Biology Lecture Chapter 2: Chemical Basis of Life: Atomic Structure, Chemical Bonding, and Properties of Water

Atomic Structure and Subatomic Particles

Definition of Atoms and Matter:
- Atoms are the smallest functional units of matter.
- Matter is defined as anything that occupies space.
Subatomic Particles:
- Protons:
  - Located in the nucleus (the center of the atom).
  - Carry a positive charge ( $+\text{charge}$ ).
  - The number of protons dictates the specific chemical identity of an element.
- Neutrons:
  - Located in the nucleus alongside protons.
  - Carry no charge (neutral).
  - Elemental Exception: Hydrogen ( $H$ ) is unique among elements because its standard atom contains no neutrons in its nucleus (consisting of only $1$ proton and $1$ electron).
- Electrons:
  - Located outside the nucleus, revolving within surrounding electron shells.
  - Carry a negative charge ( $-\text{charge}$ ).
Electron Shell Configuration and the Octet Rule:
- Innermost Shell (1st Shell):
  - Positioned closest to the nucleus.
  - Holds a maximum capacity of $2$ electrons.
- Subsequent Shells (2nd and 3rd Shells):
  - The second shell holds a maximum capacity of $8$ electrons.
  - The third shell also holds up to $8$ electrons.
- The Octet Rule:
  - States that atoms are most stable when their outermost electron shell is filled with $8$ electrons (or $2$ electrons for the innermost shell).
  - The quantity of $8$ determines the type of chemical bonds an element can form and which atoms can interact with one another.
Electron Energy States and Thermal Excitation:
- Electrons residing in the innermost shell can absorb external energy, become excited, and jump to a higher, outer electron shell.
- Excited electrons on outer shells lose energy and fall back to lower shells closer to the nucleus.
- Thermal Excitation Example: A car sitting under direct sunlight on a hot summer day feels warm to the touch because solar energy excites electrons in the metallic atoms of the car hood. These electrons absorb solar energy, become excited, and then lose that excitement by releasing the absorbed energy as heat.

Organization of the Periodic Table and Elemental Metrics

Periods (Rows):
- Rows running horizontally across the periodic table are called periods.
- The period number indicates the total number of electron shells an element possesses.
  - Row $1$ elements possess $1$ shell.
  - Row $2$ elements possess $2$ shells (1st shell holds max $2$ electrons; 2nd shell holds up to $8$ electrons).
  - Row $3$ elements possess $3$ shells (with an additional outer shell holding up to $8$ electrons).
Groups (Columns):
- Columns running vertically down the periodic table are called groups.
- The group tells the number of electrons residing in an atom's outermost shell (valence electrons).
  - Hydrogen ( $H$ ): Possesses $1$ shell with $1$ electron, leaving room for $1$ additional electron to fill the shell.
  - Lithium ( $Li$ ): Contains $3$ total electrons ( $2$ in the inner shell, $1$ in the outer shell), leaving room for $7$ more electrons or readily donating its $1$ outer electron.
  - Carbon ( $C$ ): Contains $6$ total electrons ( $2$ in the innermost shell, $4$ in the second shell). Because its outer shell has room for $4$ more electrons to satisfy the octet rule, Carbon can form up to $4$ single bonds.
Atomic Number ( $Z$ ):
- The top number listed on an element's entry on the periodic table.
- Represents the total number of protons in the nucleus, defining the elemental identity.
Atomic Mass ( $A$ ):
- Represents the sum of protons plus neutrons in the nucleus: $\text{Atomic Mass} = \text{Protons} + \text{Neutrons}$
- Calculating Neutrons: $\text{Neutrons} = \text{Atomic Mass} - \text{Atomic Number}$
- Helium ( $He$ ) Calculation:
  - Atomic Mass = $4$ , Atomic Number = $2$
  - $\text{Neutrons} = 4 - 2 = 2$
  - Helium contains $2$ protons, $2$ neutrons, and $2$ electrons.
- Carbon-12 ( $^{12}C$ ) Calculation:
  - Atomic Mass = $12$ , Atomic Number = $6$
  - $\text{Neutrons} = 12 - 6 = 6$
  - Carbon-12 contains $6$ protons, $6$ neutrons, and $6$ electrons.
- Carbon-14 ( $^{14}C$ ) Calculation:
  - Atomic Mass = $14$ , Atomic Number = $6$
  - $\text{Neutrons} = 14 - 6 = 8$
  - Carbon-14 contains $6$ protons, $8$ neutrons, and $6$ electrons.

Isotopes, Radioisotopes, and Practical Applications

Isotopes and Radioisotopes:
- Isotopes are forms of an element that maintain the same number of protons (same atomic number) but differ in their number of neutrons (different atomic mass).
- Radioisotopes are unstable, slightly radioactive forms of elements that spontaneously decay, emitting subatomic particles and energy over time.
- Half-Life: The specific duration of time required for $50\%$ of a radioisotope sample to undergo decay.
Applications of Radioisotopes:
- Radiocarbon Dating ( $^{14}C$ ):
  - Carbon-14 ( $^{14}C$ ) possesses a long half-life, making it useful for measuring the decay rate in organic remains to determine the age of ancient biological artifacts such as fossils or mummies.
- Positron Emission Tomography (PET Scans):
  - Patients are administered a special form of glucose tagged with a radioactive isotope, such as radioactive Fluorine.
  - Cancer cells possess a significantly higher metabolic rate than surrounding healthy tissue, metabolizing glucose much faster.
  - On a PET scan image, malignant tumor sites absorb larger quantities of radioactive glucose, lighting up brighter in red or orange colors.

Chemical Elements of Life and Compound Formation

Abundant Biological Elements:
- Four main elements constitute the vast majority of living organisms:
  1. Hydrogen ( $H$ )
  2. Oxygen ( $O$ )
  3. Carbon ( $C$ )
  4. Nitrogen ( $N$ )
- Trace Elements:
  - Present in living systems in extremely small quantities.
  - Selenium Example: Selenium ( $Se$ ) is essential for health in trace amounts, but becomes highly toxic in elevated concentrations.
Molecules, Compounds, and Emergent Properties:
- Molecule: Formed when two or more atoms are bonded together chemically.
- Compound Molecule: A molecule composed of two or more distinct chemical elements.
- Emergent Properties:
  - Refers to new properties that arise when distinct elements combine, forming a compound with unique characteristics completely different from its individual raw components.
  - Sodium Chloride Example: Pure elemental Sodium ( $Na$ ) is a soft, metallic solid (resembling a block of salami) that is violently reactive. Pure Chlorine ( $Cl_2$ ) is a highly lethal, toxic gas. When combined in a $1:1$ ratio, they yield Sodium Chloride ( $NaCl$ ), ordinary table salt, an edible and non-toxic compound essential for life.

Electronegativity and Types of Chemical Bonds

Electronegativity Principles:
- Electronegativity measures the attraction or pull an atom exerts on shared electrons within a chemical bond.
- Periodic Table Trends:
  - Electronegativity increases moving from left to right across a period.
  - Electronegativity decreases moving from top to bottom down a group.
  - Elements in the upper-right corner (such as Oxygen and Fluorine) possess the highest electronegativity values and strongest electron pulls.
- Noble Gases (Helium, Neon, Argon):
  - Located in the far right column of the periodic table.
  - Possess completely filled outer electron shells ( $2$ or $8$ electrons).
  - Highly stable, unreactive, and do not donate, accept, or share electrons.
Bonding Capacity across Groups:
- Hydrogen ( $H$ ): Valence shell holds $1$ electron (capacity $2$ ); forms $1$ single bond.
- Oxygen ( $O$ ): Valence shell holds $6$ electrons (capacity $8$ ); forms up to $2$ single bonds or $1$ double bond.
- Carbon ( $C$ ): Valence shell holds $4$ electrons (capacity $8$ ); forms up to $4$ single bonds.
- Nitrogen ( $N$ ) and Phosphorus ( $P$ ): Positioned in the same group; each possesses $5$ valence electrons and can form up to $3$ single bonds.
Nonpolar Covalent Bonds:
- Formed when electrons are shared equally between two atoms with identical or nearly identical electronegativities.
- Electronegativity Difference: Less than $0.4$ ( $\Delta\text{EN} < 0.4$ ).
- Symmetrical electron density distribution.
- Molecular Oxygen ( $O_2$ ) Example: Two oxygen atoms share electron pairs equally ( $\Delta\text{EN} = 3.5 - 3.5 = 0$ ), creating a nonpolar covalent double bond ( $\text{O=O}$ ).
- Carbon-Hydrogen ( $C-H$ ) Example: Carbon ( $2.5$ ) and Hydrogen ( $2.1$ ) have a difference of $2.5 - 2.1 = 0.4$ , forming a nonpolar covalent bond.
- Methane ( $CH_4$ ) Example: Symmetrical molecular geometry with nonpolar covalent bonds.
- Covalent bonds are the strongest category of chemical bonds.
Polar Covalent Bonds:
- Formed when electrons are shared unequally between two atoms due to a moderate difference in electronegativity.
- Electronegativity Difference: Between $0.4$ and $1.8$ ( $0.4 \le \Delta\text{EN} \le 1.8$ ).
- Asymmetrical electron distribution. The more electronegative atom pulls electron density toward itself, acquiring a partial negative charge ( $\delta^-$ ). The less electronegative atom acquires a partial positive charge ( $\delta^+$ ).
- Water ( $H_2O$ ) Example: Oxygen ( $3.5$ ) and Hydrogen ( $2.1$ ) have an electronegativity difference of $3.5 - 2.1 = 1.4$ .
- In a water molecule, Oxygen holds a partial negative charge ( $\delta^-$ ), while each Hydrogen holds a partial positive charge ( $\delta^+$ ).
Ionic Bonds:
- Formed when the difference in electronegativity between two atoms is so large that electrons are completely transferred from one atom to another rather than shared.
- Electronegativity Difference: Greater than $1.8$ ( $\Delta\text{EN} > 1.8$ ).
- Sodium Chloride ( $NaCl$ ) Example: Chlorine ( $3.0$ ) and Sodium ( $0.9$ ) have a difference of $3.0 - 0.9 = 2.1$ .
- Sodium ( $Na$ ) completely donates its single 3rd shell electron to Chlorine ( $Cl$ ), which requires $1$ electron to complete its outer shell of $8$
- Cation: A positively charged ion ( $Na^+$ ) formed when an atom donates/loses an electron (the letter "t" in cation acts as a plus sign $+$ ).
- Anion: A negatively charged ion ( $Cl^-$ ) formed when an atom accepts/gains an electron.
- Opposites attract: Ionic bonds rely on full electrostatic attractions between positive cations and negative anions.
- Ionic compounds form three-dimensional salt crystals that are stable in dry environments, but dissolve readily when placed in water.

Weak Interactions: Hydrogen Bonds and Van der Waals Forces

Hydrogen Bonds:
- Relatively weak non-covalent interactions driven by electrostatic attraction between a partial positive hydrogen atom ( $\delta^+$ ) of a polar molecule and a partial negative atom ( $\delta^-$ ) of another polar molecule (such as Oxygen or Nitrogen).
- Intermolecular Interaction Example: Interaction between Water ( $H_2O$ ) and Ammonia ( $NH_3$ ), where the partial positive hydrogen ( $\delta^+$ ) of water attracts the partial negative nitrogen ( $\delta^-$ ) of ammonia.
- Energy and Bond Strength: Covalent bonds are strong and require high energy to break. Hydrogen bonds are weak and easily broken with minor energy inputs.
- Biological Functions of Hydrogen Bonds:
  1. Protein Structure: Form between amino acid backbone groups to stabilize secondary structures, such as $\alpha\text{-helices}$ and $\beta\text{-pleated sheets}$ .
  2. DNA Structure: Hold complementary nucleotide base pairs together across the central axis of the DNA double helix, allowing the two strands to be separated for replication and transcription.
Van der Waals Interactions:
- Weak electrical attractions caused by transient, temporary localized concentrations of moving electrons within molecules (governed by quantum mechanical uncertainty principles).
- Biological Application: Geckos possess millions of microscopic hair-like structures on their feet. The additive sum of Van der Waals forces between these microscopic hairs and wall surfaces enables geckos to climb smooth vertical walls.

Physical and Chemical Properties of Water

Structure Determines Function:
- Water's bent molecular geometry and polar covalent bonds produce asymmetrical partial charges ( $\delta^-$ on Oxygen, $\delta^+$ on Hydrogens), enabling water to interact dynamically with charged solutes and other polar molecules.
Hydrogen Bonding Capacity of Water:
- A single water molecule can form up to a maximum of $4$ hydrogen bonds simultaneously with adjacent water molecules.
- $2$ hydrogen bonds are accepted by the Oxygen atom ( $\delta^-$ ).
- $1$ hydrogen bond is donated by each of the $2$ Hydrogen atoms ( $\delta^+$ ).
States of Matter and Hydrogen Bond Dynamics:
- Solid State (Ice):
  - Water molecules lock into a rigid, highly ordered crystalline lattice held together by maximum, unbroken hydrogen bonds.
  - The crystal lattice pushes water molecules slightly further apart than in liquid form, making ice less dense than liquid water, allowing ice to float.
- Liquid State:
  - Thermal energy causes hydrogen bonds to continuously break, reform, and slip past one another at high speeds.
- Gaseous State (Water Vapor):
  - Elevated thermal energy completely breaks hydrogen bonds, permitting individual water molecules to disperse into the gas phase.
Water as a Universal Solvent:
- Solvent: The liquid dissolving medium (in biological systems, water serves as the universal solvent).
- Solute: The substance being dissolved (such as table salt, $NaCl$ ).
- Solution: The resulting homogeneous mixture of solvent and solute.
- Mechanism of Dissolution: Water breaks ionic crystal lattices (like $NaCl$ ) apart by forming hydration shells. The partial negative Oxygen atoms ( $\delta^-$ ) surround positive sodium cations ( $Na^+$ ), while the partial positive Hydrogen atoms ( $\delta^+$ ) surround negative chloride anions ( $Cl^-$ ).
Participation in Chemical Reactions:
- Water acts as a direct reactant or solvent medium in biological biochemical processes, including cellular respiration and the enzymatic breakdown of nutrients.