Chapter 1 Notes: The Biochemical Basis of Life

The Fundamental Chemistry of Life

Living organisms are studied from microscopic subatomic levels to macroscopic organism levels.
Organisms are uniquely adapted with specific structures and functions.
The properties of life arise from a hierarchical arrangement of chemical parts.
Matter makes up everything, including living organisms, and is composed of elements.
An element is a pure substance that cannot be broken down by ordinary chemical or physical techniques.
The smallest particle of an element is an atom.
Elements differ in their atomic structure.
Atoms bind chemically in fixed numbers and ratios to form molecules (e.g., oxygen gas O_2).
A chemical compound is a stable combination of different elements held together by chemical bonds.
Organic compounds in living organisms primarily contain carbon (C), hydrogen (H), and oxygen (O), and often nitrogen (N).
These four elements make up 96% of a living organism's weight.
Other elements (calcium (Ca), phosphorus (P), potassium (K), sulfur (S), sodium (Na), chlorine (Cl), and magnesium (Mg)) compose most of the remaining 4%; often occur as ions or in inorganic compounds.
Trace elements (e.g., iodine (I) and iron (Fe)) are required in small amounts (<0.1%).
Deficiency in trace elements can lead to health problems.

Atomic Structure

Elements consist of individual atoms, which are the smallest units retaining an element's chemical and physical properties.
An atom is composed of protons, neutrons, and electrons.
The number of protons defines the element's identity.
Protons and neutrons are located in the nucleus.
Electrons are located in the region surrounding the nucleus.
Protons have a positive charge, neutrons have no charge, and electrons have a negative charge.
An atom has no net charge because the number of protons equals the number of electrons.
The atomic number of an element is equal to the number of protons in the element (e.g. carbon has an atomic number of 6).
The mass number of an atom is the total number of protons and neutrons in the nucleus.
Electrons are not included in the mass number because their mass is negligible compared to protons and neutrons.

Isotopes and Radioisotopes

Isotopes are different forms of the same element with different atomic masses due to varying numbers of neutrons.
Isotopes of the same element behave the same in chemical reactions because they have the same number of protons and electrons.
Radioisotopes are unstable isotopes that decay, giving off particles of matter and energy (radioactivity), transforming into another element.
Example: Carbon-12 (^{12}C) has 6 neutrons, 6 electrons, and 6 protons and accounts for 99% of carbon in nature.
Carbon-13 (^{13}C) has 7 neutrons, 6 electrons, and 6 protons and is a stable isotope.
Carbon-14 (^{14}C) has 8 neutrons, 6 electrons, and 6 protons and is an unstable radioisotope of carbon, decaying into nitrogen-14 (^{14}N).
Radioactive decay occurs at a steady rate.
The rate of decay of a radioisotope is independent of chemical reactions or environmental conditions.
Radiation from decaying isotopes may damage molecules in living cells.
Radioisotopes are useful in geological and biological research for dating organic materials, rocks, and fossils.
They also have medical applications and are used to elucidate structures of unknown compounds.
Radioactive tracers are radioisotopes used to follow a specific chemical through a chemical reaction.
Radioisotopes are detectable in cells and are used in biological, chemical, and medical research.
Examples include using ^{14}C-labelled molecules to determine photosynthesis reactions and using ^{3}H (tritium) as tracers.
Radioisotopes are also used in nuclear medicine for diagnosis and treatment (e.g., iodine-131 for thyroid gland scans).

Electron Arrangements

Electron arrangement determines an atom's chemical properties, as only electrons are directly involved in chemical reactions.
The number of electrons in an atom equals the number of protons, making the atom electrically neutral.
Electrons move around the atomic nucleus in specific regions called orbitals.
An orbital is a region of space that one or two electrons can occupy; the most stable condition occurs when the orbital contains two electrons.
Electron orbitals are grouped into energy levels or energy shells numbered 1, 2, 3, etc., indicating their relative distance from the nucleus.
The lowest energy shell is closest to the nucleus and can hold a maximum of two electrons.
The second and third energy shells can hold up to eight and 18 electrons, respectively.
The first electron shell (1s orbital) is a single spherical orbital.
Atoms with more than two electrons have higher energy levels.
The second energy level consists of a 2s orbital (spherical) and three 2p orbitals (balloon-shaped).
The farther an electron is from the nucleus, the greater its energy.
In large atoms, higher-energy electrons occupy d and f orbitals with more complex shapes.
Valence electrons are the electrons in an atom’s outermost energy shell, or valence shell.
Atoms with an incomplete outermost energy shell are chemically reactive, while atoms with a completely filled outermost energy level are chemically inactive or inert.
Atoms with incomplete outer shells tend to gain, lose, or share electrons to complete their outermost shell.

Chemical Bonds

Inert elements occur naturally in single-atom form, while reactive elements combine with each other to form compounds.
Chemical bonds are stable attractions between atoms.
Four types of chemical bonds are important in biological molecules: ionic bonds, covalent bonds, and intermolecular forces.

Ionic Bonds

Ionic bonds form between atoms that have lost or gained electrons to become charged ions.
Ions of opposite charge are strongly attracted to one another, forming an ionic bond.
A positively charged ion is called a cation; a negatively charged ion is called an anion.
Ions are very strongly attracted to water molecules.
Ionic compounds tend to dissociate and dissolve in water, forming hydrated ions.

Covalent Bonds

Covalent bonds form when atoms share one or more pairs of valence electrons.
The strength of a covalent bond depends on the electronegativity of the atoms involved.
Electronegativity is the measure of an atom’s attraction for additional electrons.
The number of covalent bonds an atom can form equals the number of additional electrons needed to fill its valence shell.
Shared orbitals that form covalent bonds extend between atoms at specific angles and directions, giving covalently bound molecules distinct three-dimensional forms.
VSEPR (valence shell electron pair repulsion) theory states that valence electron pairs repel one another and move as far apart as possible.

Polar Molecules

Polar covalent bonds result from the unequal sharing of electrons between two atoms with different electronegativity.
Electronegativity is influenced by the atomic number and the distance between the valence electrons and the nucleus of an atom.
Oxygen and nitrogen form polar bonds with atoms of most other elements due to their relatively high electronegativities.
The atom that attracts valence electrons more strongly carries a partial negative charge, resulting in the other atom carrying a partial positive charge.
This non-uniform charge distribution is the polarity of the molecule.
Polar molecules attract and align themselves to other polar molecules and tend to be soluble in water.
Polar molecules, including water molecules, tend to exclude non-polar molecules, such as oils and fats.
Non-polar molecules have very low solubility in polar liquids.

Intermolecular Forces

Intermolecular forces, also known as van der Waals forces, are forces of attraction between molecules.
They influence the physical properties, such as solubility, melting point, and brittleness, of a substance.
The strength of van der Waals forces depends on the size, shape, and polarity of molecules.

Hydrogen Bonds

A hydrogen atom covalently bonded to a strongly electronegative atom in one molecule can be attracted to a strongly electronegative atom in a different molecule, resulting in a hydrogen bond.
Hydrogen bonds may form between atoms in the same or different molecules.
Hydrogen bonds are the strongest and most biologically significant form of van der Waals forces.
Water's unique properties (high heat capacity, melting and boiling points, cohesion, adhesion, surface tension) are due to hydrogen bonds between water molecules.
Trees depend on cohesion to help transport water through xylem tissue up from their roots.
The adhesion of water to the xylem cell walls of a plant helps to counteract the downward pull of gravity.
Hydrogen bonds also give water an unusually high surface tension, causing it to behave as if it were coated with an invisible film.
Weaker attractive force of hydrogen bonds makes them easier to break than covalent or ionic bonds, especially with increased temperature.

Other van der Waals Forces

Other van der Waals forces result from momentary attractions of the electrons of one molecule to the nuclei of another molecule.
These forces develop between all molecules, but they are only significant where other, stronger bonds are not prominent (e.g., between non-polar molecules).
The size and shape of a molecule influences the number and total strength of van der Waals forces of attraction:
- Larger molecules have larger forces of attraction.
- Linear molecules align more easily, resulting in stronger van der Waals forces.
Polar molecules experience attractive forces between the positive and negative ends of interacting molecules.

Chemical Reactions

All chemical reactions involve the breaking and formation of chemical bonds, changing the arrangements of atoms and ions.
Four major types of chemical reactions are common in biological processes: dehydration, hydrolysis, neutralization, and redox reactions.

Dehydration Reactions

Dehydration reactions (also called condensation reactions) consist of the removal of -OH and -H from two reactant molecules, forming water and joining the molecules.
They are commonly used by cells to assemble macromolecules from smaller molecules.

Hydrolysis Reactions

Hydrolysis reactions are the reverse of dehydration reactions, where water acts as a reactant to split or lyse a larger molecule into smaller subunits.

Neutralization Reactions

Neutralization reactions occur between acids and bases to produce salts and often water.

Redox Reactions

Redox reactions involve the transfer of electrons from one atom to another (reduction and oxidation).
Oxidation refers to the loss of electrons, while reduction refers to the gain of electrons.
Oxidation: X_{e-} \rightarrow X + e^-
Reduction: Y + e^- \rightarrow Y_{e-}
The oxidizing agent is the molecule or atom being reduced; the reducing agent is the molecule or atom being oxidized.

Water: Life’s Solvent

Living organisms depend on water for survival.
Up to 60% of human body weight is water; essential for cellular processes.
Water is a ubiquitous substance with simple structure and complex functions.
More substances dissolve in water than any other liquid solvent because of its polarity.

Properties of Water

Water has a high specific heat capacity, absorbing large amounts of thermal energy while its temperature increases moderately due to extensive hydrogen bonding.
It also has a high specific heat of vaporization, requiring large amounts of thermal energy for liquid water to become vapor, facilitating evaporation and cooling.
Solid water (ice) is less dense than liquid water due to the hydrogen bonds forming a rigid crystalline lattice.

Aqueous Solutions

Water readily surrounds polar and charged molecules and ions, forming surface coats that reduce attractions between molecules of other substances promoting separation and dissolution.
Polar or charged ions strongly attracted to water are hydrophilic, while non-polar molecules are hydrophobic.
Water dissolves thousands of solutes necessary for life within cells and blood.

Ionization and pH

Pure water consists of H2O molecules, OH^- ions, and H_3O^+ ions.
Autoionization: 2H2O \rightleftharpoons H3O^+ + OH^-
Acids have higher H3O^+ concentrations than OH^-, while bases have higher OH^- concentrations than H3O^+.
Acidity is measured using the pH scale: pH = -\log_{10}[H^+]
Neutral solution pH 7 (equal concentrations of H_3O^+ and OH^-).
Acidic solutions pH < 7; basic solutions pH > 7.
pH within cells and in the environment is crucial for optimal life functioning; changes of even 0.1 pH can drastically affect reactions.
Increasing ocean acidification caused by increased atmospheric CO_2 harms marine organisms.

Strong and Weak Acids and Bases

Strong acids and bases completely dissociate in an aqueous solution, while weak acids and bases only partially ionize.
The reaction of a weak acid and base in water is reversible.

Neutralization Reactions and Buffers

Neutralization: Acid + Base → Water + Salt
Buffers minimize pH changes by absorbing or releasing hydrogen ions.
Carbonic acid (H2CO3) buffer system maintains blood pH.
H2CO3 \rightleftharpoons HCO_3^- + H^+

The Carbon Chemistry of Life

Carbon atoms are the base of every organic molecule due to their bonding properties, allowing them to form chain and ring structures.
Carbon has four electrons in its valence shell, capable of forming four covalent bonds with other atoms.
Molecules consisting only of carbon atoms bonded to hydrogen atoms are called hydrocarbons.
Hydrocarbons Example: Methane (CH_4) one carbon atom bound to four hydrogen atoms.
The chain of carbon atoms in a biochemical molecule is the carbon skeleton of the molecule; it can be linear, branched, or form a closed ring shape.

Functional Groups

Small reactive groups of atoms that are part of larger biochemical molecules.
Functional groups (unlike non-polar hydrocarbon chains) are usually ionic or strongly polar, interacting with other molecules and introducing different types of bonding.
Polar groups (i.e. Hydroxyl (-OH) group on Ethanol) often act as “handles” on a large molecule. 4.1)

Dehydration/Hydrolysis

Dehydration removes components of a water molecule, usually during the assembly of a larger molecule from smaller subunits. Example: sugar molecules into a starch molecule.
Hydrolysis adds components of water the molecule, as they break into smaller subunits. Example: Breakdown of Starch.

Carbohydrates

Carbohydrates consist of carbon, hydrogen, and oxygen. They serve as energy sources, building materials, and for cell communication.
Monosaccharides. are the simplest form of carbohydrate, having a C:H:O ratio of 1:2:1; i.e. formula (CH2O)n, Example: Glucose, Fructose
Disaccharides: consist of two monosaccharides Examples: maltose, sucrose. Join with a glycosidic bond, the chemical shorthand for representing a glycosidic bond between a 1-carbon and a 4-carbon is 1→4.
Polysaccharides. is a chain of monosaccharides, also known as macromolecules, assembly of these form polymerization.
Examples. Plant Starches, glycogen and cellulose, assemble from hundreds or thousands of glucose units.

Lipids

Are made of hydrogen, carbon, and lesser amounts of oxygen. They are the main component of all plasma membranes.
1. Fatty Acids. A structural backbone, and have a single hydrocarbon chain with a carboxyl functional group
2. Fats. Made from two types of molecules: fatty acid and a glycerol molecule. Triglycerides are the most well-known fats, contain three fatty acid chains.
3. Phospholipids. Primary Lipids, the glycerol forms the backbone molecules, consisting of hydrophobic and hydrophilic regions known as Amphipathic molecules.

Steriods. Structurally based on a framework of four fused carbon rings, Sex Hormones (i.e Testosterone,Progesterone)
1. Waxes. Large Lipid molecules made of long, are hydrophobic non-polars. Forms a water resistance and protection (i.e Cutin).

Proteins

Is a polymer with three-dimensional structure that specify its function
All are polymers that are composed of amino acid monomers
Have a central carbon atom attached to an amino, carboxyl group, hydrogen atom is a variable side group.

Peptide Bond links many amino acids into chains of subunits that make proteins.
Has an -NH2 group at one end, called the N-terminal end, and a -COOH group at the other end, called the C-terminal end.
Polypeptides, is that is greater that 50 Amino Acids, that makes a 3 dimensional protein.
Have a structure of 4 levels.
1.Primary Structure. Unique Linear Sequence. Side not There are essential amino acids that must be obtained by our diet.
2. Secondary structure. Interactions the Backbone. Examples alpha Helix and Beta Sheet.
3. Tertiary Structure. Ranges of bonding interactions with Amino and R groups ( i.e non-polar). The intermolecular reactions.
4 Quaternary Structure. Are composed of with a final functional protein. Examples: Hydrogen, Polar, Disulfide.
Have the following Prosthetic groups, hemoglobin the four polypeptides, surrounded and held by the polypeptides.

Nucleic Acid

Forms the structure of all protein in living organisms. Two types exist ( DNA and RNA), They include Nucleotides .
Pyrimidines ( Uricil, Thymine and Cytosine)
Purines.(Adenine and Guanine)

Are polymers of the nucleotides, consists of Phosphate (linking to the next sugar), and contains Polynucleotide Chains.
DNA consists of molecules in cells. - (G) Form 3 bond with C, ( A) form 2 bonds with T.
Double helix configuration from molecule structure. Are often short linear forms, and fold in clover, or hairpin formation.

Enzymes (1.6)

Almost always speeds up a chemical reaction, almost all are proteins. Act as - catalyst that speeds up the reaction, usually catalyze in cellular reactions.
Are made of substrate, which recognize and binds to the same enzyme, (i.e, enzyme-substrate complex).
Consists of 4 mechanisms

Induced-fit - the enzye prior to biding changes its shape to become pressiese with a Substrate.
Enzyme activity. Concentration of a pH. - i.e temp and what can and will modify the activity.
Competitive inhibition: Blocks Enzyme activity- molecules bind to each other for competition
Allosteric Regulation.
4.) In negative feed back Loop - where regulatory inhibitor is the first Enzye for Inhibition to occur.