Basic Concepts of Chemical Bonding Chapter 8 Basic Concepts of Chemical Bonding

6.3 Describing Chemical Bonding Page: 288 - 330

CHEMICAL BONDS IONIC COVALENT METALLIC Electrostatic attraction between ions. COVALENT Sharing of electrons. METALLIC Metal atoms bonded to several other atoms.

The Ionic Bonding

Na(s) + ½Cl2(g)  NaCl(s) 8.2 The Ionic Bonding Na(s) + ½Cl2(g)  NaCl(s)

transfer occurs readily? The Ionic Bonding Na: loss of an electron Cl: gain of an electron How Can you Know Whether this electron transfer occurs readily?

4. Electronegativity Difference transfer occurs readily? The Ionic Bonding 1. Ionization energy 2. Electron Affinity 3. Lattice Energy 4. Electronegativity Difference How Can you Know Whether this electron transfer occurs readily?

One species must have very low ionization energy (Na) The Ionic Bonding Na: loss of an electron Cl: gain of an electron Ionization Energy One species must have very low ionization energy (Na)

One species must have very high electron affinity (Cl) The Ionic Bonding Na: loss of an electron Cl: gain of an electron Ionization Energy Electron Affinity One species must have very high electron affinity (Cl)

ELECTROSTATIC ATTRACTIONS The Ionic Bonding Na(s) + ½Cl2(g)  NaCl(s) Δ Hf˚= -410.9 kJ ELECTROSTATIC ATTRACTIONS

LE is a measure of a stability of ions arranged within an ionic solid Lattice Energy Attraction between oppositely charged ions  release of energy  formation of lattice LE is a measure of a stability of ions arranged within an ionic solid

ΔEN > 1.7 An Ionic Bond would form between these species Electronegativity difference If the difference in electronegativity (ΔEN) between the species is: ΔEN > 1.7 An Ionic Bond would form between these species

The Ionic Bonding SUMMARY

6.3 THE COVALENT BONDING

Covalent Bonding

Covalent Bonding

ELECTRONEGATIVITY is the ability of atoms in a molecule to attract electrons to themselves.

DIPOLE and ELECTRONEGATIVITY DIPOLE: a partial separation of charge One end of the molecule has slightly positive charge The other end of the molecule has slightly negative charge

DIPOLE and ELECTRONEGATIVITY Slight excess of a negative charge Slight excess of a positive charge Slight excess of a positive charge

DIPOLES are caused by ELECTRONEGATIVITY Hydrogen has lower electronegativity The electrons spend LESS time around Hydrogen Chlorine has higher electronegativity The electrons spend MORE time around Chlorine

BOND POLARITY Homo nuclear diatomic compounds: In different compounds, electrons are not shared equally Homo nuclear diatomic compounds: H2, Cl2, O2,N2 … share electrons equally = NON POLAR COVALENT BOND

BOND POLARITY δ+ = slightly positive δ- = slightly negative Hetero nuclear diatomic compounds: HF, HCl, NO, NaCl, LiO… DO NOT share electrons equally POLAR COVALENT BOND(the attraction of one of the atoms for the bonding electrons is LARGE) δ+ = slightly positive δ- = slightly negative

BOND POLARITY H – F δ+ δ- Hetero nuclear diatomic compounds: HF, HCl, NO, NaCl, LiO… DO NOT share electrons equally POLAR COVALENT BOND(the attraction of one of the atoms for the bonding electrons is LARGE)

BOND POLARITY H – F δ+ δ- Hetero nuclear diatomic compounds: HF, HCl, NO, NaCl, LiO… DO NOT share electrons equally MOSTLY COVALENT BOND (the attraction of one of the atoms for the bonding electrons is SLIGHTLY GREATER)

POLAR, NONPOLAR, or ionic? BOND POLARITY Hetero nuclear diatomic compounds: HF, HCl, NO, NaCl, LiO… DO NOT share electrons equally POLAR COVALENT BOND(the attraction of one of the atoms for the bonding electrons is LARGE) IONIC BONDS (the attraction of one of the atoms for the bonding electrons is VERY LARGE) HOW DO YOU KNOW IF THE GIVEN BOND IS POLAR, NONPOLAR, or ionic?

ELECTRONEGATIVITY The greater the difference in electronegativity, the more polar the bond is

ΔE = 0 ΔE > 1.7 ELECTRONEGATIVITY 0 < ΔE < 0.4 NON POLAR COVALENT BOND ΔE = 0 H2, Cl2, O2,N2 MOSTLY COVALENT BOND 0 < ΔE < 0.4 H2O, CO2, HF POLAR COVALENT BOND 0.4 < ΔE < 1.7 H2O, CO2, HF IONIC BOND ΔE > 1.7 MgO, NaCl, LiF

NON POLAR COVALENT BOND ELECTRONEGATIVITY NON POLAR COVALENT BOND H2, Cl2, O2,N2 POLAR COVALENT BOND H2O, CO2, HF IONIC BOND MgO, NaCl, LiF electronegativity difference > 2.1

In each case, which bond is more polar: EXAMPLE In each case, which bond is more polar: a) B – Cl or C – Cl, b) P – F or P – Cl? Indicate in each case which atom has the partial negative charge. B – Cl 1.0 F – P 1.9 δ+ δ- δ- δ+ C – Cl 0.5 P – Cl 0.9 δ+ δ- δ- δ+

IONIC vs COVALENT Compounds Bonding COVALENT Bonding Metal + nonmetal High melting point Lattice(crystal) structures Strong electrolytes Nonmetal + nonmetal Low melting point Low boiling point Non - electrolytes

HOMEWORK PAGE: 324 - 323 PROBLEMS: all even

6.4 LEWIS STRUCTURE DIAGRAMS

LEWIS DIAGRAMS - REVIEW Show only an atom’s valence electrons and the chemical symbol.

LEWIS DIAGRAMS

Rule # 1 Dots representing valence electrons are placed around the element symbols

Rule # 2 Electron dots are placed singly until the fifth electron is reached then they are paired

Lewis Diagrams of IONS and IONIC BONDS For positive ions, one electron dot is removed from the valence shell for each positive charge. For negative ions, one electron dot is added to each valence shell for each negative charge. Square brackets are placed around each ion to indicate transfer of electrons.

LEWIS STRUCTURE DIAGRAMS for MOLECULES Lewis structures are representations of molecules showing all valence electrons: bonding and nonbonding

LEWIS STRUCTURE DIAGRAMS for MOLECULES Lewis structures are representations of molecules showing all valence electrons: bonding and nonbonding a lone pair

Lewis Structures and Multiple Bonds When two electron pairs are shared, two lines are drawn, i.e. double bond ::O :: C :: O:: or ::O = C = O::

Lewis Structures and Multiple Bonds When three electron pairs are shared, three lines are drawn, i.e. triple bond :N ::: N: or :N ≡ N:

N – N N = N N ≡ N BOND LENGTHS > > BOND STRENGTH? The length of the bond between two atoms decreases as the number of shared electrons increases N – N > N = N > N ≡ N BOND STRENGTH?

If it is cation  subtract e- STEPS TO FOLLOW… Find the sum of valence electrons of all atoms in the polyatomic ion or molecule If it is anion  add e- If it is cation  subtract e- PCl3 5 + 21 valence electrons = 26 v.e.

PCl3 5 + 21 valence electrons = 26 v.e. STEPS TO FOLLOW… 2. Write the symbols for the atoms Make one of the atoms a central atom (usually the least electronegative atom) connect it with the other atoms by single bonds PCl3 5 + 21 valence electrons = 26 v.e.

26 – 6 valence electrons = 20 v.e. left STEPS TO FOLLOW… 2. Write the symbols for the atoms Subtract those electrons from your total number of valence electrons 26 – 6 valence electrons = 20 v.e. left

3. Fill the octets of the outer atoms STEPS TO FOLLOW… 3. Fill the octets of the outer atoms 20 valence electrons

3. Fill the octets of the outer atoms STEPS TO FOLLOW… 3. Fill the octets of the outer atoms 20 valence electrons

20 valence electrons STEPS TO FOLLOW… Subtract the added electrons from your total number of valence electrons 20 valence electrons

20 – 18 valence electrons = 2 v.e. left STEPS TO FOLLOW… Subtract the added electrons from your total number of valence electrons 20 – 18 valence electrons = 2 v.e. left

4. Fill the octet of the central atom STEPS TO FOLLOW… 4. Fill the octet of the central atom 2 valence electrons

4. Fill the octet of the central atom STEPS TO FOLLOW… 4. Fill the octet of the central atom 0 valence electrons

5. If you run out of electrons before the central atom has an octet… STEPS TO FOLLOW… 5. If you run out of electrons before the central atom has an octet… …form multiple bonds until it does

Draw Lewis Structures for WORKSHEET EXAMPLE Draw Lewis Structures for CH2Cl2 C2H4 BrO3-

Draw Lewis Structures for WORKSHEET EXAMPLE Draw Lewis Structures for NO BF3 PF5

EXCEPTIONTS TO THE OCTET RULE 1. For molecules and polyatomic ions containing an odd number of valence electrons 2. For molecules and polyatomic ions in which an atom has fewer than an octet of valence electrons

EXCEPTIONTS TO THE OCTET RULE 3. For molecules and polyatomic ions in which an atom has more than an octet of valence electrons

1. For molecules and polyatomic ions containing an odd number of valence electrons Complete pairing of valence electrons is impossible due to the odd number of valence electrons E.g.: ClO2, NO, NO2, O2-

Mostly Boron or Beryllium compounds 2. For molecules and polyatomic ions in which an atom has fewer than an octet of valence electrons Not very common Mostly Boron or Beryllium compounds

3. For molecules and polyatomic ions in which an atom has more than an octet of valence electrons Very common Such molecules/ions are called HYPERVALENT Only for atoms of 3rd period or higher WHY? 1. They have available and unfilled d orbitals for bonding 2. Their central atom (P, S, I, Xe…) is large enough to be bonded to even five different atoms (Cl, F or O)

Draw a Lewis Structure for ion: ICl4- EXAMPLE Draw a Lewis Structure for ion: ICl4-

Draw a Lewis Structure for the thiocyanate ion: NCS- EXAMPLE Draw a Lewis Structure for the thiocyanate ion: NCS-

WHICH ONE IS THE MOST IMPORTANT? WHICH ONE IS CORRECT ? WHICH ONE IS THE MOST IMPORTANT?

Calculate THE FORMAL CHARGE of each ion to find out…

FORMAL CHARGE The charge the atom would have if all the atoms in the molecule had the same electronegativity ALL unshared electrons are assigned to the atom on which they are found For a single/double/triple bond, half of the bonding electrons is assigned to to each atom in the bond

FORMAL CHARGE OF AN ATOM = # of VALENCE ELECTRON – # OF ELECTRONS ASSIGNED ALL unshared electrons are assigned to the atom on which they are found For a single/double/triple bond, half of the bonding electrons is assigned to to each atom in the bond

EXAMPLE What are the formal charges of C and N in the cyanide ion: CN-? ALL unshared electrons are assigned to the atom on which they are found 2. For a single/double/triple bond, half of the bonding electrons is assigned to to each atom in the bond [:C ≡ N:]- # of VALENCE ELECTRON – # OF ELECTRONS ASSIGNED

EXAMPLE What are the formal charges of C and N in the cyanide ion: CN-? ALL unshared electrons are assigned to the atom on which they are found 2. For a single/double/triple bond, half of the bonding electrons is assigned to to each atom in the bond [:C ≡ N:]- C N Valence e- 4 5 Assigned e- Formal Charge -1 -1 # of VALENCE ELECTRON – # OF ELECTRONS ASSIGNED

What are the formal charges on the thiocyanate ions? WORKSHEET EXAMPLE What are the formal charges on the thiocyanate ions? 1. The most important (most dominant) Lewis structure is the one which has its value closest to 0 (the one with the fewest charges) 2. The most important (most dominant) Lewis structure is the one which has any negative charges reside on the more electronegative atoms WHICH STRUCTURE IS THE MOST IMPORTANT (DOMINANT) ONE?

What are the formal charges the thiocyanate ions? WORKSHEET EXAMPLE What are the formal charges the thiocyanate ions? 1. The most important (most dominant) Lewis structure is the one which has its value closest to 0 (the one with the fewest charges) 2. The most important (most dominant) Lewis structure is the one which has any negative charges reside on the more electronegative atoms WHICH STRUCTURE IS THE MOST IMPORTANT (DOMINANT) ONE?

The cyanate ion, NCO-, has three possible Lewis structures. WORKSHEET EXAMPLE The cyanate ion, NCO-, has three possible Lewis structures. Draw these three structures assign formal charges in each. Which Lewis structure is dominant?

Draw a Lewis Structure of OZONE, O3 8.6 RESONANCE STRUCTURES Draw a Lewis Structure of OZONE, O3

Not a single resonance structure RESONANCE STRUCTURES Not a single color OZONE: A mix of different resonance structures GREEN PAINT: A mix of different colors Not a single resonance structure

Somewhere between single bond and double bond RESONANCE STRUCTURES Same bond lengths Somewhere between single bond and double bond

Is the use of two or more Lewis structures to represent a molecule RESONANCE STRUCTURES Is the use of two or more Lewis structures to represent a molecule this is because that molecule can not be represented by only a single Lewis structure

RESONANCE STRUCTURES In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygen atoms and the carbon. They are not localized; they are delocalized.

RESONANCE STRUCTURES The organic compound benzene, C6H6, has two resonance structures. It is a hexagon with a circle inside to signify the delocalized electrons in the ring.

What are the resonance structures of nitrate ion: NO3- EXAMPLE What are the resonance structures of nitrate ion: NO3-

WORKSHEET EXAMPLE Using Formal Charges and the concept of Resonance, show why the correct Lewis structure of BF3 is the one in which B has an incomplete octet and only single bonds are present.

8.8 STRENGTHS OF COVALENT BONDS Learn on your own

HOMEWORK PAGE: 325 - 326 PROBLEMS: 49, 51, 53, 55, 57, 61, 67, 69, 71, 73,

ANSWERS TO WORKSHEET EXAMPLES

Copy this table and fill out the missing information. WORKSHEET EXAMPLE Copy this table and fill out the missing information.

WORKSHEET EXAMPLE Which substance would you expect to have the greatest lattice energy, MgF2, CaF2, and ZrO2 ? ZrO2

WORKSHEET EXAMPLE Ionizing an H2 molecule to H2+ changes the strength of the bond. Based on the description of covalent bonding given previously, do you expect the bond H—H in H2+ to be weaker or stronger than the bond in H2? Weaker. In both H2 and H2+ the two H atoms are principally held together by the electrostatic attractions between the nuclei and the electron(s) concentrated between them. H2+ has only one electron between the nuclei whereas H2 has two and this results in the H—H bond in being stronger.

WORKSHEET EXAMPLE How does the ELECTRONEGATIVITY of an element differ from its ELECTRON AFFINITY? Electron affinity measures: the energy released when an isolated atom gains an electron to form a 1- ion. The electronegativity measures: the ability of the atom to hold on to its own electrons the ability of the atom to attract electrons from other atoms in compounds.

In each case, which bond is more polar: WORKSHEET EXAMPLE In each case, which bond is more polar: a) B – Cl or C – Cl, b) P – F or ? Indicate in each case which atom has the partial negative charge.

The cyanate ion, NCO-, has three possible Lewis structures. WORKSHEET EXAMPLE The cyanate ion, NCO-, has three possible Lewis structures. Draw these three structures assign formal charges in each. Which lewis structure is dominant?

WORKSHEET EXAMPLE Using Formal Charges and the concept of Resonance, show why the correct Lewis structure of BF3 is the one in which B has an incomplete octet and there are only single bonds present. Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine. This would not be an accurate picture of the distribution of electrons in BF3. Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.