Chapter 10: Chemical Bonding – Structured Study Notes

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Chemical Bonding

Bonding Models and Drug Design

Bonding theories are essential for understanding how atoms combine to form molecules, which is crucial in fields such as drug design. For example, the structure of HIV-protease was used to design drugs that inhibit its function, preventing the virus from replicating in the human body. These drugs, known as protease inhibitors, have significantly improved the life expectancy of HIV-infected individuals.

HIV-protease and drug molecule simulation

Bonding Theories

Overview of Bonding Theories

Bonding theories predict how atoms bond to form compounds and explain why certain combinations are stable. They also help determine molecular shapes, which influence physical and chemical properties.

Lewis Theory: Uses dot structures to represent valence electrons and predict molecular stability.
Advanced Theories: Used for more accurate predictions but Lewis theory is preferred for quick, everyday applications.

Lewis Theory: Representing Valence Electrons

In Lewis theory, valence electrons are shown as dots around the element symbol. The number of valence electrons for main-group elements (except helium) equals the group number.

Electron configuration of oxygen Lewis structure of oxygen

Dot Structure: Each dot represents a valence electron, with a maximum of two dots per side.
Octet: Atoms with eight valence electrons are particularly stable.

Lewis structures for Period 2 elements

Writing Lewis Structures for Elements

To write a Lewis structure, place the correct number of dots (valence electrons) around the element symbol.

Example: Phosphorus (Group 5A) has five valence electrons.

Lewis structure of phosphorus

Helium: An Exception

Helium has two valence electrons, forming a stable duet rather than an octet.

Electron configuration and Lewis structure of helium

Lewis Theory of Chemical Bonding

Types of Chemical Bonds

Lewis theory states that chemical bonds involve the sharing or transfer of electrons to achieve stable electron configurations.

Ionic Bond: Electrons are transferred (metal to nonmetal).
Covalent Bond: Electrons are shared (nonmetal to nonmetal).

The octet rule states that atoms transfer or share electrons to obtain eight electrons in their outer shell, except for hydrogen, lithium, and beryllium, which achieve stability with two electrons (duet).

Lewis Structures of Ionic Compounds

When metals bond with nonmetals, electrons are transferred, forming cations and anions. The attraction between these ions forms ionic compounds.

Lewis structures of potassium and chlorine Electron transfer from potassium to chlorine

Example: Potassium transfers its valence electron to chlorine, forming K+ and Cl−.

Lewis structure of KCl with charges

Lewis Structure of MgO

Magnesium loses two electrons to oxygen, forming Mg2+ and O2−.

Lewis structures of magnesium and oxygen Lewis structure of MgO with charges

Predicting Chemical Formulas for Ionic Compounds

Lewis theory helps predict the correct formula by balancing electron transfer.

Example: Sodium (Na) and sulfur (S) form Na2S, as two sodium atoms are needed to provide two electrons for sulfur.

Lewis structures of sodium and sulfur Lewis structure of Na2S with charges

Covalent Lewis Structures: Electrons Shared

Covalent Bonding

Nonmetals form molecular compounds by sharing electrons, allowing each atom to achieve an octet (or duet for hydrogen).

Lewis structures of hydrogen and oxygen Lewis structure of water (H2O) Duet and octet in water molecule

Bonding Pair Electrons: Shared between atoms, represented by dashes.
Lone Pair Electrons: Not shared, located on one atom.

Bonding pair and lone pair in water Lewis structure of water with dashes

Halogens and Diatomic Molecules

Halogens form diatomic molecules by pairing together to achieve octets.

Lewis structure of chlorine Lewis structure of Cl2

Hydrogen as H2

Hydrogen atoms share electrons to form H2, achieving a stable duet.

Lewis structure of H2

Double and Triple Bonds

Atoms can share more than one pair of electrons to achieve octets. Double bonds involve two pairs, triple bonds involve three pairs.

Example: O2 forms a double bond; N2 forms a triple bond.

Lewis structure of oxygen Double bond formation in O2 Lewis structure of O2 with double bond Octet in O2 molecule Lewis structure of nitrogen Triple bond formation in N2

Writing Lewis Structures for Covalent Compounds

Steps for Writing Lewis Structures

Write the correct skeletal structure for the molecule.
Calculate the total number of valence electrons.
Distribute electrons to give octets (or duets for hydrogen).
If any atoms lack an octet, form double or triple bonds as needed.

Example: CO2 Lewis Structure

CO2 has a linear structure with two double bonds between carbon and oxygen.

CO2 skeletal structure CO2 with lone pairs CO2 with double bonds

Polyatomic Ions and Resonance

Lewis Structures for Polyatomic Ions

For polyatomic ions, adjust the electron count for the ion's charge and enclose the structure in brackets with the charge indicated.

Example: CN− ion has a triple bond between C and N.

CN- skeletal structure CN- with lone pairs CN- with triple bond and charge

Exceptions to the Octet Rule

Some molecules have odd numbers of electrons (free radicals) or fewer/more than eight electrons around the central atom (e.g., BF3, BH3, SF6, PCl5).

Lewis structure of NO Lewis structures of BF3 and BH3 Lewis structures of SF6 and PCl5

Resonance Structures

Some molecules, like SO2, have multiple valid Lewis structures. The actual structure is an average (resonance) of these forms.

SO2 with bonding electrons SO2 with lone pairs SO2 with double bond SO2 resonance structures SO2 resonance with double-headed arrow

Predicting Molecular Shapes

VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on the repulsion between electron groups around the central atom.

Linear Geometry: CO2 has a bond angle of 180°.
Trigonal Planar: H2CO has bond angles of ~120°.
Tetrahedral: CH4 has bond angles of 109.5°.

CO2 linear geometry Lewis structure of CO2 Lewis structure of H2CO H2CO trigonal planar geometry CH4 tetrahedral geometry Space-filling model of CH4

Molecules with Lone Pairs

Lone pairs affect molecular geometry. For example, NH3 is trigonal pyramidal, and H2O is bent.

Lewis structure of NH3

Electronegativity and Polarity

Electronegativity

Electronegativity is the ability of an atom to attract electrons in a bond. Differences in electronegativity lead to polar covalent bonds, with partial charges and dipole moments.

Example: In water, oxygen is more electronegative than hydrogen, resulting in a partial negative charge on oxygen and a partial positive charge on hydrogen.

Bond Types Based on Electronegativity Difference

Pure Covalent: Identical electronegativities (e.g., Cl2).
Ionic: Large difference (e.g., NaCl).
Polar Covalent: Intermediate difference (e.g., HF).

Polar Bonds and Polar Molecules

Polar bonds do not always result in polar molecules. The overall molecular geometry determines whether dipole moments cancel or add.

Example: CO2 is nonpolar due to linear geometry; H2O is polar due to bent geometry.

Everyday Chemistry: Soap and Molecular Interactions

How Soap Works

Soap molecules have a polar head and a nonpolar tail, allowing them to interact with both water (polar) and grease/oil (nonpolar), enabling the removal of grease from surfaces.

Chapter 10 Review

Lewis Theory: Explains ionic and covalent bonding.
Molecular Shapes: Predicted by Lewis and VSEPR theories.
Electronegativity: Determines bond polarity.
Polarity: Polar bonds may or may not result in polar molecules.

Chemical Skills Learning Objectives

Write Lewis structures for elements and compounds.
Predict chemical formulas for ionic compounds.
Write resonance structures.
Predict molecular shapes.
Determine molecular polarity.

Free Radicals and Aging

Free Radicals

Free radicals are molecules with an odd number of valence electrons and an unpaired electron, making them highly reactive. They are implicated in aging and disease processes.