CBC2

Overview of Cell Chemistry and Bioenergetics

Vital Force: Living organisms are fundamentally chemical systems not governed by a "Vital Force", which emphasizes the biochemical mechanisms that define life. Life fundamentally depends on a vast array of chemical reactions primarily involving carbon-based organic molecules alongside water, which is crucial for various physiological processes.

Aqueous solutions: cells are made of 70 % water, which serves as a solvent for biochemical reactions and is essential for maintaining cellular structure and function.

Cells and Their Composition: Living organisms are made of only a small selection of the 92 naturally occurring elements (4) . Cells, the basic units of life, are composed mainly of four key elements: Carbon (C), Hydrogen (H), Nitrogen (N), and Oxygen (O), accounting for roughly 96.5% of an organism's weight. In addition to these major elements, trace elements such as Magnesium (Mg), Calcium (Ca), and Iron (Fe) play essential roles in various biochemical functions, including enzyme activity and structural integrity of biomolecules.

Cell Chemistry

Acids and Bases

The pH scale is used to quantify the acidity or basicity of a solution, where lower pH values indicate stronger acids and higher pH values indicate stronger bases.

  • Acid:

    • H3O+ (hydronium ion) is the species formed when an acid donates a proton to water, thus increasing the concentration of hydronium ions in the solution. Also, a substance dissolves in water to produce hydronium ions, thereby contributing to the overall acidity of the solution.

      • Higher the H3O+, the more acidic the solution becomes, leading to a lower pH value.

      • Even in pure water, there are small concentrations of H3O+ ions present due to the self-ionization of water, which establishes a neutral pH of 7.

      • Acids are characterized as being strong or weak, depending on how readily they give up their protons to water.

        • Strong acids, such as hydrochloric acid (HCl), lose their protons quickly (favorable)

        • Weak acid, like Acetic acid, holds on to its proton more tightly when dissolved in water. Many of the acids important in the cell—such as molecules containing a carboxyl (COOH) group—are weak acids (unfavourable)

      • Acids—especially weak acids—will give up their protons more readily if the concentration of H 3O+ in solution is low and will tend to receive them back if the concentration in solution is high

  • Base:

    • Any molecule capable of accepting a proton from a water molecule is called a base.

      • Sodium hydroxide (NaOH) is basic (the term alkaline is also used) because it dissociates readily in aqueous solution to form Na + ions and OH– ions.

        • Because of this property, NaOH is called a strong

          base. More important in living cells, however, are the weak bases—those that have a weak tendency to reversibly accept a proton from water.

        • Many biologically important molecules contain an amino (NH2) group. This group is a weak base that can generate OH – by taking a proton from water: –NH 2 + H 2O → –NH3+ + OH

  • As a result…

  • H+ are often being tossed back and forth, so there is an abundance of H+ in the cell at any point in time. Weak acids contribute to a dynamic equilibrium of H+ ions, creating a balance where H+ is neither too scarce nor too abundant—perfect for the needs of the cell! so that an acid can be a base and vice versa all depending on the environment– this is extremely important to the chemistry of the cell.

  • Buffers: Substances that help maintain the pH of a solution by resisting changes in acidity or alkalinity when small amounts of acid or base are added.

    • Weak acids and bases that can release or take up protons near pH 7 keep the environment of the cell relatively constant under a variety of conditions.

Cells are made of Molecules

  • A molecule consists of one or more atoms

    bonded together

    • Atom: equal # protons & neutrons bundled

      together into a nucleus

    • (-) charged electrons orbit the nucleus

      – same #: neutral

      – more: anion (-)

      – fewer: cation (+)

  • Electron Shells (2n²)

    • Electrons occupy discreet spaces around the nucleus is called shells

    • Each shell can hold a specific maximum number of electrons

    • Innermost shell

      • Shells closest to the nucleus must fill first because they require the least amount of energy.

      • Once the first shell is full, electrons move to the next shell.

        • 1st shell (n=1): can hold 2 electrons

        • 2nd shell (n=2): can hold 8 electrons

        • 3rd shell (n=3): can hold 18 electrons

        • 4th shell (n=4): can hold 32 electrons, and so on.

    • Outermost shell

      • Called valence shells: the electrons in this shell are important because they determine how the atom will react chemically with other atoms.

      • The outermost shells of many atoms are only

        partially “filled”

      • Atoms with unfilled outer shells are inherently

        less stable than atoms with filled outer shells

    • Stability

      • Atoms with unfilled outer shells try to reduce instability

        • by finding ways to fill their outermost shell

        • Interacting with another atom is one way to get closer to a full

          outer shell

        • 2 important interactions

          – covalent bonds

          – ionic bonds

  • Ionic Bonds

    • Ionic bonds are a type of electrostatic attraction

      • Same type of attractive force as hydrogen bond but

        much stronger due to full rather than δcharge

      • very strong and pliable

Bonds and Electronegativity

  • Water is held by Hydrogen Bonds

    • When a positively charged region of one water molecule (that is, one of its H atoms) approaches a negatively charged region (that is, the O) of a second water molecule, the electrical attraction between them can result in a hydrogen bond

    • Individually Weak, Collectively Strong:

      • A single hydrogen bond is weak and can be easily broken.

      • However, when many hydrogen bonds occur together, they create a cumulative strength that stabilizes structures and interactions.

      • This "strength in numbers" makes hydrogen bonding a key player in biology.

    • Why weak Individually?

      • Hydrogen bonds rely on electrostatic attractions, which are not as strong as covalent or ionic bonds.

      • This weakness allows for dynamic interactions, such as breaking and reforming, which is crucial for processes like DNA replication, protein function, and molecular recognition.

        • These bonds are broken by the heat energy of the molecules, which increases molecular motion and disrupts the attractive forces, leading to hydrogen bonds being continually broken and formed. It is only because of the hydrogen bonds that link water molecules together that water is a liquid at room temperature—with a high boiling point and high surface tension—rather than a gas

          • hydrophilic: Water-loving

          • Hydrophobic: water-avoiding (hydrocarbon)

Covalent Bonds

atoms with nearly filled outer shells share an outer shell electron: a covalent bond forms, and a molecule is formed. Covalent bonds are relatively strong and relatively stable.

  • Covalent bonds can be “single”

    • single bonds can rotate

    • single bonds are flexible

      • a macromolecule allows rotation of the atoms they join, giving the polymer chain great flexibility. In principle, this allows a macro-

        molecule to adopt an almost unlimited number of shapes, or conformations, as random thermal energy causes the polymer chain to writhe and rotate.

    • Covalent bonds can be “double”

      • double bonds are shorter, stronger, and less flexible

    • Triple bonds exist but are uncommon in cells

      • nitrogen gas N2

      • acetylene C2H2

  • Covalent Bonds Are Not All Equal

    • The “strength” of a bond is the difference in energy between the free atoms and the energy of the molecule they form

    • Therefore, different covalent bonds have different strengths

      • stronger bonds are more stable, possess less energy e.g., CO2, H2O

    • Weaker covalent bonds are less stable

      • possess more energy

      • molecules with weaker covalent bonds can be a form of stored energy

        • e.g., CH2O

  • Electronegativity

    • Electronegativity is the tendency of an atom to attract shared electrons in a covalent bond.

    • Differences among atoms

      • Atoms like oxygen (O) and nitrogen (N) are highly electronegative, meaning they pull electrons toward themselves strongly.

      • Atoms like hydrogen (H) or carbon (C) are less electronegative, meaning they don't pull as strongly on shared electrons.

    • Unequal Sharing of Electrons

      • If the atoms have different electronegativities, the electrons are not shared equally.

        The more electronegative atom pulls the electrons closer to itself, creating an unequal distribution of electrical charge.

    • Polar Molecules

      • A polar molecule is formed when this unequal sharing creates a partial positive charge (S delta+) on one side of the molecule and a partial negative charge (S delta-) on the other

        • Example: Water (H2O)

          • Oxygen is more electronegative than hydrogen.

          • Oxygen pulls the shared electrons closer, making oxygen S delta− and the hydrogens S delta+.

          • This results in a polar molecule with an uneven charge distribution.

          • Polar molecules are very important in biological

            systems.

    • Polar Covalent Bonds

      • if the polar difference is small, then attractions can develop between atoms on to different polar molecules

        • Most important—the hydrogen bond

      • Electrostatic attraction between an electropositive hydrogen atom on one polar molecule and an electronegative atom on another are Individually weak, but many such bonds together can be strong

        • These interactions are weaker than covalent or ionic bonds but still play a vital role in stabilizing molecular structures and facilitating biological processes.

      • Additional Infomation

        • If the Polar Difference is Large...

          • The molecules behave as acids or bases

          • Acids give up the H+ in their polar bond to a base

          • In aqueous environments (i.e. the cell) the “base” that picks up this H+ is often H2O, resulting in H3O+

          • The stronger the acid, the more likely its Hδ+ will exist

            mostly as part of a hydronium ion

Cells Need “Fixed” Nitrogen and Carbon

To make a living cell requires matter, as well as free energy. DNA, RNA, and

protein are composed of just six elements: hydrogen, carbon, nitrogen, oxygen,

sulfur, and phosphorus.

  • Carbon

    • Importance: Carbon is small and has four electrons and four vacancies in its outermost shell, a carbon atom can form four covalent bonds with other atoms.

    • The carbon compounds made by cells are called organic molecules. In contrast, all other molecules, including water, are said to be inorganic.

    • Four major families: of small organic molecules: the sugars,

      the fatty acids, the nucleotides, and the amino acids.

    • Macromolecules: are the most abundant carbon-containing molecules

      in a living cell

  • C and N are common

    • but common forms (CO2 and N2) are very stable

    • not very “accessible”

  • Phototrophs “fix” carbon

    • pry it off the very stable CO2

    • make it readily available (in the form of CH2O)Carbs

  • Several species of bacteria synthesize an enzyme that catalyzes the conversion of N2 to ammonia

    • also lightning, geothermal events

Wait….