7: Molecular Geometry

Mid-Semester Feedback

• Mid-semester feedback is available.
• A QR code is provided for accessing feedback.

Lecture 7: Announcements

• Specific sections of Brown's Chemistry textbook are covered in the lecture (8.1-8.8, 9.1, 9.2).
• Problem Set 6 is due the next day; Problem Set 7 is due next week.
• Study Center hours are Wednesdays from 6 pm to 8 pm in ETA F 5.
• Professor Norris and Brisby's office hours are Thursdays from 5 pm to 6 pm, in LEE P 210.
• Next week's lecture will cover sections 11.1-11.6 of Brown's Chemistry textbook.

Review

• Electronic structure and periodic table were discussed in Lecture 6.
• Light waves, quanta of energy, and photons were discussed.
• Line spectra, atomic lines, and the Bohr model of hydrogen were covered.
• Beyond the Bohr model, quantum mechanics was discussed along with quantum numbers n, l, ml, and ms.
• Orbital shapes (s, p, d, f) and spin, Pauli Exclusion Principle, Hund's rule, multi-electron atoms, were also covered.
• Electron configurations, screening, effective nuclear charge, atomic radii, ionic radii, ionization energy, and electron affinity are also key concepts for the lecture.

Basics of Chemical Bonding

• Atoms or ions are held together by bonds in molecules and solids.
• Bonds involve valence electrons, typically those in the outermost shell.
• Types of bonding include ionic, covalent, and metallic.
• lonic bonding involves electrostatic attraction between ions.
• Covalent bonding involves sharing of valence electrons.
• Metallic bonding involves sharing of valence electrons with the entire solid.
• Examples of bonding types are provided for each example.

Lewis Symbols

• For each atom, draw the atomic symbol and add valence electrons as dots.
• Valence electrons are placed around the symbol, with a maximum of two dots per side of the symbol.
• The arrangement of dots is equivalent top, bottom, left and right, to visually identify valence electrons.

Octet Rule

• Atoms tend to gain, lose, or share electrons until they have eight valence electrons.
• This stable configuration is like that of noble gas atoms.
• Exceptions to the octet rule exist, particularly for atoms with valence electrons in higher energy levels (d and f subshells).

Ionic Bonding

• Metal atoms lose valence electrons, and nonmetal atoms gain valence electrons.
• This results in ions with opposite charges that attract each other.
• Ionic bonding is strong, due to the electrostatic attraction between the ions.
• Lattice energy quantifies the strength of ionic bonding, representing the energy required to separate ions to infinite distance..
• An example of calculation is provided in the lecture.

Calculating the Lattice Energy

• Lattice energy (ΔHlat), is calculated by using Hess's law, enthalpy of formation (ΔHf), ionisation energy, and electron affinity (EA).

Covalent Bonding

• Two atoms share their valence electrons to form a covalent bond.
• The shared electron pair is the "glue" that holds the atoms together.
• Each atom achieves a noble gas configuration by sharing electrons.
• Examples of covalent bonding are provided.
• Shared electrons are drawn as a line to symbolise a "bond".

Multiple Covalent Bonds

• Single pairs of shared electrons indicate a single covalent bond.
• Two shared pairs indicate a double covalent bond.
• Three shared pairs indicate a triple bond.
• In general, bond lengths decrease with an increasing number of shared electron pairs.

How Are Valence Electrons Shared in Covalent Bond?

• Bond polarity refers to the distribution of the bonding pair.
• Nonpolar covalent bonds share electrons equally.
• Polar covalent bonds share electrons unequally.
• Electronegativity is the property of an atom to attract shared electrons in a covalent bond toward itself.
• The electronegativity difference between atoms determines the bond polarity.
• The higher the difference, the more polar the bond.

Bond Polarity

• Bond polarity describes whether an electron pair is shared equally or unequally in a covalent bond.
• Molecules like F₂ share electrons equally, showing no polarity, whereas HF (hydrogen fluoride) shows polarity.
• Electronegativity values can be used to estimate the polarity of a bond. Values are provided in a table.

Electronegativity Table

• Electronegativity values are used to predict bond polarity.
• Values are provided in a periodic table format.

Calculate Bond Polarity?

• Electronegativity difference can be used to determine if a bond is polar covalent or ionic.
• A rule of thumb exists for electronegativity difference (≥ 2) indicating an ionic bond.
• Partial charges are used to represent polar covalent bonds (δ+ and δ-).

Dipole Moments

• Partial charges in a molecule lead to a dipole moment, a vector from negative to positive charge.
• Magnitude quantifies the separation of charges by distance.
• The unit for quantifying dipole moment is the Debye (D).

Ionic versus Covalent

• The nature of bonding falls along a continuum, not a strict line between ionic and covalent bonding.
• Generally, greater electronegativity difference leads to more ionic character (higher melting/boiling temp), lower for more covalent character.
• Oxidation number of a metal influences the bonding character.

Drawing Lewis Structures

• Rules to draw Lewis structures are provided. This allows predicting the shape, connectivity, and overall structure of molecules and ions like HCN and NH3.

Alternative Lewis Structures

• Multiple Lewis structures can satisfy the octet rule for a molecule.
• Formal charges are used to determine the most likely structure. The structure with formal charges closest to zero is favored. This is critical to accurately depict molecular structure and bonding.

Alternative Lewis Structures:

• Formal charges help determine which alternative Lewis structure is more probable. Negative formal charges tend to be on most electronegative atoms. Additional examples are given for determining most probable alternative structures.

Oxidation Number vs. Formal Charge vs. Partial Charge

• Different methods to assess formal charge using different perspectives. Oxidation numbers assume the bond is ionic, formal charges are completely shared, and partial charges account for variability between the two. Examples exist for comparing the results of calculations for the same molecules.

Resonance Structures

• Molecules can have multiple equivalent or resonance structures that accurately depict overall structure and bonding.

Resonance Structures

• Molecules can exhibit resonance, shown by more than one possible Lewis structure.
• Bond lengths within a resonant molecule can be approximately intermediate between single and multiple bonds.

What About Molecular Shape?

• Lewis structures are 2-dimensional; need 3-dimensional representations for molecules. The shape is determined by valence-shell electron-pair repulsion (VSEPR) theory. Real-world examples (e.g., tetrachloromethane (CCl4), and NH3) are presented to illustrate that structures must consider 3-dimensional bonding and geometries.

Predicting Shapes

• Shapes of simple molecules are predicted using VSEPR theory, which accounts for the interplay of repulsive forces among electron pairs in the molecules. Several molecular shapes are presented.

Why Remove Atoms?

• Valence-shell electron-pair repulsion (VSEPR) explains molecular shapes based on electron-pair repulsion; electron pairs are positioned as far apart as possible. Several shapes are presented again based on the VSEPR theory.

Repulsion of Electron Domains

• Electron repulsion leads to specific molecular shapes like linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral.

Possible Molecular Geometries

• Additional examples are given for the molecule phosgene (Cl2CO) illustrating the influence of electron domains and multiple bonds on molecular geometry.

Predicting Shape with VSEPR Model

• A step-by-step process for determining molecular shape using VSEPR is illustrated with the example, NH3 (ammonia), which has a trigonal pyramidal shape.

Possible Molecular Geometries

• Hypervalent atoms beyond period 3 on the periodic table can form molecules with more than 4 electron pairs. Examples are provided to illustrate molecules beyond period 3 on the periodic table and their 3D shapes using the VSEPR model.

What We Learned

• Summary of key concepts and techniques covered in Lecture 7.

Test your knowledge on the molecular geometry of carbon tetrachloride (CCl4), including its bond angles, bond lengths, and bonding types. Explore key concepts like the octet rule and Lewis symbols related to this tetrahedral molecule.