Chemistry and Chemical Bonds for Biology

Lab Preparations and Course Outlook

• Upcoming labs will utilize vocabulary and concepts covered in lectures and readings.
• Prior knowledge will enable faster completion of lab activities, allowing students to finish and depart.
• The course will proceed continuously into future topics.

Introduction to Basic Chemistry for Biology

• A foundational understanding of chemistry is essential for any conversation in biology.
• The main chemistry and cell biology content will be covered in "Principle Two" next semester.
• This section assumes no prior chemistry knowledge and builds concepts from the ground up.

Chemical Bonds: The Basics

• Definition: A chemical bond is an abstract concept representing the force that holds atoms together within a molecule.
• Formation Rationale: Most individual atoms lack both electrical neutrality and full outer electron shells, making them unstable.
  • By forming molecules, atoms can achieve both electrical neutrality and full outer shells, leading to a more stable structure than their constituent atoms alone.

Ionic Bonds

• Mechanism: One atom donates an electron to another atom.
  • One atom gives, the other receives.
  • This process creates two oppositely charged ions.
• Attraction: These ions then stick together due to electrostatic attraction.
• Example: Table Salt (Sodium Chloride, $ext{NaCl}$)
  • A sodium ($ext{Na}$) atom easily gives up its lone outer electron.
  • A chlorine ($ext{Cl}$) atom easily receives an extra electron into its nearly full outer shell.
  • Process: When a $ext{Na}$ atom meets a $ext{Cl}$ atom, sodium donates an electron to chlorine.
  • Result: This creates a positively charged sodium ion ($ext{Na}^+$) and a negatively charged chloride ion ($ext{Cl}^-$).
  • $ext{Na}^+$ and $ext{Cl}^-$ ions are then attracted to each other, forming $ext{NaCl}$.

Covalent Bonds

• Mechanism: Atoms share electrons with other atoms.
  • This electron sharing results in molecules where all participating atoms achieve full outer electron shells.
• Contrast with Ionic Bonds: Unlike ionic bonds where electrons are transferred, covalent bonds involve sharing.
• Characteristics: Covalent bonds are often very strong.
• Conditions for Formation: These bonds typically form when atoms have outer electron shells that are neither almost empty nor almost full.
• Example 1: Oxygen ($ext{O}_2$)
  • Atomic Number: Oxygen has an atomic number of $8$, meaning $8$ protons and $8$ electrons in a neutral atom.
  • Electron Configuration: $2$ electrons are in the first shell, and $6$ electrons are in the second shell.
  • Shell Status: The second (outer) shell needs $2$ more electrons to be full ($8$ total). It is not nearly empty or nearly full, making ion formation difficult.
  • Molecule Formation: Two oxygen atoms can come together and share two pairs of electrons (a double covalent bond) to form an $ext{O}_2$ molecule.
  • Stability: The $ext{O}_2$ molecule has full shells for both oxygen atoms and is therefore very stable and hard to break apart. It is the form of oxygen commonly found in air ($21$ of the atmosphere).
• Example 2: Nitrogen ($ext{N}_2$)
  • Atomic Number: Nitrogen has an atomic number of $7$, meaning $7$ protons and $7$ electrons.
  • Electron Configuration: $2$ electrons are in the first shell, and $5$ electrons are in the second shell.
  • Shell Status: The outer shell needs $3$ more electrons to be full.
  • Molecule Formation: Two nitrogen atoms can share three pairs of electrons (a triple covalent bond) to form an $ext{N}_2$ molecule.
  • Stability: $ext{N}_2$ is extremely stable due to its three covalent bonds and is very difficult to break apart. It constitutes approximately $78$ of the Earth's atmosphere.
  • Biological Implication: Despite the abundance of nitrogen in the air, most living organisms cannot break these strong $ext{N}<em>2$ bonds to access nitrogen for their molecules. This leads to most ecosystems being nitrogen-starved, and agricultural fertilizers often contain usable forms of nitrogen (not $ext{N}</em>2$).
• Definition of Covalent Bond Unit: Each pair of shared electrons in a covalent bond is considered a single covalent bond.

Representing Molecules: Structural Formulas

• Purpose: Structural formulas provide an easier, more abstract model to represent molecules compared to detailed electron configurations.
• Components:
  • Nuclei: Represented by atomic symbols (e.g., $ext{O}$, $ext{N}$, $ext{H}$).
    • Knowing the atomic symbol implies knowledge of the number of protons (and thus the nuclear charge).
  • Bonds: Each covalent bond (a pair of shared electrons) is represented by one line joining the atomic symbols.
• Examples:
  • $ext{O}_2$ molecule: $ext{O}= ext{O}$ (two lines for two shared pairs).
  • $ext{N}_2$ molecule: $ext{N} ext{☰} ext{N}$ (three lines for three shared pairs).
  • Water molecule ($ext{H}_2 ext{O}$): $ext{H}- ext{O}- ext{H}$ (with hydrogens offset, as in a bent shape).

Polarity of Covalent Bonds

• Nonpolar Bonds: Occur when shared electrons are equally shared between two atoms.
  • Typically seen when the bonded atoms are identical (e.g., $ext{O}<em>2$, $ext{N}</em>2$).
• Polar Bonds (Polar Covalent Bonds): Occur when shared electrons are not equally shared.
  • Electrons spend more time closer to one atom than the other, creating partial charges.
  • These bonds are somewhat analogous to ionic bonds, acting as an intermediate state where sharing is unequal.
• Example: Water ($ext{H}_2 ext{O}$)
  • Each of the two hydrogen atoms shares a pair of electrons with the single oxygen atom.
  • However, the shared electrons spend most of their time (are "housed") near the oxygen atom.
  • Consequence: The oxygen end of the molecule becomes slightly (partially) negatively charged, and the hydrogen ends become slightly (partially) positively charged.
  • Dissociation: These polar covalent bonds in water are so much like ionic bonds that water molecules can easily dissociate (break apart) into ions: a hydrogen ion ($ext{H}^+$) and a hydroxide ion ($ext{OH}^-$).
  • In a glass of pure water, a small proportion of the water molecules are constantly dissociating and re-associating into these ions, while most remain as intact $ext{H}_2 ext{O}$ molecules.

The Continuum of Polarity in the Real World

• Revision of Simple Model: The initial model suggesting three distinct types of bonds (nonpolar covalent, polar covalent, ionic) is an oversimplification.
• Reality: Chemical bonds exist on a continuum of polarity.
  • This ranges from completely nonpolar (rare) to slightly polar, moderately polar, very polar, and finally to completely ionic.
  • The degree of polarity depends on the types of atoms involved in the bond.
  • Most covalent bonds exhibit at least some degree of polarity.

Polar Molecules

• Definition: A polar molecule is electrically neutral overall, but different parts of it (specifically, one end or side) carry partial positive charges, and another part carries partial negative charges.
• Formation: Polar molecules typically form when a molecule contains ionic or polar covalent bonds.
• Example: Water Molecule
  • Due to the unequal sharing of electrons, the oxygen end of a water molecule tends to be negatively charged, while the hydrogen ends tend to be positively charged.
  • Thus, a water molecule is a highly polar molecule.

Consequences of Molecular Polarity

• Intermolecular Attraction: The positively charged part of one polar molecule is attracted to the negatively charged part of another polar molecule.
  • This electrostatic attraction allows electrically neutral polar molecules to stick together.
• Solubility in Water: Substances that are polar (or contain ionic bonds) are strongly attracted to water molecules and can easily dissolve in water.
  • Example: Dissolving $ext{NaCl}$ in Water
    • A crystal of table salt ($ext{NaCl}$) consists of tightly bound $ext{Na}^+$ and $ext{Cl}^-$ ions.
    • When placed in water, the negatively charged oxygen ends of water molecules surround and attract the positively charged $ext{Na}^+$ ions.
    • Simultaneously, the positively charged hydrogen ends of water molecules surround and attract the negatively charged $ext{Cl}^-$ ions.
    • These attractions are stronger than the ionic bonds holding the salt crystal together, causing the $ext{Na}^+$ and $ext{Cl}^-$ ions to break off (dissociate) from the crystal and become surrounded (solvated) by water molecules.
    • The salt effectively dissolves and dissociates into separate ions in the water.
• **