Bonding: General Concepts

Bonding: General Concepts

• Source: Chapter 3 Zumdahl, CHEMISTRY: An Atoms First Approach, Third Edition. © 2021 Cengage.

Table of Contents

• Section 3.1: Types of chemical bonds

• Section 3.2: Electronegativity

• Section 3.3: Ions: Electron configurations and sizes

• Section 3.4: Partial ionic character of covalent bonds

• Section 3.5: The covalent chemical bond: A model

• Section 3.6: The localized electron bonding model

• Section 3.7: Lewis structures

• Section 3.8: Exceptions to the octet rule

• Section 3.9: Resonance

• Section 3.10: Naming simple compounds

Questions to Consider

• What is meant by the term “chemical bond”?

• Why do atoms bond with each other to form compounds?

• How do atoms bond with each other to form compounds?

Section 3.1: Types of Chemical Bonds

• Chemical Bond Definition: The force that holds atoms together is called a chemical bond.

• Covalent Bond:

  • Formed when atoms share electrons.

  • A collection of atoms makes a molecule.

• Molecular Representations:

  • Chemical Formula: Example CO2

  • Structural Formula: Example H—O—H

  • Space-filling Model: Indicates both relative sizes of atoms and their orientation in the molecule.

  • Ball and Stick Model: A three-dimensional model using spheres (atoms) and rods (bonds).

Attributes of Molecules

• Physical Properties include:

  • Melting Point

  • Hardness

  • Electrical and Thermal Conductivity

  • Solubility

  • Electric Charge

  • Bond Energy: The amount of energy needed to break a bond.

Types of Chemical Bonding

• Ionic Bonding:

  • Occurs between an atom that easily loses electrons (typically metals) and an atom that has a high affinity for electrons (typically non-metals).

  • Ionic compounds are typically formed when metals react with non-metals.

Coulomb’s Law

• Used to calculate the energy of interaction between ions:

  • Formula:
    E = k rac{Q1 Q2}{r}
    where:

  • E: Energy, units in joules

  • r: Distance between ions in nanometers

  • Q1 and Q2: Numerical charge values of the ions.

Stability of Atoms in Bonding

• Example:

  • In the H2 molecule, electrons reside in the space between two nuclei leading to increased stability compared to single H atoms.

  • Increased attractive forces result in a decrease in the potential energy of electrons.

Polar Covalent Bond

• Definition: Unequal sharing of electrons between atoms in a molecule.

  • Results in charge separation within the bond, creating partial positive (δ+) and partial negative (δ−) charges.

Section 3.2: Electronegativity

• Definition: The ability of an atom in a molecule to attract shared electrons to itself.

• Pauling Method:

  • For a molecule HX, relative electronegativities of H and X are assessed by comparing bond energies:

  • Actual Bond Energy vs. Expected Bond Energy.

Determining Values of Electronegativity

• Charge distribution based on electronegativity differences can be understood through variations in bond energy:

  • Formula:
    

  \Delta E = ext{Actual Bond Energy - Expected Bond Energy}
  

  • If Δ = 0, H and X have identical electronegativities.

  • If the electronegativity of X exceeds that of H, shared electrons are closer to atom X, influencing charge distribution.

Table 3.1: Relationship Between Electronegativity and Bond Type

• List of Bonds Ordered by Polarity based on their electronegativity differences (values provided in parentheses):

  • H—H (2.1), S—H (2.1), Cl—H (3.0), O—H (3.5), F—H (4.0)

Section 3.3: Ions - Electron Configurations and Sizes

• Electron Configuration in Compounds:

  • In a reaction between two nonmetals, electron sharing enables both to achieve a complete valence electron configuration, imitating noble gas states.

  • In reactions involving nonmetals and representative-group metals, binary ionic compounds are formed; ions achieve noble gas configurations.

Concept Check: Isoelectronic Series

• Task: Choose ions from alkali metals, alkaline earth metals, noble gases, and halogens that form an isoelectronic series.

• Concepts Involved:

  • Electron configurations, number of electrons and protons, and physical properties including ionic radius and ionization energy.

Predicting Formulas of Ionic Compounds

• Definition of Ionic Compounds:

  • Chemists define these as compounds in solid states where the attraction between positive and negative ions is maximized due to close proximity.

  • Example of valence electron configurations:

  • Calcium: ext{Ca} [ ext{Ar}]4s^2

  • Oxygen: ext{O} [ ext{He}]2s^22p^4

• Since oxygen has a higher electronegativity than calcium, the prediction of compound formulas can be established for consistency with stable noble gas configurations.

Sizes of Ions

• Factors influencing ionic size:

  • Size of the parent atom.

  • Position in the periodic table: Isoelectronic ions contain the same number of electrons.

Section 3.4: Partial Ionic Character of Covalent Bonds

• Percent Ionic Character Formula:

  • No bond is 100% ionic; thus the ionic character is calculated through comparative studies.

  • Example of calculating ionic character from electronegativity differences across various bonds and determining the effect of bond type on ionic character.

Ionic Compounds

• Identification Problems:

  • No bond considered fully ionic; many substances comprise polyatomic ions.

  • Operational definition: Any compound conducting electricity when melted is classified as ionic.

Section 3.5: The Covalent Chemical Bond - A Model

• Models Overview:

  • Used to explain microscopic behaviors based on macro observations and properties.

  • Provide frameworks that help understand chemical behaviors, emphasizing that models simplify reality.

Section 3.6: The Localized Electron Bonding Model

• Localized Electron (LE) Model:

  • Molecules consist of atoms bound through shared electron pairs from atomic orbitals.

  • Lone pairs: Non-bonding electron pairs on atoms.

  • Bonding pairs: Found in space between atoms.

Lewis Structures

• Definition: Illustrates valence electron arrangement in molecules, named after G. N. Lewis.

• Principles of Writing Lewis Structures:

  • Sum of valence electrons from all atoms.

  • Use a pair of electrons to bond atoms.

  • Arrange remaining electrons to fulfill the duet rule for hydrogen and octet rule for row 2 elements.

Exceptions to the Octet Rule

• Boron Compounds: Tends to have fewer than eight electrons around it (boron fluoride example).

  • Adding double bonds can achieve an octet.

• Exceeding the Octet Rule: Seen in elements from Period 3 or beyond, e.g., sulfur hexafluoride.

Resonance

• Definition: More than one valid Lewis structure for some molecules.

  • Example: Nitrate ion (NO3-); true molecular structure is a superposition of all valid structures.

Formal Charge

• Definition: Used for estimating the charge on atoms in Lewis structures:

  • Formula:
    ext{Formal charge} = ( ext{valence electrons}) - ( ext{assigned electrons})

• Assumptions: Lone pairs belong entirely to an atom, while shared electrons are divided equally.

Naming Simple Compounds

• Binary Ionic Compounds (Type I):

  • Compounds consist of a cation followed by an anion.

  • Rules: Cation named first, anion named second as -ide root.

  • Example Table:

  • Cation: H+, Li+, Na+, etc.

  • Anion: H− (hydride), F− (fluoride), etc.

Binary Ionic Compounds (Type II)

• Metals forming more than one type of positive ion are named using Roman numerals to signify charge.

• Naming rules adjust based on higher (-ic) or lower (-ous) charge nomenclature.

Oxyanions

• Specific naming conventions depending on oxygen presence:

  • -ate for higher, -ite for lower, -hypo and -per for additional variations in oxygen content.

Acids

• Naming rules based on anion suffixes:

  • -ide becomes hydra- and -ic, while -ate becomes -ic and -ite becomes -ous.

  • Example acids:

  • HCl (hydrochloric acid), H2SO4 (sulfuric acid).