Chapter 8

1. Chapter 8 BY: Angela, Charlie, Henry, Taylor, &Yvonne

2. Chemical Bonds • Chemical Bond- whenever two atoms or ions are strongly attached to each other. • Ionic Bond- bond between metals and nonmetals • Covalent bond- bond between nonmetals • Metallic bond- bond between metals 8.1 Chemical Bonds, Lewis Symbols, and The Octet Rule

3. Lewis Symbols • The chemical symbol for an element, with a dot for each valence electron.

4. The Octet Rule • The Octet Rule: Atoms tend to gain, lose, or share electrons until they surrounded by eight valence electrons. • An octet of electrons consists of full s and p subshells in an atom. • *There are exceptions to the rule.

5. 8.2 Ionic Bonding • Electron transfer • The ionization energy- indicates how easily an electron can be removed from an atom. • Electron affinity- measures how much an atom wants to gain an electron. • Electron transfer which forms oppositely charged ions occurs when one of the atoms readily gives up an electron(I. E.) and the other atom readily gains an electron(E.A.)

6. Energetics of Ionic Bond Formation Electron configurations of Ions • Exothermic process- the gain of electrons to an atom • Endothermic process- the loss of electrons from an atom • Lattice energy- is the energy required to completely separate a mole of a solid ionic compound into gaseous ions. • Ions tend to have noble-gaselectron configurations • For example : Sodium (Na) Neon (Ne)

7. 8.3 Covalent Bonding Characteristics of Ionic Substances -usually brittle, high melting points, crystalline, ionic crystals can be cleaved; break apart along smooth flat surfaces -these characteristics result from electrostatic forces that maintain the ions in a rigid, well-defined, 3-d arrangement

8. 8.3 Covalent Bonding • Covalent Bond- a chemical bond formed by sharing of ONE AND ONLY pair of electrons • Example H2, hydrogen molecule • To depict this bond one would use the Lewis Structure, which is defined as a diagram that shows the bonds between atoms (as a line) and the unshared electron pairs as dots ExampleCO2, double bond. For nonmetals only- the number of valence electrons in a neutral atom is the same as the group number Multiple bonds 1. Single bond - sharing of only one pair of electrons 2. Double bond - sharing of two electron pairs 3. Triple bond - corresponds to the sharing of three pairs of electrons *bonds are shown as a line* *non-bonds are shown as dots* *bond length - the distance between the nuclei of the atoms in a bond*

9. 8.4 Bond Polarity and Electronegativity Bond Polarity- helps describe the sharing of electrons between atoms There are 2 types of Covalent bonds: 1. Non-polar - one in which the electrons are shared equally between two atoms 2. Polar-one of the atoms exerts a greater attraction for the bonding electrons than the other *if the difference in relative ability to attract electrons is large enough, an ionic bond is formed* Electronegativity- a quantity that is used to estimate whether a bond will be non-polar covalent, polar covalent, or ionic • Also defined as the ability of an atom in a molecule to attract electrons to itself *the greater an atom's electronegativitythe greater is it's ability to attract electrons to itself.* *in relationto the ionization of energy, which is a measure of how strong a gaseous atom holds on to its electrons*

10. 8.4 Bond Polarity and Electronegativity Electron Affinity- a measure of how strongly an atom attracts additional electrons *an electron with very negative electron affinity values and high ionization energy will both attract electrons from other atoms andresist having its electrons attract away - it will be highly electronegative* *electronegativitydecreases with increasing atomic number in any one group* Examples: Non-polar covalent F24.0-4.0=0 Polar Covalent HF4.0-2.1=0 Ionic Bond LiF4.0-1.0=3.0 *the greater the difference in electronegativity between two atoms, the more polar the bond*

11. 8.4 Bond Polarity and Electronegativity Polar Molecule Dipole- the centers of positive and negative charge do not coincide a distance separates two electrical charges of equal magnitude but opposite sign • Basically, this is a force that results in bonding between two electrons Dipole Moment – quantitative measure of the magnitude of the dipole (µ) *the dipole moment increases as the magnitude of charge and the distance between the charges that is separated increases µ=Qr The magnitude of the dipole moment is the product of Q and R Dipole moments are reported in debyes (D)

12. Quite simply, a way of representing molecular compounds (Steps to be explained) Examples: NH3CO2 8.5: Drawing Lewis Structures H N H O C O H

13. Defined as: charge atom would have if all atoms in molecule had same electronegativty(each bonding pair shared =ly) TO CALCULATE! Start with a Lewis structure. All nonbonding e- are assigned to atom on which found For any bond (x1,x2,x3) ½ of bonding e- are assigned to each atom in bond Then, subtract the # of e- assigned to each atom in bond If it sounds confusing, everything becomes much simpler by drawing a diagram! 8.5: Formal Charge

14. (Using a Lewis Structure…) Valence e-: 6 4 6 - (e- assigned to atom): 6 4 6 _________________________________________________________________ Formal Charge: 0 0 0 8.5: Formal Charge O C O

15. Important: Formal charge does not represent actual charge (although it adds up to it) But why is it useful?– It can be used to determine the best Lewis Structure because… Usually, we use the L.S. with the F.C. closest to 0 Usually, we use the L.S. in which – charges reside on more electronegative atoms. 8.5: Formal Charge

16. L.S.’s in which the placement of the atoms is the same but the placement of the electrons is different. • Shown by drawing the possible structures and connecting with arrow(s) Example! Ozone (03) 8.6: Resonance Structures

17. Exceptions to the Octet Rule

18. Impossible to pair NO, NO2, ClO2, O2- Highly reactive Double bonds make resonance structures Bonding cannot be expressed by a single Lewis formula Odd Numbers of Electrons

19. Boron and Beryllium Double bonds make resonance structures In the example with double bond: Fluorine is sharing extra electrons giving it a +1 charge and Boron a -1 charge BF3 reacts strongly with compounds with unshared pairs of electrons such as ammonia, NH3 Less than an octet of valence electrons

20. Octet rule is based ns and nporbitals Third period elements often exceed the Octet Rule NCl5 and OF4 do not exist The central atom is forced to “expand” in order to add more than 8 electrons The larger the central atom, the more electrons can surround it Expansion occur most often when the central atom is bonded to small electronegative atoms, such as F, Cl and O. More than an octet of valence electrons

21. Bond Enthalpy- the enthalpy change, ∆H, for the breaking of a particular bond in one mole of a gaseous substance. • Ex. Enthalpy change of the bonds of chlorine atoms in the Cl2 molecule when 1 mol of Cl2 is dissociated into chlorine atoms: Cl-Cl2Cl • The bond enthalpy is always a positive quantity • You can use average bond enthalpies to estimate the enthalpies of reactions in which bonds are broken and new bonds are formed. The equation is: ∆HRxN=∑(bond enthalpies of broken bonds)- ∑(bond enthalpies of bonds formed) • Ex. The gas-phase reaction between methane, CH4, and chlorine to produce methyl chloride, CH3Cl, and hydrogen chloride, HCl: H - CH3(g) + Cl – Cl(g)  Cl – CH3(g) + H – Cl(g), ∆HRxN=? Bonds broken: 1 mol C – H, 1 mol Cl – Cl Bonds made: 1 mol C – Cl, 1 mol H – Cl ∆HRxN= [D(C – H) + D(Cl – Cl)] – [D(Cl – Cl) + D(H – Cl)] = (413 kJ + 242kJ) – (328kJ + 431kJ)= -104kJ Section 8.8