Basic Structure of Carbon and Chemical Bonding Study Notes

Unit I: Basic Structure of Carbon and Chemical Bonding

Prepared by: Ghellyn T. Gajeles, Ph.D., RCh.

Main Branches of Chemistry

  • Organic chemistry: Study of hydrocarbons and their various derivatives.
  • Inorganic chemistry: Study of all substances other than hydrocarbons and their derivatives.

Types of Compounds

  • Compounds are categorized into two types:
    • Organic compounds: Derived from living organisms (with a vital force).
    • Inorganic compounds: Derived from minerals (without a vital force).
Definition of Organic Chemistry
  • Organic chemistry is defined as the study of compounds that contain carbon.

Characteristics of Organic Compounds

  • Key attributes of organic compounds:
    • Are made from carbon atoms.
    • Contain one or more carbon atoms.
    • Typically have many hydrogen atoms.
    • May also contain oxygen (O), sulfur (S), nitrogen (N), phosphorus (P), and halogens (F, Cl, Br, I).

Properties of Organic Compounds

  • Typical properties include:
    • Contain carbon.
    • Have covalent bonds.
    • Generally characterized by low melting points.
    • Typically have low boiling points.
    • Are flammable.
    • Soluble in nonpolar solvents.
    • Not soluble in water (example: oil vs. water).

Organic vs. Inorganic Compounds

  • Example Comparison:
    • Propane (C₃H₈): Organic compound used as fuel.
    • NaCl (Sodium Chloride): Inorganic compound made up of Na⁺ and Cl⁻ ions.
Question to Consider
  • Why is propane classified as an organic compound while NaCl is not?

Comparing Properties of Organic and Inorganic Compounds

  • Table 11.1: Some Properties:
    • Organic Examples: C and H, sometimes O, S, N, P, halogens (e.g., C₂H₈).
    • Inorganic Examples: Most metals, Na, and Cl.
    • Particles: Organic (mostly molecules) vs. Inorganic (mostly ions).
    • Bonding: Mostly covalent in organics, primarily ionic in inorganics, with some covalent.
    • Polarity of Bonds: Organics are nonpolar unless there is a strongly electronegative atom; most inorganics are ionic or polar covalent.
    • Melting/Boiling Points: Organics usually low (e.g., -188 °C for melting, -42 °C for boiling), compared to high values for inorganics (e.g., 801 °C for melting, 1413 °C for boiling).
    • Flammability: High for organics (burns in air) vs. low for inorganics (does not burn).
    • Solubility in Water: Not soluble (organics unless a polar group is present) vs. often soluble for inorganics unless nonpolar.

Learning Check

  • Identify characteristics typical of inorganic or organic compounds:
    • 1. High melting point: Inorganic (I)
    • 2. Not soluble in water: Organic (O)
    • 3. Formula C₃H₇: Organic (O)
    • 4. Formula MgCl₂: Inorganic (I)
    • 5. Burns easily in air: Organic (O)
    • 6. Has covalent bonds: Organic (O)

What Makes Carbon Special?

  • Key Properties:
    • Atoms to the left of carbon give up electrons.
    • Atoms to the right of carbon accept electrons.
    • Carbon shares electrons.

The Structure of an Atom

  • Subatomic particles:
    • Protons: Positively charged particles.
    • Neutrons: Neutral particles.
    • Electrons: Negatively charged particles.
  • Atomic number: Equal to the number of protons.
  • Carbon’s atomic number: 6.

Atomic Symbol Convention

  • Mass number (A): Total number of protons and neutrons (A = Z + number of neutrons).
  • Atomic number (Z): Number of protons.
General Notes
  • Number of protons = number of electrons in neutral atoms.
  • Charge of an ion = number of protons – number of electrons (applicable for monoatomic ions).

Isotopes

  • Definition: All carbon atoms share the same atomic number; however, they can have different mass numbers due to differing numbers of neutrons.
  • Naturally occurring carbon isotopes:
    • C-12 (6 protons, 6 neutrons, 6 electrons)
    • C-13 (6 protons, 7 neutrons, 6 electrons)
    • C-14 (6 protons, 8 neutrons, 6 electrons)

The Distribution of Electrons in an Atom

  • Shell Structure:
    • 1st shell has a capacity for 2 electrons (1s).
    • 2nd shell has a capacity of 8 electrons (2s, 2p).
    • 3rd shell can have up to 18 electrons (3s, 3p, 3d).
Electron Configuration
  • Aufbau principle: Electrons fill atomic orbitals of the lowest energy first.
  • Pauli exclusion principle: No more than two electrons can occupy a single orbital (must have opposite spins).
  • Hund’s rule: Electrons will occupy empty degenerate orbitals before pairing.

Terms in Electron Configuration

  • Ground-state configuration: Orbitals occupied by electrons when they are at lowest energy level.
  • Excited-state configuration: Occurs when one or more electrons move to higher-energy orbitals.
  • Core electrons: Inner shell electrons; do not participate in bonding.
  • Valence electrons: Electrons in the outermost shell that participate in bonding.

Condensed Electron Configurations

  • Abbreviated configurations that indicate innermost electrons using noble gas symbols.
  • The structure of 1s² 2s² 2p⁶ is contained within noble gas notation.

Lewis Symbols

  • Developed by G. N. Lewis to represent bonding electrons using dots around the elemental symbol.
  • Octet rule: Atoms tend to gain, lose, or share electrons to have eight valence electrons.

Electrons and Bonding

  • Atoms on the Left Side of the Periodic Table: Tend to lose electrons (e.g., lithium loses an electron to form Li⁺).
  • Atoms on the Right Side of the Periodic Table: Tend to gain electrons (e.g., fluorine gains an electron to form F⁻).

Models of Chemical Bonding

  • Ionic bonding: Attraction between oppositely charged ions (e.g., Na⁺ and Cl⁻).
  • Covalent bonding: Sharing of electrons between atoms.
  • Metallic bonding: Involves a sea of delocalized electrons around positively charged ions.

Covalent Bonds

  • Formed by sharing electrons.
  • Types of covalent bonds:
    • Nonpolar covalent bond: Atoms are identical.
    • Polar covalent bond: Atoms are different and result in unequal sharing of electrons.
  • Bond Polarity: Depends on the difference in electronegativity.
Bond Type Indicators
  • Polar Covalent Bond: Electronegativity difference between atoms is 0.4-1.9.
  • Ionic Bond: Electronegativity difference > 2.0.
  • Nonpolar Covalent Bond: Electronegativity difference is 0-0.39.

Dipole Moments and Bond Polarity

  • Dipole moment is a measure of the polarity in a bond.
  • Table 1.4 provides dipole moments for various bonds, illustrating that greater electronegativity differences lead to higher dipole moments.

Formal Charge

  • Formal Charge Formula:
    Formal Charge=number of valence electrons−(number of lone pair electrons+12number of bonding electrons)\text{Formal Charge} = \text{number of valence electrons} - (\text{number of lone pair electrons} + \frac{1}{2} \text{number of bonding electrons})

Structural Representation of Molecules

  • Lewis Structures represent how atoms are connected in a molecule including details about bonding and lone pairs.
  • Expanded structural formulas, Kekulé structures, and condensed structures illustrate molecular connectivity and functional groups.

Acid and Base Theories

  • Arrhenius Acid: A substance that produces H⁺ ions in solution.
  • Arrhenius Base: A substance that produces OH⁻ in solution.
Bronsted-Lowry Theory
  • Bronsted–Lowry acid: A proton donor.
  • Bronsted–Lowry base: A proton acceptor.
Conjugate Acid-Base Pairs
  • These are formed when an acid donates a proton, resulting in its corresponding conjugate base, and vice versa.

Acid Strength

  • Equilibrium Position: A strong acid dissociates completely in water, whereas a weak acid does not.
  • Strengths of acids are often determined by their dissociation constants (Ka).
  • Higher Ka values correspond to stronger acids.

Acids and Bases Classification

  • Monoprotic: Supplies one proton (e.g., HCl).
  • Diprotic: Supplies two protons (e.g., H₂CO₃).
  • Triprotic: Supplies three protons (e.g., H₃PO₄).
  • Polyprotic: Supplies two or more protons (includes diprotic and triprotic).

Factors Affecting Acidity

  • Electronegativity: More electronegative elements stabilize negative charges better, leading to stronger acids.
  • Inductive Effects: Electron-withdrawing groups can stabilize conjugate bases, increasing acidity.
  • Hybridization: Acidity increases as hybridization becomes more s character.

Conclusion

  • This unit covered the basics of carbon structure and bonding, different types of chemical reactions, and acid-base theories among other chemistry principles. Additional study and practice are recommended to better understand these concepts and their applications.