l2 Atoms, Compounds, and Chemical Bonding

Synopsis of Lecture 2

• This lecture focuses on atoms, compounds, and chemical bonding
• The structure of atoms
• Comparing different elements - core properties of the periodic table
• How atoms interact – Part I: ionic interactions
• How atoms interact – Part II: covalent interactions
• Molecular orbitals
• Features of chemical bonds

Learning Outcomes from Lecture 2

• Describe the structure of an atom in terms of orbitals
• Compare the electronic properties of different elements
• Sketch simple MO diagrams and assess the resulting bonding interactions
• Compare and contrast different chemical bonds
• Explain how delocalisation stabilises molecular structures

How My Lectures Are Built

• Lectures center on core concepts applied to explain diverse chemical properties
• Lectures provide core concepts plus background ideas and applications
• Explain lecture examples using core concepts
• Apply these concepts to similar chemical compounds
• Lectures use chemistry-specific terminology
• Correct terminology helps describe chemical ideas and communicate them
• Material in the first two lectures can be found in chapters 2, 3 and 11 of Chemistry for the Biosciences

Why Consider the Structure of Atoms

• Boyle proposed in "The Sceptical Chymist" (1661) that matter consists of particle clusters
• Avogadro introduced the term "molecule" in 1811, defining it as a group of atoms united by attraction
• Perrin proved the existence of molecules in the early 20th century
• Quantum mechanics predicts chemical properties: Linus Pauling showed that it accurately deduces molecular structures

Atomic Structure

• A positively charged nucleus is surrounded by electrons
• Positive and negative interaction enables structure
• Nucleus contains neutrons & protons; proton number determines the element (atomic number)
• Neutron variance causes isotopes
• The number of electrons is equal to protons

The Quantum Mechanics Picture

• An electron is described by four quantum numbers
• No two electrons share all four identical quantum numbers simultaneously (Pauli Exclusion Principle)
• Electrons with same quantum numbers behave similarly
• The Principal quantum number, n, describes the electron shell and distance from nucleus starting from 1
• The Azimuthal quantum number, l, indicates orbital type; ranges 0 to n-1; electrons sharing are in a subshell
• Magnetic quantum number (ml) specifies orbital within a subshell; values range -l to +l, creating 2l+1 orbitals per subshell
• Spin magnetic quantum number (ms) specifies electron spin orientation (±½), allowing two electrons per orbital

Orbitals

• Orbitals solve Schrödinger's equation, assuming independent electrons
• Orbitals are functions describing electron behavior
• Plotting orbitals visualizes electron location probability density: convention shows orbitals containing 90% probability density
• s orbitals (l=0) are spherical: one s-orbital per subshell
• p orbitals (l=1) are dumbbell-shaped: three per subshell, oriented along x, y, z directions
• d orbitals (l=2) have complex shapes: play important roles in metals

Important Features of Orbitals

• Orbital lobes mark high electron location probability
• Different colors indicate two wavefunction phases
• Probability relies on the squared value
• Nodal planes mark zero electron location probability
• Nodal planes separate the phases
• Nodal planes are important in molecular orbital formation

Orbitals and Electron Configuration

• Aufbau principle dictates electrons fill orbitals starting with lowest energy subshell
• Approximate energy order follows the n+l rule
• Electrons fill subshells with lowest n+l value
• For equal n+l values, subshell with lower n fills first
• Considering subshells provides the energy order as 1s, 2s, 2p, 3s, 3p, 4s, 3d
• We can fill the subshells to obtain the electron configuration

Examples of Electron Configuration

• Oxygen (8 electrons): 1s² 2s² 2p⁴
• Magnesium (12 electrons): 1s² 2s² 2p⁶ 3s² = [Ne] 3s²
• Copper (29 electrons): [Ar] 3d¹⁰ 4s¹

Distributing Electrons in Orbitals

• Hund's first rule says degenerate orbitals are partially filled (one electron per orbital) before adding a second electron
• This rule results in low-energy, spin-parallel conformations

Core and Valence Electrons

• Valence electrons take part in chemical bonding
• Core electrons reside in inner shells, valence electrons in outermost shells
• Atoms require a filled outer shell, and achieve this by forming bonds

Valency

• There are four electrons needed to fill the second shell
• Forming four chemical bonds helps C share and fill its shell
• Example can be seen in methane: four hydrogens bind to one carbon
• Valency (IUPAC): The maximal number of univalent atoms that may combine with an element's atom or a fragment for it to be substituted

Periodic Table: Effective Charge

• Effective nuclear charge (Zeff) is the charge experienced by outermost electrons from the nucleus
• Electrons in core shells shield electrons, the valence electrons do not shield each other well
• Charge experienced by electrons varies among elements, Zeff increases along periods

Periodic Table: Examples Of C and O

• Oxygen's electronic configuration is 1s² 2s² 2p⁴
• Carbon's configuration is 1s² 2s² 2p²
• Oxygen has 8 protons
• Carbon has 6
• Valence electrons in the second shell
• p electrons don't shield each other well, so oxygen feels more charge and is pulled closer
• Radii: carbon 0.77 Å, oxygen 0.66 Å (predicted effective charges: 3.25, 4.55)

Periodic Table: Covalent Radius

• Radius is half the length of symmetric, homonuclear bonds
• Concentrate on atomic radii: key for measure ionic radii, they determine bonding
• Atomic radii decrease across a period and increase a group

How Atoms Interact: Attraction of Nuclei

• Periodic table's effective charge changes cause varied electron "pull" by atoms
• When proximity occurs with one atom with strong poll, and the other with a weaker pull results in atoms pulling, leading to polarization to stronger atom
• Electronegativity is the tendency of an atom to attract electrons shared within chemical bonds
• The effective charge and valence electrons' nucleus distance influences electronegativity

Differences in Electronegativity:

• When differences are less than 0.7
  • Both atoms have equal strength, they attract shared electrons to create covalent bonding
• When intermediate differences occur
  • Electrons still require to be shared, yet there is a tendency to polarize shared electrons
  • Polarized bonds are created
• Differences than 1.7 results in
  • One atom pulling, and sharing electrons away to create ions

Ionic Interactions and Bonds

• Ionic bonds form due to large electronegativity differences
• Result in very stable bonds due to charge interactions. Sodium chloride crystals serve as an example
• Ionic interactions occur between charged groups/ions such as biomolecular systems
• Interactions aren't pure ionic bonds, still produce strong interactions
• These interactions occur beside hydrogen bonding interactions
• Interactions occur with charged/polarized groups are determined by Coulomb’s law, the strength in the distance dependence
• Example: Salt Bridges

Covalent Interactions

• LCAO – Linear Combination of Atomic Orbitals
• Add and subtract atomic orbitals to form molecular orbitals (MOs)
• In-phase and out-of-phase combination creates both, termed bonding and antibonding molecular orbitals

Orbital Symmetry and Bonding

• The importance is determined by symmetry
• Interactions are based on constructive, in-phase of atomic orbitals
• The interactions lower the energy of electrons; bond atoms with shared electron density.
• Only orbitals forming constructive overlap is based on symmetry, bonding occurs

Formation of Molecular Orbitals

• Anti-bonding orbitals form via out-of-phase orbital overlap.
• The energy destabilizing anti-bonding MOs
• Anti-bonding MOs become more destabilized than bonding MOs are stabilized.
• Bonding orbitals form via in phase overlap
• Energy stabilizes bonding MOs.

Two Examples of the Formation of Molecular Orbitals

• AOs are partially filled: Only bonding MOs get filled creating stable bonds (e.g., H₂)
• AOs are filled: Both MOs get filled creating unstable bond (e.g. He₂)

Orbital Size and Energy Matching

• Size impacts energy: atomic orbitals overlap better and produce lower energy MOs if their size matches
• Example: Two 2s orbital will overlap, more than a 2s and 3s
• Smaller atomic orbitals overlap better- Increased efficiency
• Example: Two 2s orbitals interact stronger than two 3s orbitals.

Energy Matching

• When interacting, two atomic oribitals form one bonding, and the other anti-bonding molecular orbital
• Bonding orbital is lower in energy than the two
• Anti-bonding orbital is higher than the two AOs; closer two AOs are in energy = stronger

MO Diagram for Hydrogen

• Use relevant AOs to match the orbital
• Must have correct symmetry
• Add the valence electrons
• Add the valence electrons from the bottom, only fill the max - Max of 2 electons

MO Diagram for Oxygen

• Use relevant AOs; those combined for bonding + antibonding MOs
• Add valence electrons to a orbital
• Make sure to select 1 electrical
• The bond order shows signifance of bonding
• This difines it as: half the number of diff in bond
• 8 in bond and 4 in anti-bonf = and therefore the bond order # 2

Core Concepts for MO Formation

• Molecular orbitals (MOs) are linear combinations of atomic orbitals (LCAO approach) formed in-phase or out-of-phase
• MOs form, same goes for the nmber of AOs combined, e.g. Combining 2 1s = one bonding and one anti-bonding MO
• Atomic orbitals form constructive overlap if symmetry allows for it
• Size that matters to energy (AOs) Determines how strong the resulting energy

Lewis Structures

• Chemical bonding for main group elements, so, aims to show electrons of nearby noble gas
• Knowns as the octet
• Each indiviual bond can be seen with picture
• Duplets: two + six e from octet
• Show bonds between atoms: - Show lines - Show line electrons - Pairs of Electrons with Dots
• Charges MUST be added
• Structures aren't uniquly difned
• The structures lead to havng resonance

Types of Molecular Orbitals

• σ molecular orbitals result from head-on AO overlap, which makes it strong
• Bonds == single bonds
• There isn't a nodal plane coincides around the atom
• π molecular orbitals form by side-on AO overlap and is a weak band
• Bond == double and triple bones
• There is nodal palne where the plane is found

Understanding Polar Bonds: General Observations

• Different atoms lead to diff charge
• Relatesd to shild from subshells due to same shell
• Forces of diff energy for diff orbitals - Nucleious with elec
• ither orbitals in 0 will be ower
• Hard to compare periodic and predict energy of h v 0

###Polarisation of Bonds: MO Explanation

• Orbital energy varies with effective charge
• Bonding MO (from two differing AOs) resembles lower E AO
• Bonding MO has increased electron density near corresponding atom, polarizing the bond.

Electrons: Electronegativity

• Difference in charge = atim with dif force strength is called Electro
• Different in electrons are pulled at same
• Therefore, band is polar, so and so.
• Help predict. The picture helps expand

Delocalised Bonds

• Examples focused electrons in bonds that are localised between two atoms
• Bonds have electron that may be spread among atoms = stablises bond
• There are difficulties in Representing a delocalised electron

Desocalisation

• Occurs for multple orbts that are able to be comboned
• These electrons ofen form π system

Structures of Bonds

• Set of structure that are limited to molecule
• Can't make structure a diagram

Group Exaples

• Example for delocisation:
• the bond and 05 o atoms
• the reultung mos of the 4 electrons

The MO of the group, the A0 is compined a obtain mol, the bond is o-a bnd o-obrl the charges the is on 0 and the orb is de