Chemical Bonding I: Basic Concepts Chapter 8 Copyright © The McGraw-Hill Companies, Inc.  Permission required for reproduction or display.

Lewis Dot Symbols Consists of the symbol of an element and one dot for each valence electron in an atom of the element. Note that (except helium) the number of valance electrons each atom has is the same as the group number of the element.

Ionic Bond Is the electrostatic force that holds ions together in an ionic compound The metal gives the electrons to the non metal Cation is the metal Anion is the non metal

The Covalent Bond A bond in which two electrons are shared by two atoms. Only in covalent compounds Non metal and non metal

The Covalent Bond

The Covalent Bond Pairs of valence electrons that are not involved in covalent bond formation are called: Lone pairs We can only draw Lewis structures for compounds that have covalent bonds

Polar Covalent Bonds Though atoms often form compounds by sharing electrons, the electrons are not always shared equally. Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.

Electronegativity The ability of an atom to attract toward itself the electrons in a chemical bond

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

Lewis Structures Is a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs of dots on individual atoms Only valence electrons are shown in a Lewis structure

Writing Lewis Structures Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. If it is an anion, add one electron for each negative charge. If it is a cation, subtract one electron for each positive charge. PCl3 Keep track of the electrons: 5 + 3(7) = 26

Writing Lewis Structures The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds. Keep track of the electrons: 26 − 6 = 20

Writing Lewis Structures Fill the octets of the outer atoms. Keep track of the electrons: 26 − 6 = 20; 20 − 18 = 2

Writing Lewis Structures Fill the octet of the central atom. Keep track of the electrons: 26 − 6 = 20; 20 − 18 = 2; 2 − 2 = 0

Writing Lewis Structures If you run out of electrons before the central atom has an octet… …form multiple bonds until it does.

The "best" Lewis structure for NO3- 1.  Determine the total number of valence electrons in a molecule Draw a skeleton

Of the 24 valence electrons in NO3-, 6 were required to make the skeleton. Consider the remaining 18 electrons and place them so as to fill the octets of as many atoms as possible

Are the octets of all the atoms filled Are the octets of all the atoms filled?   If not then fill the remaining octets by making multiple bonds Check that you have the lowest FORMAL CHARGES possible for all the atoms, without violating the octet rule; (valence e-) - (1/2 bonding e-) - (lone electrons).

Writing Lewis Structures The best Lewis structure… …is the one with the fewest charges. …puts a negative charge on the most electronegative atom.

The Concept of Resonance Resonance structure, is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure Resonance means the use of two or more Lewis structures to represent a particular molecule

Exceptions to the Octet Rule The expanded octet Atoms in the 2nd period (that are in families 3A-7A) cannot have more than 8 valance electrons around the central atom Atoms of elements in and beyond the 3rd period (that are in families 3A-7A) form some compound in which more than 8 electron surround the central atom Odd- electron molecules Some molecules contain an odd number of electrons. We need an even number of electron for complete pairing the octet rule clearly cannot be satisfied with all the atoms in any of these molecules Incomplete octet The number of electrons surrounding the central atom in a stable molecule is fewer than eight

The modern periodic table. Figure 2.10 The modern periodic table. +1 +2 +3 NC -3 -2 -1

Bond Enthalpy A measure of the stability of a molecule Thus bond enthalpy is the enthalpy change required to break a particular bond in 1 mole of gaseous molecules Ho(reaction) = sum of the bond energies of bonds being broken - sum of the bond energies of the bonds being formed. Ho(reaction) = H(reactant bonds broken) - H(product bonds formed)

Example of Using Bond Energies to Calculate Heat (enthalpy) of Reaction Use the bond energies provided in the table above to calculate the heat (enthalpy) of reaction, Ho, for the reaction: CH4(g) + 4Cl2(g) -----> CCl4(g) + 4HCl(g) Write the balanced chemical equation, with all reactants and products in the gaseous state. CH4(g) + 4Cl2(g) -----> CCl4(g) + 4HCl(g) Write the general equation for the heat (enthalpy) of reaction: Ho(reaction) = H(reactant bonds broken) - H(product bonds formed)

Bonds Broken Bonds Formed bond type bond energy bond type bond energy Substitute bond energy values into the equation and solve for Ho(reaction) Bonds Broken Bonds Formed bond type bond energy bond type bond energy 4 x C – H 4 x 413= 1652 4 x C - Cl 4 x 328 = 1312 4 x Cl – Cl 4 x 243 = 972 4 x H - Cl 4 x 432 = 1728 H(reactant bonds broken) = 2624 H(product bonds formed) = 3040 Ho(reaction) = 2624 - 3040 = -416 kJ mol-1