Bonding and Water: Atomic Structure, Bond Types, and Emergent Properties flashcards

Atomic structure recap: nucleus, electrons, and shells

At the center, the nucleus contains protons (positive charge) and neutrons (neutral).
Surrounding the nucleus is the electron cloud, where electrons spin rapidly.
Electrons arrange themselves in electron shells which contain orbitals.
Bohr-like simplification used in lecture:
- First electron shell: $1\text{ orbital}, 2\text{ electrons}$
- Second electron shell: $4\text{ orbitals}, 8\text{ total electrons}$
- Third electron shell: $4\text{ orbitals}, 8\text{ total electrons}$
In this model, shells are described in terms of orbitals and electrons, affecting how atoms interact chemically.

Electron interactions and electronegativity: how atoms interact

When atoms interact, it is the electrons (in the electron clouds) that collide and rearrange.
Electronegativity is the key to predicting bond type:
- Bond type depends on the difference in electronegativity between two atoms.
- If electronegativities are equal or close, electrons are shared between atoms.
- If the difference is large, electrons are pulled toward the more electronegative atom, forming ions.
Note on classroom behavior: not memorizing exact electronegativity numbers is encouraged; a chart may be provided when needed.
Equal or near-equal electronegativity leads to covalent bonding (sharing electrons).
Unequal sharing can still occur with covalent bonds (polar covalent) where electrons spend more time around the more electronegative nucleus.
If the difference is large, one atom can take an electron completely, forming ions (cation and anion).
- Example concept: sodium (Na) tends to lose an electron to achieve a filled shell; chlorine (Cl) tends to gain one.
- The species with the extra electron becomes negatively charged (anion); the one that loses becomes positively charged (cation).
Opposite charges attract: ionic bonds form due to electrostatic attraction between ions.
When electrons are shared, they can still be unevenly shared (polar covalent) or evenly shared (nonpolar covalent).

Covalent, polar covalent, and ionic bonds: what determines bond character

Covalent bond includes electron sharing between atoms.
Polar covalent bond: unequal sharing due to different electronegativities; results in partial charges:
- Slight negative on the more electronegative atom (δ−)
- Slight positive on the less electronegative atom (δ+)
- Represented using δ+ and δ−; H-bonds are not covalent bonds, but hydrogen is involved.
Nonpolar covalent bond: nearly equal sharing; no significant partial charges.
Ionic bond: large electronegativity difference; one atom effectively donates an electron to another, creating ions that attract each other.
Example: water (H–O–H) shows polar covalent bonds with oxygen pulling electrons more strongly than hydrogen.
Example: diatomic oxygen (O=O) has equal electronegativities, leading to nonpolar covalent bonding.
Example: sodium chloride (NaCl) features an ionic bond due to a large difference in electronegativity between Na and Cl.
Water as a key example: O is more electronegative than H, creating a partial negative on O and partial positive on H.

The math of electronegativity differences (illustrative)

For the O–H bond example:
- Difference: $3.4 - 2.2 = 1.2$
- This falls into the polar covalent range (unequal sharing but not full electron transfer).

Bohr model, valence shells, and bonding capacity

Bohr model helps visualize electron shells and how many electrons are needed to fill a shell.
For hydrogen: the first shell needs two electrons to be full.
Atoms tend to form bonds to fill their valence shells for stability.
When two atoms come together, their electrons can fill their valence shells via sharing or transfer.
In general, the central atom in many biological molecules tends to have four pairs of valence electrons to achieve stability (either by sharing or by having lone pairs).
Helium already has a full outer shell, so it usually doesn’t form bonds in this context.
The three-dimensional arrangement of electron pairs around a central atom leads to specific molecular geometries.

Molecular shapes and bond angles: why 3D structure matters

The distribution of electron pairs around a central atom leads to three-dimensional geometry.
For tetrahedral arrangements (four electron pairs around a central atom), the ideal bond angle is about \theta \approx 109.5^\8, but in molecules with lone pairs (like water), the angle can be reduced (water ~104.5°) due to lone-pair repulsion.
Methane (CH₄) is a classic tetrahedral molecule: carbon forms four C–H bonds with bond angles of about $104.5^\circ$ between hydrogens around carbon in the tetrahedral geometry.
The fourth electron pair in a tetrahedral arrangement points away from the other three to minimize repulsion, which explains why multiple bonds (e.g., quadruple) are not observed under natural conditions.
The geometry of a molecule determines its properties and its interactions with other molecules (structure determines function).
Simple examples:
- Methane: central carbon with four hydrogens forms a three-dimensional tetrahedron.
- Water: oxygen with two hydrogen atoms and two lone pairs forms a bent shape, not a straight line.
A useful analogy: the shape of a molecule is like the shape of a tool (screwdriver) determining how it can interact with materials; changing the shape can change function.

Hydrogen bonds: inter-molecular interactions and bio-relevance

Hydrogen bond: an interaction where a hydrogen atom covalently bonded to a highly electronegative atom (usually O, N, or F) forms an attraction with a lone pair on another electronegative atom in a different molecule (or different part of the same molecule).
In diagrams, hydrogen bonds are shown as dotted or dashed lines; covalent bonds are shown as solid lines.
Commonly discussed between water molecules: the slight positive on hydrogen of one molecule is attracted to the slight negative on oxygen of another molecule.
Hydrogen bonds are weaker than covalent bonds but crucial for structure and function in biology (e.g., proteins, DNA, and water’s properties).
In water, the hydrogen bond network leads to emergent properties such as cohesion and adhesion (described below).
In molecules like ammonia (NH₃) or alcohols, hydrogen bonding can occur between O–H or N–H groups in different molecules.
The concept of partial charges in hydrogen bonding uses δ+ and δ− notation to indicate slight charges rather than full charges.

Emergent properties of water: cohesion, adhesion, and surface tension

Emergent properties: new functions arising from interactions among parts (water-water interactions create cohesion; water interacting with surfaces creates adhesion).
Cohesion: water molecules hydrogen-bond to each other, creating a tightly connected network.
Surface tension: caused by cohesive forces at the air-water interface; allows small organisms to rest on the surface and explains why spiders can traverse water surfaces.
Adhesion: water hydrogen-bonds with polar surfaces or other molecules (e.g., paper towels), enabling capillary action and movement against gravity.
Example demonstrations in class:
- Paper towel climbing in water shows adhesion due to hydrogen bonding with the paper.
- The surface tension of water supports small insects and droplets and creates a “skin” at the liquid’s surface.
The environment can modulate interactions: saltwater vs freshwater changes ionic content, ionic strength, and can influence how molecules interact and fold in solution.

Energy, metabolism, and the conservation of matter in chemistry

Life requires energy; energy changes are governed by chemical reactions.
The distribution and movement of energy and electrons underlie biological processes like photosynthesis and respiration.
Conservation of matter: in any chemical reaction, atoms are rearranged but not created or destroyed. The total number and type of atoms on the left side (reactants) equal those on the right side (products).
Conservation of energy: energy is transformed but not created or destroyed.
Examples of chemical reactions relevant to biology:
- Combustion (general chemistry example):
- $\mathrm{CH4 + 2\,O2 \rightarrow CO2 + 2\,H2O}$
- Aerobic cellular respiration (biological energy extraction):
- Glucose oxidation roughly represented as: $\mathrm{C6H{12}O6 + 6\,O2 \rightarrow 6\,CO2 + 6\,H2O}$
- Photosynthesis (the energy source for most life on Earth):
- Simplified overall reaction: $6\,CO2 + 6\,H2O \rightarrow C6H{12}O6 + 6\,O2$
These reactions illustrate how energy carriers and redox processes facilitate the flow of energy in biology.
Fermentation briefly mentioned as an alternative metabolic pathway: produces energy without oxygen; notable types include alcohol fermentation and lactic acid fermentation, but these were not the focus here.
The three-dimensional arrangement and emergent properties of molecules enable complex biological functions (structure-function relationship). If the environment changes (e.g., salt concentration, temperature), structure and thus function can change.

Connections to biology: building toward larger biological molecules

Molecules assemble into larger structures (proteins, nucleic acids, lipids) whose shapes determine function.
Three-dimensional folding (driven by bond types, angles, and hydrogen bonding) leads to active sites, binding pockets, and macromolecular function.
The discussion of water and hydrogen bonding sets the stage for chapter 3 (water) and subsequent topics in biology (protein folding, cell membranes, etc.).
The foundational idea: structure and function are linked; understanding bond types, molecular geometry, and intermolecular forces explains how biology works at the cellular and organism level.

Quick reference: key terms and concepts from this lecture

Nucleus; protons (+); neutrons (neutral); electrons in an electron cloud.
Electron shells and orbitals: first shell 1 orbital; second shell 4 orbitals; third shell 4 orbitals (simplified model).
Electronegativity: tendency of an atom to attract electrons in a bond.
Bond types: covalent (shared electrons), polar covalent (unequal sharing; partial charges), nonpolar covalent (equal sharing), ionic (electron transfer; ions).
Partial charges in polar bonds: δ+ and δ−.
Hydrogen bond: intermolecular interaction between a hydrogen attached to an electronegative atom and a lone pair on another electronegative atom (usually O, N, F); shown as dotted lines.
Emergent properties: properties that arise from interactions, such as cohesion, adhesion, and surface tension in water.
Molecular geometry: three-dimensional arrangement of atoms; driven by electron pair repulsion; tetrahedral electron geometry with bond angles ~109.5°, bent geometry for H₂O due to lone pairs.
Conservation of matter and energy: matter is neither created nor destroyed; reactions rearrange atoms; energy transforms but is conserved.
Examples:
- Water: $\mathrm{H_2O}$ , polar covalent bonds, hydrogen bonding between molecules.
- Oxygen gas: $\mathrm{O_2}$ , nonpolar covalent bond.
- Methane: $\mathrm{CH_4}$ , tetrahedral, nonpolar covalent bonds.
- Sodium chloride: $\mathrm{NaCl}$ , ionic bond.