Chapter 2. Classical Atom Theory
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4. Drawing Molecular Structures
1. History of Atomic Models
⑴ 1st. Democritus
① Particle Theory (5th century BC): Proposed that all matter is composed of indivisible particles
② Atom: Derived from the Greek word “atomos,” meaning indivisible
③ Refuted by Plato and Aristotle
⑵ 2nd. Dalton
① Atomic Theory: Accurate definition of the atom (1808)
② Hypothesis 1: All matter consists of indivisible particles
○ Counterexamples: Atoms can be divided into a nucleus and electrons
○ Counterexamples: Atoms can be further broken down by nuclear fission
③ Hypothesis 2: Same types of atoms have the same size and properties; different types have different sizes and masses
○ Counterexample: Isotopes
④ Hypothesis 3: Atoms are neither created nor destroyed and do not change into different kinds
○ Counterexample: Nuclear reactions
⑤ Hypothesis 4: Different atoms combine in fixed ratios to form new substances
⑶ 3rd. Thomson
Figure 1. Thomson’s Cathode Ray Experiment
① Thomson’s Cathode Ray Experiment: In 1897, using a vacuum discharge tube, discovered that the flow of cathode rays was the flow of particles known as electrons
② Rejection of Dalton’s Atomic Theory: Atoms can be divided into negatively charged electrons and positively charged protons
③ Proposal of the Plum Pudding Model
○ Overall structure resembles a positively charged sphere
○ Electrons with equal negative charges are scattered throughout
○ Significance: Partial explanation of atomic electric properties
○ Limitation **: ** Rutherford’s alpha particle scattering experiment
⑷ 4th. Millikan
① Millikan Experiment: Determined the quantized charge of electrons
② Calculated the mass of electrons from their charge
⑸ 5th. Rutherford
Figure 2. Rutherford’s Experiment
① Purpose of Rutherford’s Experiment: To verify Thomson’s Plum Pudding Model
② Experimental Process (1911)
○ Created a thin gold foil with a thickness of 1/20,000 cm
○ Fired alpha particles from radioactive sources at the gold foil
○ Placed a zinc sulfide screen behind the foil to detect small flashes produced by alpha particles hitting
③ Expected Results: According to Thomson’s model, there should be no deflection of alpha particles
④ Experimental Results
Figure 3. Model of α-Particle Scattering
○ A few alpha particles scattered in unexpected directions upon impact (failed to pass through the foil)
○ About 1 out of 8,000 alpha particles scattered at 180° (backward scattering)
⑤ Interpretation of Results
Figure 4. Rutherford’s Atomic Model
○ Alpha particles not passing through the 1/20,000 cm gold foil ≒ Bullets not passing through thin tissue paper
○ Presence of a heavy, dense structure with significant positive charge: Introduction of atomic nucleus
○ In other words, alpha particles scatter because their mass is smaller than the gold foil’s atomic nucleus
○ Solar system model proposed (1910): Atoms with a central nucleus around which electrons orbit
⑥ Limitations 1. Stability of Atoms
○ Cannot explain stability of electrons in circular orbits
○ Accelerating electrons emit light energy and eventually fall into the nucleus
⑦ Limitations 2. Hydrogen Gas Line Spectrum: According to the solar system model, a continuous spectrum should be observed
⑹ 6th. Bohr Atomic Model
⑺ 7th. Discovery of Protons and Neutrons
① In 1896, Goldstein (E. Goldstein) discovered positive rays
② 1914, Rutherford Experiment: Suggested the existence of protons
③ In 1932, Chadwick (J. Chadwick) discovered particles (neutrons) in the atomic nucleus without a charge
○ Chadwick was a student of Rutherford
2. Classical Atomic Model
⑴ Atom
① Atomic Theory: Introduced by Dalton
② Atom = Nucleus + Electrons, Origin: Indivisible (a-) particles (tom)
③ Atomic Nucleus = Protons + Neutrons
④ Atomic Number: Number of protons = Number of electrons in a neutral atom
○ In ions, number of protons ≠ number of electrons
⑤ Atomic Mass Number = Number of protons + Number of neutrons
⑥ Atomic Representation (Atomic Symbol)
○ A: Atomic Mass
○ B: Atomic Number
○ C: Charge
○ D: Number of atoms
⑦ Isotopes
○ Elements with the same atomic number but different mass numbers: Difference in number of neutrons
○ All elements have two or more isotopes
⑵ Periodicity
① Periodic Table: Electrons, protons, neutrons, atomic mass
Figure 5. Periodic Table
○ Period: Number of electron shells. 1st period, 2nd period, ···
○ Group: Set of atoms with the same number of outermost electrons
○ Metal: Group of atoms with a strong tendency to donate electrons
○ Nonmetal: Group of atoms with a strong tendency to gain electrons
○ Metalloid: Possesses properties of both metals and nonmetals, 7 elements (B, Si, Ge, As, Sb, Te, Po)
○ Main Group Elements: s-block, p-block
○ Transition Elements: d-block
○ Lanthanides: f-block
○ Common oxidation state is +3
○ Typically form coordination compounds with coordination numbers greater than 6
○ React with acids to release hydrogen
○ Form stable complexes using oxygen as a ligand
○ Actinides: f-block
② Periodic Behavior of Atoms: Involves forces between nucleus and electrons (number of protons in the nucleus, nuclear-electron distance), electron-electron repulsion
③ Atomic Radius
○ Atomic size cannot be precisely defined due to the probabilistic nature of electron clouds
○ Definition 1: Half the distance between the nuclei of two atoms bonded together (excluding noble gases)
○ Rational to define the atomic radii of O2 and N2 as half the distance between their nuclei
○ Unreasonable to define the atomic radius of HF using this approach
○ Definition 1 concept can also be applied to metallic bonding
○ Definition 2: Defined using Van der Waals radius (applies to noble gases)
○ Ion Radius Definition: Consideration of distance and weighting between ions in ionic bonding
○ Trend 1: Atomic radius decreases with increasing atomic number in the same period
○ Trend 2: Atomic radius increases with increasing number of electron shells in the same group
④ Effective Nuclear Charge: The positive charge felt by an electron, a measure of the attractive force considering electron repulsion
○ Comparison of effective nuclear charge within an atom: Inner orbitals experience greater effective nuclear charge
○ Comparison of effective nuclear charge of outermost electrons between atoms
○ Within the same period, atomic radius decreases as atomic number increases, leading to an increase in effective nuclear charge
○ Within the same group, effective nuclear charge slightly increases as the number of electron shells increases (opposite trend to ionization energy)
○ With an increase in the number of electron shells, shielding effect weakens, strengthening nuclear repulsion, as understood
○ Comparison of effective nuclear charge of specific orbitals between atoms
○ With an increase in atomic number, the effective nuclear charge of a specific orbital mentioned in the problem increases
○ Slater’s Rule: A formula for effective nuclear charge
○ If the number of valence electrons in the outermost shell is denoted as n, then the S constant values for n, n-1, n-2, n-3, ··· are 0.35, 0.85, 1, 1, ···
○ When calculating the effective nuclear charge for n-level electrons, subtract one because oneself does not influence the calculation
○ For example, applying S = 0.3 to the effective nuclear charge calculation for 1s orbit electrons
Table 1. Example of Applying Slater’s Rule
⑤ First Ionization Energy (IE)
○ Definition
○ X(g) → X+(g) + e-(g), IE = E(X+) - E(X). IE > 0.
○ Orbital interpretation: Energy needed to move an electron from E1 to infinity
○ Trend 1: Increases as the number of electron shells decreases, contrary to the trend of effective nuclear charge
○ Trend 2: Increases with increasing period
Figure 6. Trends in Ionization Energy
○ Comparison of 2nd and 13th groups
○ Orbital energy level is ns < np, so ionization energy is lower in the 13th group
○ For the same reason, the 2nd ionization energy is lower in the 14th group compared to the 13th group
○ Comparison of 15th and 16th groups
○ Ionization energy is lower in the 16th group due to repulsion between paired electrons
○ Nitrogen is stable as [He] 2s22p3, while oxygen as [He] 2s22p4 is less stable
○ Element with the highest ionization energy: Helium
○ Sequential Ionization Energies: Requires very high energy to remove inner electrons, allowing determination of the atomic number
○ 1st Ionization Energy (1st I.E.): E(X+) - E(X)
○ 2nd Ionization Energy (2nd I.E.): E(X2+) - E(X+)
○ 3rd Ionization Energy (3rd I.E.): E(X3+) - E(X2+)
○ Tip: Sequential ionization energy trends according to atomic number parallel the trend of 1st ionization energy
○ Measurement: Franck-Hertz apparatus
Figure 7. Measurement of ionization energy through the Franck-Hertz apparatus
○ Situation: Low-energy electrons are accelerated between the electrodes inside a tube filled with gas.
○ Result: Accelerated electrons can transfer energy to gas particles through collisions.
○ Interpretation: Accelerated electrons can lose a specific amount of kinetic energy (4.9 eV) through collisions with gas particles, and this value can be considered the first ionization energy of gas particles.
⑥ Electron Affinity (EA): The energy released when one mole of gaseous atoms gains one mole of electrons.
○ Definition
○ X-(g) → X(g) + e-(g), EA = E(X) - E(X-). EA can be greater than 0 or less than 0.
○ Ionization energy of X = Electron affinity of X+
○ All ionization energies are greater than electron affinities (∵ E(X+) ≫ E(X), E(X-))
○ Trend 1. Generally increases going up and to the right on the periodic table.
○ Trend 2. Electron affinities of elements B to F in the 2nd period are exceptionally lower than those of elements in the same group in the 3rd period.
○ Electron affinities of elements with maximum single electron configurations in groups 2, 15, and 18 are very small.
○ Element with the highest electron affinity: Cl
○ O, S become O2- and S2- ions, which is unexpected as absorbing the second electron requires energy.
⑦ Electronegativity (EN): A combination of ionization energy and electron affinity concepts.
○ Mathematical expression: E.N = (IE + EA) / 2
○ Electronegativity increases with increasing atomic number in the same period.
○ Electronegativity decreases with increasing electron shells in the same group.
○ Tip: Electronegativity increases as you get closer to F.
○ Essential memorization: H(2.1), B(2.0), C(2.5), N(3.0), O(3.5), F(4.0), Cl(3.0), Br(2.8)
Figure 8. Electronegativity
3. Classical Molecular Models
⑴ Overview: Lewis Structure - Valence Shell Electron Pair Repulsion Theory - Atomic Bonding Theory - Molecular Orbital Theory
① Simplest molecular model
② 2D model
③ Localized electron model
⑵ Molecular Theory
① Molecule: The smallest unit of a substance that can exist independently in a stable state.
○ Rare for atoms to exist independently in nature.
○ Exception: Inert gases (He, Ne, Ar, Kr, Xe, Rn)
② Ion: An atom or group of atoms with a net positive or negative charge.
③ Introduction of Avogadro’s Number: Can directly count the number of carbons in a carbon crystal.
④ Law of Constant Composition: Elements combine in fixed integer ratios.
○ Discovered by Joseph Louis Proust.
⑤ Law of Multiple Proportions: Different elements combine in whole-number ratios.
○ Except when two elements form more than two compounds.
○ Established by Dalton.
⑶ Octet Rule
① Overview
○ Definition: Atoms tend to gain or lose electrons to achieve an electron configuration of eight electrons.
○ Most organic atoms, excluding H, C, N, O, X, Na, K, satisfy the octet rule.
② Reasons for the Octet Rule
○ 2nd and 3rd period atoms have only s and p orbitals, providing eight electron positions.
○ Eight electrons satisfy normal wave conditions, leading to more stable molecules.
③ Exception 1. Radicals: Chemical species with unpaired electrons. Molecules with two radicals are called biradicals.
④ Exception 2. Expanded Valence Shells
○ Transition elements
○ Temporary excitation in 3rd period atoms → electrons placed in vacant 3d orbitals → varied valence shell electrons → expanded octet rule
⑤ Exception 3. Incomplete Octet
○ BF3:BF3 exhibits ionic bonding where B has a 3+ charge and F has a -1 charge.
⑷ Intramolecular Forces (Chemical Bonds)
① Bond: Interaction between two atoms leading to stability.
Figure 9. Principles of Chemical Bonding
○ Principle 1. Energy decreases as two atoms approach each other (stability).
○ Principle 2. Energy increases significantly when atoms get closer than the bonding point (instability).
○ Inert gases exist as single-atom molecules due to satisfying the octet rule.
○ Inert gases usually do not participate in bonding (exceptions exist).
② Ionic Bonding: Occurs when the electronegativity difference between two atoms is greater than 1.9.
○ Metal atoms lose electrons to become cations, nonmetal atoms gain electrons to become anions, held together by electrostatic forces.
○ Examples: NaCl, LiCl, KF
③ Covalent Bonding: Bonding where two nonmetal atoms share electron pairs to form a bond.
④ Metallic Bonding: Metal atoms form bonds by sharing a “sea” of delocalized electrons, stabilizing the structure.
⑤ Bond Energy (Bond Dissociation Energy, BDE)
○ Definition: Energy required to break a bond between A and B, stronger bonds have higher BDE (more stable).
○ For ionic bonds: Energy required to separate ions in an ionic crystal lattice.
○ For covalent bonds: Energy required to break covalent bonds and produce radicals.
○ Trends
○ As the electronegativity difference between bonding atoms increases, bond energy increases.
○ Example 1: HF > HCl > HBr > HI: Can be explained by electronegativity difference or size difference.
○ Example 2: F2 < Cl2 < Br2 < I2: Can be explained by electron repulsion; F2 has more repulsion and higher reactivity.
○ Bond Comparison
○ Generally: Ionic bond > Covalent bond > Metallic bond ( controversial )
○ In aqueous environments like living organisms: Covalent bond > Ionic bond > Metallic bond, influenced by solvent effects.
○ Boiling and melting points comparison
○ Ionic bond, metallic bond > Covalent bond
⑸ Covalent Bonding
① Unbonded atom electrons exist as nonbonding electron pairs.
② Covalent bonds categorized based on electronegativity difference
○ Polar covalent bond: Electronegativity difference between two atoms is 0.5 to less than 1.9 (0.5 is not a strict number).
○ δ+ and δ- indicate stronger attraction and weaker attraction, respectively.
○ Polar covalent bonds create dipole moments due to opposing charges.
○ Nonpolar covalent bond: Electronegativity difference between two atoms is less than 0.5.
○ Ionic bonding can also be seen as polar covalent bonding with an electronegativity difference of 1.9 (1.7 to 2.0).
○ Implies that covalent bonding principles can apply to metallic bonding.
③ Covalent bond types based on the number of shared electrons: Single bond, double bond, triple bond.
④ Coordinate Bonding: An atom provides electron pairs to another atom’s vacant orbital for shared bonding.
○ Examples: H+ + H2O → H3O+
○ Examples: H+ + NH3 → NH4+
⑹ Intermolecular Forces
① Intermolecular forces are divided into van der Waals forces and hydrogen bonds, which are weaker than chemical bonds.
② Van der Waals forces are categorized as Keesom force, Debye force, London force, and repulsion force.
③ Keesom Force: Permanent dipole-permanent dipole interaction between polar molecules.
○ Attraction between δ+ and δ- in polar molecules.
④ Debye Force: Permanent dipole-induced dipole interaction between polar and nonpolar molecules.
○ Polar molecules induce polarization in nearby nonpolar molecules.
⑤ London forces or dispersion forces: These are forces of attraction between instantaneous dipoles caused by the temporary polarization (fluctuation) of electrons.
○ They are very weak forces that arise due to temporary dipoles (induced dipoles) created by the polarization of electrons.
○ These forces are the only type of attraction found in nonpolar compounds.
○ They operate in all molecules and are stronger when the atomic radius and surface area are larger (for example, the van der Waals forces in I2 are much stronger than those in F2).
○ They can explain the liquefaction of noble gases.
⑥ Repulsion Force: Nucleus-nucleus repulsion when molecules get too close.
○ Trend 1. Repulsion increases as atomic number increases in the same period.
○ Result of increased electron shielding due to higher electron count.
○ Trend 2. Repulsion increases as atomic number increases in the same group.
○ Result of decreased electron shielding due to lower electron count.
⑦ Hydrogen Bonding: Strong attraction between δ+ in H bonded to F, O, N, and δ- in F, O, N.
○ Hydrogen bonding is one of the van der Waals forces, but it should not be equated with other dipole-dipole interactions.
○ Principle 1. Hydrogen exposes its proton due to lack of core electrons, resembling ionic bonding.
○ Principle 2. Strength of hydrogen bonds
Figure 10. Boiling points based on electronegativity of atoms bonded to hydrogen
○ Boiling points are significantly affected by hydrogen bond strength.
○ Factors of hydrogen bond strength: Electronegativity of atoms bonded to hydrogen, number of hydrogen bonds.
○ Number of hydrogen bonds: The number of non-shared electron pairs of atoms with high electronegativity and the number of hydrogens within a molecule are important[Note:6]
○ H2O: O has two non-shared electron pairs. There are 2 hydrogens in the molecule, so there are 4 hydrogen bonds per molecule.
○ HF: F has three non-shared electron pairs. There’s 1 hydrogen in the molecule (limiting factor), so there are 2 hydrogen bonds per molecule.
○ NH3: N has one non-shared electron pair (limiting factor). There are 3 hydrogens in the molecule, so there are 2 hydrogen bonds per molecule.
○ Urea: Forms 6 hydrogen bonds per molecule.
○ Specific strength of hydrogen bonding:
○ F-H ······ :F = 39 kcal/mol
○ O-H ······ :N = 7 kcal/mol
○ O-H ······ :O = 5 kcal/mol
○ N-H ······ :N = 3 kcal/mol
○ N-H ······ :O = 2 kcal/mol
○ Can form multimers in liquids and solids
⑧ Relative strength: Covalent bond, Ionic bond > Metallic bond > Hydrogen bond > Dipole-dipole force ≒ Dispersion force
○ Ion-ion bond strength: 250 kJ/mol
○ Ion-dipole bond strength: 15 kJ/mol
○ Permanent dipole-permanent dipole bond strength (stationary condition): 2 kJ/mol
○ Permanent dipole-permanent dipole bond strength (rotating condition): 0.3 kJ/mol
○ Permanent dipole-induced dipole bond strength: 2 kJ/mol
○ Dispersion force: 2 kJ/mol
○ Hydrogen bond: 20 kJ/mol
⑨ Force-distance relationship:
○ Ion-ion
○ Ion-dipole: Hydration reaction is typical
○ Induced dipole-dipole (stationary)
○ Induced dipole-dipole (rotating)
○ Dipole-induced dipole
○ Induced dipole-induced dipole
○ Example: In the case of BaCl2, since Z is large, it is very stable and doesn’t hydrate easily.
⑩ If the bond within a molecule is strong, the bond between molecules is weak.
○ Example: If the hydrogen bond within a molecule is strong, the hydrogen bond between molecules is weak.
⑺ Phenomena due to water’s hydrogen bonding:
③ Increase in volume upon condensation
○ Reason why ice floats on water
○ A single water molecule can form 4 hydrogen bonds, so ice forms a hexagonal crystal structure.
4. Drawing molecular structures
① Lewis structure and Kekulé structure:
○ Lewis structure: Dot structure
○ Kekulé structure: Line structure that simply represents a pair of covalent bond electrons
○ Usually, both structures are combined and referred to as the Lewis structure.
② Drawing:
○ 1st. Draw the basic skeleton of the compound from the molecular formula
○ Generally, the central atom has a low electronegativity, and the surrounding atoms have high electronegativity.
○ Elements in the 2nd period should satisfy the octet rule (exceptions: Be, B)
○ 2nd. Count the valence electrons of all atoms
○ 3rd. Connect single bonds between atoms (each bond consumes 2 electrons)
○ 4th. Indicate non-shared electron pairs on all atoms
○ 5th. Indicate charges considering the formal charge
③ Limitations:
○ Cannot understand the bond shape of orbitals: Emergence of 3D structures and VSEPR theory
○ Cannot distinguish between sigma and pi bonds: Orbital theory emerges
○ Limitations in explaining bond lengths
○ Actual structures that cannot be explained with one structure: Represented by resonance structures
⑵ Bond-line formula or skeletal formula:
① Introduced because representing all organic compounds with Lewis structures was cumbersome
② 1st. Carbon atoms are generally not shown, but sometimes indicated for emphasis
③ 2nd. Hydrogens bonded to carbon are not shown; since carbon always has a valency of 4, the number of hydrogens can be inferred
⑶ 3D structures:
① Definition: A method to represent structures in three dimensions
② Bonds protruding from the surface are indicated with wedges, and bonds going into the surface are indicated with dashed lines
③ Bonds on the surface are usually indicated with straight lines
5. Formal charge
⑴ Definition: A fictitious charge that arises when a shared electron pair is evenly divided
⑵ Calculation: Valence electrons - Bonded electrons ÷ 2 - Non-shared electrons
⑶ Rules
① In nature, molecules should have a sum of formal charges equal to 0
② If multiple Lewis structures can be drawn, the structure with the smaller formal charge is more stable
③ If a formal charge arises, structures in which elements with high electronegativity have a (-) formal charge are more stable
Input: 2018.12.28 23:07