CHEMICAL BONDING

Chemical Bonding

Formation of Ions

·        The stable number of electrons in the outermost shell. Except for the 1^st shell is 8. An atom tends to complete the outermost shell either by gaining or losing or sharing electrons in order to have a stable or closed configuration. Once atom gains or loses electrons it becomes charged. A charged atom is called an ion.

Formation of Cations

·        Atoms which have 1, 2, or 3 valence electrons tend to lose electrons and become positively charged atoms called cations. The number of electrons lost is the positive valence of the element.

Example:       ₁₁Na – 1s² 2s² 2p⁶ 3s¹

                                    Valence e^- = 1

                                    Valence    = +1

·        Cations are smaller than their corresponding neutral atoms.

Example Problem

            Predict the ions that magnesium and aluminum are most likely to form.

Formation of Anions 

·        Atoms which have 5, 6, or 7 valence electrons tend to gain electrons and become negatively charged atoms called anions. The number of electrons gained is the negative valence of the element.

Example:       ₁₇Cl – 1s² 2s² 2p⁶ 3s² 3p⁵

                                    Valence e^- = 7

                                    Valence    = -1

·        Nonmetals have negative electron affinities and generally form anions.

·        Anions are larger than their corresponding neutral atoms.

The Ionic Bond

·        An ionic bond is formed by the electrostatic attraction of oppositely charged ions.

§  Ionic compounds form between metals and nonmetals.

§  Ionic compound usually consist of elements that are widely separated in the periodic table: a metal from left hand side and nonmetal from the right.

§  Forming an ionic bond between a metal and nonmetal usually requires energy to form the ion pair.

§  Ionization energies are positive.

§  Electron affinities for nonmetals are negative.

The Covalent Bond

·        A covalent bond is based on the sharing of pairs of electrons between two atoms.

§  Ionic bonding lowers energy by transferring electrons between a metal and nonmetal.

§  Covalent bonding lowers energy by sharing electrons between two nonmetals.

§  Bond energy – energy released when isolated atoms form a covalent bond.

§  Formation of bonds always releases energy.

Chemical Bonds and Structure of Molecules

·        During ionic bond formation, the cations and anions achieve np⁶ electronic configurations (noble gas configurations).

§  Metals lose electrons.

§  Nonmetals gain electrons.

§  During covalent bond formation, electrons are shared between two atoms.

§  Shared electrons are available to both bonding atoms.

§  Sharing leads to 8 valence electrons around each atom.

·        Octet rule – an atom will form covalent bonds to achieve a complement of eight valence electrons.

§  For the n = 1 shell, hydrogen violates the octet rule and shares only 2 electrons.

·        Lewis dot symbols keep track of valence electrons, especially for main group elements, allowing prediction of bonding in molecules.

§  To draw a Lewis dot symbol, the valence electrons are represented by dots and are placed around the element symbol.

§  The first four dots are placed singly.

§  Starting with the fifth dot, they are paired.

The second period Lewis symbols are shown below:

§  Elements within a group have the same number of valence electrons and identical Lewis dot symbol.

·        Lewis dot structures show how electrons are shared in a molecule.

§  A pair of shared electrons between two atoms is a bonding pair.

§  Bonding pairs represented by a line between two atomic symbols.

Pairs of electrons associated with one atom are nonbonding or lone pair electrons

·        By sharing an electron from each atom, two hydrogen atom can from a covalent bond.

§  Hydrogen violates the octet rule by sharing only two electrons.

When two fluorine atoms combine, they form a stable covalent bond.

§  By sharing a pair of electrons, each atom is surrounded by eight valence electrons.

·        Bonding atoms in molecules can share more than one bonding pair of electrons.

§  A single bond results when one bonding pairs are shared.

§  A double bond results when two bonding pairs are shared.

§  A triple bond results when three bonding pairs are shared.

Example: HF (single), O₂ (double), N₂ (triple)

Example Problem

            Draw the Lewis structure of the following ionic compound:

                        a. K₂O                                                 b. MgCl₂

Electronegativity and Bond Polarity

• ·      Bonding between the two ends of the bonding continuum, ionic and covalent bonding, is described using electronegativity and bond polarity.

§  Electronegativity is the attraction of an atom for the shared electrons in a covalent bond

§  The higher the electronegativity value, the more likely an element will attract extra electron density during compound formation.

• ·        Electronegativities increase from left to right across a period and from the bottom to top for a group.

• ·        Fluorine is the most electronegative element, with an electronegativity of 4.0.

• ·        The greater the electronegativity difference, the more polar the bond. In polar bond electron cloud is not equally distributed among the atoms.

            Ex: HCI: 3.0-2.1=0.9

• ·        When the electronegativity difference is zero, the bond is classified as nonpolar covalent. Electron cloud is equally distributed among the atoms.

Ex: all di atomic molecules like Cl₂ : 3.0-3.0=0

• ·        When the electronegativity difference exceeds 2.0, the bond is classified as ionic.

Ex: NaCl: 3.0-0.9=2.1 

   ∆EN                                                    Ionic character

  > 1.7                                                   Mostly ionic

  0.4 – 1.7                                             Polar covalent

  <0.4                                                    Mostly Covalent

     0                                                          Non-polar

·        The formation of the polar covalent HF bond

§  The more electronegative F has a partial negative charge.

§  The less electronegative H has a partial positive charge.

Example Problem

            Which bond is the most polar: C-H, O-H, or H-Cl?

            Electronegativity values: H = 2.1, C = 2.5, O = 3.5, Cl = 3.0

Keeping Track of Bonding: Lewis Structures

·        Lewis structures indicate how many bonds are formed and between which elements in a compound.

·        Step 1 – Count the total valence electrons in the molecule or ion.

§ Sum the number of valence electrons for each element in a molecule.

§ For ions, add or subtract valence electrons to account for the charge.

§ For the compound OF₂, the number of valence electron is 20.

F        2 x 7 = 14

O       1 x 6 = 6

          Total = 20

·        Step 2 – Draw the “skeletal structure” of the molecule.

§  The element written first in the formula is usually the central atom, unless it is hydrogen.

§  Usually, the central atom is the least electronegativity.

         

            F     O

               F

     Step 3 – Place single bonds between all connected atoms in the structure by drawing lines between them

§  A single line represents a bonding pair.

§  Four electrons are placed in bonds.

§  Sixteen electrons are left to place.

·        Step 4 – Place the remaining valence electrons not accounted for on individual atoms until the octet rule is satisfied. Place electrons as lone pairs whenever possible.

§  Place electrons first on outer atoms, then on central atoms.

§  Six electrons are placed as lone pairs on each F satisfies the octet rule for each F.

§ The four remaining electrons are placed on the O to satisfy the octet rule for each O.

 

·        Step 5 – Create multiple bonds by shifting lone pairs into bonding positions as needed for any atoms that do not have a full octet of valence electrons.

§  Correctly choosing which atoms to form multiple bonds between comes from experience.

§  Multiple bonds are not required for OF₂, as the octet rule is satisfied for each atom.

Example Problem

            Draw the Lewis structure of dichlorofluromethane, CF₂Cl₂, also known as DuPont’s Freon-12.

            Calcium phosphate is an important precursor for the formation of bioceramic coatings. Draw the Lewis structure of the phosphate ion, PO­₄^3-.

Practice Problem

            Poly (vinyl alcohol) is used in several biomaterials applications, including surgical structures. Draw the Lewis structure of vinyl alcohol, CH₂CHOH, the monomer from which poly (vinyl alcohol) is made.

Example Problem 7.6

            Draw the Lewis structure for the NH₄⁺ and ClO₂^- ions.

Steps in solving bonding pair and lone pair:

• Solve for all the valence electrons of the combining atom.

•  Determine the number of electron required to complete the octet.

• Determine the bonding electrons by subtracting answer in step 1 from answer in step 2.

• Determine the nonbonding electrons by subtracting answer in step 3 from step 1.

 

Solve for the bonding pair and lone pair of the following:

1. H₂O

     2.      HF

  

     3. CCl₄

4        4. HCN

RECORDED LECTURES

Lewis Diagrams Made Easy: How to Draw Lewis Dot Structures

https://www.youtube.com/watch?v=cIuXl7o6mAw&t=55s

Lewis Dot Structures

https://www.youtube.com/watch?v=Sk7W2VgbhOg&t=10s

The Chemical Bond: Covalent vs. Ionic and Polar vs. Nonpolar https://www.youtube.com/watch?v=PoQjsnQmxok&t=10s

ONLINE PUBLISHED RESEARCHES

Hypervalency” and the chemical bond 

https://doi.org/10.1016/j.comptc.2019.02.014. 

(http://www.sciencedirect.com/science/article/pii/S2210271X19300593)

Chemical bonding without orbitals

https://doi.org/10.1016/j.comptc.2018.10.004.(http://www.sciencedirect.com/science/article/pii/S2210271X18305656) 

Chemical Bonding—The Formation of Materials

https://doi.org/10.1016/B978-0-12-810425-5.00003-5.

(http://www.sciencedirect.com/science/article/pii/B9780128104255000035)

The Chemical Bond: The Perspective of NMR Spectroscopy

https://doi.org/10.1016/bs.arnmr.2016.07.001.

(http://www.sciencedirect.com/science/article/pii/S0066410316300278)