General Biology: Matter, Elements, and Chemical Bonding

Essential Understanding: To comprehend biology, one must first understand chemistry. Dr. David Essegg emphasizes that biological concepts are fundamentally rooted in chemical principles.
Living systems are entirely composed of matter, making chemistry the foundational language of biological study.

Definition of Matter: Matter is defined as anything that possesses mass and volume (takes up space).
States of Matter: Matter exists in three primary states in biological contexts: - Gas - Liquid - Solid
Elements: Elements are the basic substances of matter that cannot be broken down further by chemical means. - There are $92$ naturally occurring elements. - Essential Elements: Only $6$ elements predominate in living systems and are considered essential. A pie chart of elemental distribution in living organisms reveals: - Oxygen ( $O$ ): The most abundant element, comprising $61\%$ of the organism. - Carbon ( $C$ ) - Hydrogen ( $H$ ) - Nitrogen ( $N$ ) - Calcium ( $Ca$ ) - Phosphate ( $P$ ) - Trace Elements: These are elements required in smaller quantities but are nonetheless vital for survival. The term "trace" may be misleading as it does not imply insignificance, only lower abundance. Examples include: - Magnesium ( $Mg$ ) - Potassium ( $K$ ) - Sodium ( $Na$ )

Molecules: A molecule is formed when two or more elements bond together. - Example: Water ( $H_2O$ ) is formed from the combination of $2$ hydrogen atoms and $1$ oxygen atom.
Chemical Bonds: These are the attractive forces that hold same or different atoms together. There are three primary types discussed: - Ionic bonds - Covalent bonds - Hydrogen bonds

Ions: Atoms or molecules that carry a net electrical charge. - Cations: Positively charged ions ( $+$ ). Mnemonic: The "t" in cation resembles a plus sign. - Anions: Negatively charged ions ( $-$ ).
Formation of Ionic Bonds: These bonds form when ions or molecules of opposite charges are attracted to one another through electrostatic forces, creating a relatively stable association.
Case Study: Sodium Chloride ( $NaCl$ ) - In their isolation or in a watery solution, Sodium ( $Na$ ) and Chlorine ( $Cl$ ) have incomplete outer (valence) shells. - Sodium has an extra electron in its outer shell; Chlorine is missing one electron to complete its shell. - Donation Process: Sodium donates its valence electron to Chlorine. - Resulting Charges: Sodium becomes a cation ( $Na^+$ ) because it has lost an electron. Chlorine gains a negative charge as it receives the electron, becoming a Chloride anion ( $Cl^-$ ). - The resulting stability is created by the electrostatic attraction between $Na^+$ and $Cl^-$ .

Definition: Covalent bonds are strong, sturdy bonds formed when electrons are shared equally between two atoms to complete the valence shells of both.
General Rule: They typically form between two nonmetal atoms.
Bond Strength and Variety: Covalent bonds provide a strong attractive force and can exist in three levels of intensity: - Single bonds - Double bonds - Triple bonds (the strongest)
Example: Methane ( $CH_4$ ) - A single carbon atom is missing $4$ electrons in its outer shell. - Carbon binds with $4$ hydrogen atoms (each possessing a single electron). - The electrons are shared equally in four different directions. - This arrangement creates high stability, exceeding that of an ionic bond.

Definition: A variation of the covalent bond where electrons are shared unequally between atoms.
Significance: This unequal sharing creates partial negative ( $\delta^-$ ) and partial positive ( $\delta^+$ ) charges, which are critical in biological chemistry.
The Water Molecule ( $H_2O$ ): - Water functions as the primary solvent for living organisms. - Structure: Oxygen contains $8$ protons and $8$ neutrons. It forms two covalent bonds by sharing electrons with two hydrogen atoms. - Electronegativity: The large oxygen nucleus attracts the shared electrons more strongly than the hydrogen nuclei. - Charge Distribution: Electrons spend more time near the oxygen end, resulting in a weak negative charge on the oxygen and a weak positive charge on the hydrogens.

Definition: These are weak electrostatic interactions between a partial positive hydrogen atom and another atom with a partial negative charge.
Characteristics: While individually weak, they are numerous and vital for stabilizing large biological structures.
Common Polar Precursors: Pairs that often lead to hydrogen bonding include: - Oxygen and Hydrogen ( $O-H$ ) as seen in water. - Nitrogen and Hydrogen ( $N-H$ ). - Carbon and Oxygen ( $C-O$ ).
Role in Water: In liquid water, molecules form complex networks where the partial positive hydrogen of one molecule interacts with the partial negative oxygen of another, conferring water's unique liquid properties.
Role in DNA: - DNA consists of two strands. Each individual strand's backbone is held together by strong covalent bonds (specifically sugar-phosphate bonds). - The two strands are held together in a double helical (alpha helical) shape by hydrogen bonds. - These bonds occur between oxygen and hydrogen atoms attached to nitrogen in the nitrogenous bases. - In DNA, these pairings occur in groups of $2$ or $3$ bonds. - Functional Importance: The relative weakness of these bonds allows the DNA strands to be separated for essential biological processes such as transcription and replication.

Polymers and Monomers: Exploring how complex structures are built, such as the DNA strand being a polymer composed of nucleotide monomers.
Repeating Units: How monomers are joined in a repeating fashion via covalent bonds (e.g., sugar-phosphate backbone).
Biomacromolecules: Introduction to the different categories of large biological molecules.

Essential Understanding: Comprehension of biology requires knowledge of chemistry; biological concepts are rooted in chemical principles.
Nature of Matter:
- Defined as anything with mass and volume.
- Exists as gas, liquid, solid. - Essential Elements: 6 major elements in living systems: Oxygen (61%), Carbon, Hydrogen, Nitrogen, Calcium, Phosphate. - Trace Elements: Needed in small quantities (e.g., Magnesium, Potassium, Sodium).
Molecules & Bonds:
- Molecules form when elements bond (e.g., Water - $H_2O$ ).
- Bonds: Ionic, covalent, hydrogen.
Ionic Bonds: Ions carry charges; cations (+), anions (-).
- Example: Sodium Chloride ( $NaCl$ ) forms from ionic bonds.
Covalent Bonds: Strong bonds from shared electrons; can be single, double, triple (e.g., Methane - $CH_4$ ).
Polar Covalent Bonds: Unequal electron sharing creates partial charges (e.g., Water).
Hydrogen Bonds: Weak interactions stabilize biological structures; crucial in water and DNA.
Future Topics: Polymers & monomers, repeating units, introduction to biomacromolecules.