CO 3 2 -) ion? benzene Carbonate ion

Lewis Structures and Exceptions to the Octet Rule 1.Incomplete Octet - less then eight electrons around an atom other than H. Let’s look at BF 3 2.Odd Electron - if there is an odd number of valence electrons it isn’t possible to give every atom eight electrons. Let’s look at NO 3.Expanded Octet - elements in 3rd period and beyond may have 10 and 12 electrons around it.

Exceptions to the Octet Rule The Incomplete Octet HHBeBeH 2 BF 3 FBF F

Exceptions to the Octet Rule Odd-Electron Molecules N – 5e - O – 6e - 11e - NO N O The Expanded Octet (central atom with principal quantum number n > 2) Electrons can be accomodated in the d orbitals SF 6 S – 6e - 6F – 42e - 48e - S F F F F F F 6 single bonds (6x2) = lone pairs (18x2) = 36 Total = 48

Exceptions to the Octet Rule For example, the bonding in phosphorus pentafluoride, PF 5, shows ten electrons surrounding the phosphorus. : F : : : F : : : : F : : P F : : :

Exceptions to the Octet Rule In xenon tetrafluoride, XeF 4, the xenon atom must accommodate two extra lone pairs. F : : : : F : : Xe F : : : : F : : : :

Expanded octet is the most common exception. Write the Lewis structures of SH 6 and SF 4

The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy. H 2 (g) H (g) +  H 0 = kJ Cl 2 (g) Cl (g) +  H 0 = kJ HCl (g) H (g) +Cl (g)  H 0 = kJ O 2 (g) O (g) +  H 0 = kJ OO N 2 (g) N (g) +  H 0 = kJ N N Bond Energy Bond Energies Single bond < Double bond < Triple bond

Bond Energies (BE) and Enthalpy changes in reactions  H 0 = total energy input – total energy released =  BE(reactants) –  BE(products) Imagine reaction proceeding by breaking all bonds in the reactants and then using the gaseous atoms to form all the bonds in the products. (Examples 9.13,9.14)

H 2 (g) + Cl 2 (g) 2HCl (g)2H 2 (g) + O 2 (g) 2H 2 O (g)

Use bond energies to calculate the enthalpy change for: H 2 (g) + F 2 (g) 2HF (g)  H 0 =  BE(reactants) –  BE(products) Type of bonds broken Number of bonds broken Bond energy (kJ/mol) Energy change (kJ) HH FF Type of bonds formed Number of bonds formed Bond energy (kJ/mol) Energy change (kJ) HF  H 0 = – 2 x = kJ

Bond Energy To illustrate, let’s estimate the  H for the following reaction. In this reaction, one C-H bond and one Cl- Cl bond must be broken. In turn, one C-Cl bond and one H-Cl bond are formed.

Bond Energy simple arithmetic yields  H.  H 0 = [BE (C-H) + BE (Cl-Cl)] – [BE (C-Cl) + BE (H-Cl)] =[( )] – [( )] kJ = -104 kJ