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2.1 Atoms, Isotopes, Ions, and Molecules: The Building Blocks
1. Matter and Elements:
● Matter: Anything that occupies space and has mass. ● Elements: Pure forms of matter that cannot be broken down further by ordinary chemical reactions. There are 118 elements, with 98 occurring naturally. ○ Four common elements in living organisms: Oxygen (O), Carbon (C), Hydrogen (H), and Nitrogen (N). ○ The distribution of elements varies between living organisms and nonliving matter (e.g., the atmosphere vs. Earth's crust).
2. Atomic Structure:
● Atom: The smallest unit of an element, retaining all chemical properties of the element. ○ Composed of: ■ Protons: Positive charge, located in the nucleus. ■ Neutrons: No charge, located in the nucleus. ■ Electrons: Negative charge, orbiting the nucleus. ○ Electron Cloud: The space around the nucleus where electrons reside.
3. Atomic Properties:
● Mass and Charge: ○ Protons and neutrons have almost identical mass (1 amu). ○ Electrons have much smaller mass (~1/1800 of a proton or neutron). ○ The number of protons = atomic number, which defines the element. ○ Atomic Mass: Sum of protons and neutrons (mass number). ● Atomic Number: Defines the element based on the number of protons. ● Isotopes: Atoms of the same element with different numbers of neutrons. ○ Example: Carbon-12 (6 protons, 6 neutrons) vs Carbon-14 (6 protons, 8 neutrons).
4. Isotopes and Radioactive Decay:
● Radioactive Isotopes: Unstable isotopes that decay over time (e.g., Carbon-14). ○ Carbon Dating: Used to determine the age of biological materials by measuring the ratio of Carbon-14 to Carbon-12. ○ Half-life: Time taken for half of the original isotope to decay.
5. The Periodic Table:
● Organized by atomic number (number of protons). ● Elements are grouped by shared chemical properties. ○ Three categories: ■ Metals (left side) ■ Nonmetals (right side) ■ Metalloids (in between) ● The lanthanides and actinides are placed below the main table to save space.
● Bohr Model: Describes electrons orbiting the nucleus in defined energy levels or shells (e.g., 1n, 2n).
6. Electron Behavior and Chemical Bonds:
● Atoms bond by sharing or transferring electrons to form molecules. ○ Electron Shells: Electrons are arranged in energy levels around the nucleus, with each level having a maximum number of electrons it can hold. ○ Atoms bond to achieve stable electron configurations (e.g., full outer shells).
Electron Configuration and Stability:
● Electrons fill orbitals from the closest to the nucleus outward. ● Electrons fill orbitals of equal energy one at a time before pairing. ● The outermost electrons (valence electrons) determine an atom’s chemical behavior and stability. ● Octet rule: Atoms are most stable when they have 8 electrons in their valence shell (except for the first shell, which holds a maximum of 2 electrons).
Bohr Diagrams:
● Group 1 elements (like H, Li, Na) have 1 valence electron, which they lose to achieve stability. ● Group 14 elements (like C) have 4 valence electrons and often form covalent bonds. ● Group 17 elements (like Cl) have 7 valence electrons and tend to gain 1 electron to complete their octet. ● Group 18 elements (like He, Ne, Ar) have full valence shells and are inert (non-reactive).
Electron Orbitals and Shells:
● Electrons exist in orbitals (s, p, d, f), not simple circular paths. ● The Bohr model represents electron shells, but orbitals provide a more accurate description of electron distribution. ● The first shell (1n) has 1s orbital (2 electrons max), second shell (2n) has 2s and 2p orbitals (8 electrons max), etc. ● Orbitals increase in number and complexity with each shell, and electrons fill the lowest energy levels first.
Chemical Bonds and Molecules:
● Atoms bond to achieve a stable electron configuration (full valence shell). ● Atoms can donate, accept, or share electrons to form bonds and molecules. ● Covalent bonds: Atoms share electrons (e.g., H2O molecule where hydrogen shares electrons with oxygen).
Types of Covalent Bonds
● Polar covalent bonds : electrons are unequally shared, creating partial charges (e.g., water molecule, oxygen is more electronegative). ○ This creates a dipole (partial positive and negative charges). ○ Polar bonds enable hydrogen bonding between molecules. ● Nonpolar covalent bonds : electrons are equally shared between atoms (e.g., methane, O2). ○ Nonpolar molecules have symmetrical shapes, resulting in no overall dipole.
Hydrogen Bonds and Van Der Waals Interactions
● Hydrogen bonds : weak bonds formed between a partially positive hydrogen and a partially negative atom (usually oxygen) in another molecule. Crucial for water's properties and stabilizing DNA and protein structures. ● Van Der Waals interactions : weak forces between molecules due to electron density fluctuations, important in protein structure and function.
Career Connection: Pharmaceutical Chemist
● Pharmaceutical chemists are involved in developing drugs, testing for safety, and synthesizing compounds for better effectiveness. ○ Example: Paclitaxel (Taxol) was derived from pacific yew tree bark to treat breast cancer. ○ Some drugs have unforeseen effects (e.g., minoxidil for hair growth, originally used for high blood pressure). ○ They also detect fraudulent or ineffective drugs (e.g., Krebiozen).
These bonds and interactions are crucial for the structure and function of biological molecules and processes.
2.2 Water
Water's Importance to Life and Its Unique Properties
- Importance of Water ○ Essential for life as we know it. ○ Composes 60–70% of the human body. ○ Serves as a medium for cellular chemistry and metabolism. ○ High heat capacity, heat of vaporization, and ability to dissolve polar molecules are critical for life processes.
- Water’s Polarity
○ Water molecules are polar, with a slightly positive hydrogen and a slightly negative oxygen. ○ Hydrogen bonds form between water molecules, creating cohesion and surface tension. ○ Hydrophilic substances interact well with water; hydrophobic substances do not.
- States of Water (Gas, Liquid, Solid) ○ Liquid water is constantly forming and breaking hydrogen bonds. ○ Ice is less dense than liquid water due to its crystalline structure, which allows it to float. ○ Ice insulates water beneath, protecting life in cold environments.
- High Heat Capacity ○ Water has the highest specific heat capacity of any liquid, meaning it resists temperature change. ○ This helps maintain a stable internal temperature in organisms.
- Heat of Vaporization ○ Water requires a lot of energy to vaporize, acting as a heat sink. ○ Evaporation of sweat cools organisms, helping maintain body temperature.
- Water as a Solvent ○ Water dissolves polar molecules and ions due to its polarity. ○ Ions in water form hydration shells, keeping particles dispersed.
- Cohesion and Adhesion ○ Cohesion: Water molecules stick together (e.g., surface tension allows a needle to float). ○ Adhesion: Water molecules stick to other surfaces (e.g., capillary action helps water travel in plants).
- pH and Buffers ○ pH measures the concentration of hydrogen ions (H+) in a solution. ○ Water naturally dissociates into H+ and OH- ions, with pure water having a neutral pH of 7. ○ Acids increase H+ concentration, bases decrease it. ○ Buffers maintain pH stability by absorbing excess H+ or OH-, essential for life processes. ○ Example: The blood's buffering system (carbonic acid/bicarbonate) regulates pH.
These properties make water essential for maintaining life, regulating temperature, supporting chemical reactions, and stabilizing internal conditions in organisms.
2.3 Carbon
Macromolecules and Organic Molecules
● Hydrogen Bonds : Between functional groups, crucial for molecule folding and function (e.g., DNA base pairing, enzyme-substrate binding).
Summary
● Macromolecules are made from carbon-based structures (backbones and functional groups). ● Hydrocarbons can form chains or rings with single, double, or triple bonds influencing the shape and properties. ● Isomers are molecules with the same chemical formula but different structures, affecting their function. ● Functional Groups confer specific chemical properties, driving reactions and interactions in biological systems.
Class Notes: 1/
I. The Atom
- Electrons arranged in energy shells
- 1st shell can hold greater than 2 electrons
- 2nd and 3rd shells can hold more than 3 electrons
- The saturation of an element’s outermost shell determines the element’s valence
- The most versatile element is carbon
Atomic Number
- Defined as the number of protons an element has
- This defines the element
- Carbon has 6 protons - C^
Atomic Mass
- Protons and neutrons each have 1 mass unit
- Element’s atomic mass is number of protons and number of neutrons added together (Example: Boron has 5 protons and 5 neutrons so the atomic mass is 10)
- Expressed in grams
- 1 mole of carbon atoms = 12 grams
Variations
- Most hydrogen has no neutron
- Isotopes are atoms with +/- neutrons
- 99% of C is C-12 AND 1% is C-13 then C-
- Ions are atoms with either a positive or negative charge
- Ions participate in biological processes
Essential Elements
- 25 essential elements
- CHNOPS are the bulk elements
Chemical Bonds & Forces
- Covalent Bonds - The sharing of electrons between two elements is called a covalent bond
- Ionic Bonds
- When two atoms lie at opposite ends of the electronegativity spectrum
- Between ions of opposite charges
- Hydrogen Bonds
- Arise from polar bonds
- Polar bonds and partial chargers form when atoms lie moderately far apart on electronegativity spectrum
Background info:
- Chemical/molecular vs structural formula
- Linear vs cyclic structures
- structural formulas should accurately reflect valences for electrons which you can determine atomic #s
- C and C-H bonds often abbreviated
- Chemical reactions written as chems
3.1 Synthesis of Biological Macromolecules
Dehydration Synthesis and Hydrolysis
Dehydration Synthesis
● Definition: A reaction where monomers combine to form a polymer, releasing a
water molecule as a byproduct.
● Process:
○ Two monomers (e.g., glucose) join, forming a covalent bond.
○ One monomer releases a hydrogen (H) atom, and the other releases a
hydroxyl group (OH).
○ This results in the formation of a water molecule (H ₂ O) and a new bond
between the monomers.
● Example:
○ Two glucose molecules combine to form maltose.
○ The reaction results in the loss of a water molecule.
● Key Idea: Dehydration synthesis builds polymers by removing water.
Hydrolysis
● Definition: The process of breaking down polymers into monomers by adding water.
● Process:
○ A water molecule is inserted across the covalent bond between monomers.
○ The bond is broken, and one part of the polymer gains a hydrogen atom
(H+), while the other part gains a hydroxyl group (OH–).
● Example:
○ The disaccharide maltose is broken down into two glucose molecules.
○ Water is used to split the bond, and the components (H and OH) are added to
the respective glucose molecules.
● Key Idea: Hydrolysis breaks down polymers by adding water.
Enzymes and Catalysis
● Both dehydration synthesis and hydrolysis reactions are enzyme-catalyzed, meaning
they are sped up by specific enzymes.
● Dehydration synthesis requires energy to form new bonds.
● Hydrolysis releases energy by breaking bonds.
● Example enzymes:
○ Amylase, sucrase, lactase, maltase: Break down carbohydrates.
○ Proteases (e.g., pepsin, peptidase), hydrochloric acid: Break down proteins.
○ Lipases: Break down lipids.
Applications in the Body
● Enzymes in the digestive system hydrolyze food into smaller molecules for
absorption.
● This allows cells to absorb essential nutrients and generate energy for cellular
processes.
3.2 Carbohydrates
Molecular Structures of Carbohydrates
● Carbohydrates have the stoichiometric formula (CH2O)n, where n is the number of
carbon atoms. The ratio of carbon to hydrogen to oxygen is 1:2:1.
● Carbohydrates are classified into three types: monosaccharides, disaccharides, and
polysaccharides.
Monosaccharides
● Simple sugars; common example: glucose.
● Carbon count ranges from 3 to 7. Common types: trioses (3 C), pentoses (5 C),
hexoses (6 C).
● Aldose: Carbonyl group (C=O) at the end of the molecule.
● Ketose: Carbonyl group in the middle.
● Most monosaccharides end in -ose.
● Isomers: Same chemical formula but different structural arrangement (e.g., glucose,
galactose, fructose – all C6H12O6).
Examples:
● Glucose (C6H12O6): Important energy source in humans, used in cellular
respiration to make ATP.
● Galactose: Part of lactose (milk sugar).
● Glycogen: In animals, stores glucose, broken down during low blood glucose levels.
● Cellulose: Provides structural support in plants, not digestible by humans but used
as fiber.
Benefits of Carbohydrates
● Energy: Carbohydrates are a primary energy source (4.3 Kcal/g).
● Fiber: Found in cellulose, promotes digestion and regular bowel movement.
○ Soluble fiber: Helps regulate blood sugar and lower cholesterol.
○ Insoluble fiber: Aids digestion and prevents constipation.
● Whole Grains & Vegetables: Contribute to fullness and provide slow, steady energy
release.
● Health Implications: Reduces cholesterol and may lower risk of colon cancer.
Dietary Considerations
● Carbohydrates should be part of a balanced diet along with proteins, vitamins, and
fats.
● Some diets recommend reducing carbohydrates, but they remain an essential part of
energy metabolism.
● Registered Dietitians play an important role in advising on balanced carbohydrate
intake for health and disease management (e.g., diabetes, obesity).
3.3 Lipids
Structure of Fats and Oils:
● A fat molecule consists of glycerol (a 3-carbon alcohol) and fatty acids (long
hydrocarbon chains with a carboxyl group).
● Fatty acids can be saturated (no double bonds between carbon atoms) or
unsaturated (contain one or more double bonds).
● Triacylglycerol (or triglyceride) is the primary form of fat, where three fatty acids
are attached to glycerol via ester bonds, with water molecules released in the
process.
Saturated vs. Unsaturated Fats:
● Saturated fats (e.g., stearic acid) have only single bonds between carbon atoms and
are typically solid at room temperature.
● Unsaturated fats (e.g., oleic acid) contain one or more double bonds.
Monounsaturated fats have one double bond, while polyunsaturated fats have
multiple.
● Cis fats (common in nature) have hydrogens on the same side of the double bond,
while trans fats have hydrogens on opposite sides and are often artificially created
by hydrogenating oils.
Health Implications:
● Trans fats, created through industrial hydrogenation, are linked to an increase in
“bad” cholesterol (LDL) and are associated with heart disease. Many products now
label trans fat content due to health concerns.
● Omega-3 fatty acids (like alpha-linolenic acid) are essential, meaning the body
cannot produce them, so they must be obtained through diet. These fats are known
to reduce heart disease risks, lower blood pressure, and have anti-inflammatory
properties.
Waxes:
● Waxes are lipids made of long fatty acid chains esterified to long-chain alcohols.
They are hydrophobic and protect plants and animals from water, such as the
coating on leaves or feathers.
Phospholipids:
● Phospholipids form the primary component of cell membranes. They have a
glycerol backbone, two fatty acids, and a phosphate group. The molecule has both
hydrophobic (fatty acid tails) and hydrophilic (phosphate head) parts, leading to the
formation of a phospholipid bilayer that forms the structure of membranes.
Steroids:
● Steroids, including cholesterol and hormones like cortisol, have a fused ring
structure. Cholesterol, while often maligned, is essential for the body, serving as a
precursor to hormones, bile salts, and vitamin D.
● Secondary Structure: Local folding (α-helix, β-pleated sheet) stabilized by hydrogen
bonds.
● Tertiary Structure: 3D shape formed by interactions between R groups (e.g., ionic
bonds, hydrophobic interactions, disulfide bridges).
● Quaternary Structure: Multiple polypeptide chains interact to form a functional
protein.
Protein Folding & Denaturation
● Denaturation: Loss of protein shape due to changes in temperature, pH, or
chemicals. Can be reversible (e.g., egg white protein) or irreversible (e.g., fried egg).
● Chaperones: Proteins that assist in proper folding by preventing aggregation during
the process.
Mutations in Protein Structure
● Sickle Cell Anemia: A single amino acid change (glutamic acid to valine) in the
hemoglobin β chain alters protein shape and causes cells to sickle, leading to
reduced oxygen capacity and complications.
Protein Synthesis and Post-Translational Modifications
● Post-Translational Modifications: After translation, proteins can undergo cleavage,
phosphorylation, and the addition of chemical groups to become fully functional.
Evolutionary Significance of Proteins
● Cytochrome c: A highly conserved protein with similar sequences across species,
used to assess evolutionary relationships.
Nucleus
Function:
● Genetic command center
● DNA and RNA synthesis
● Assembly of ribosimal RNA subunits
Nucleolus
- A spherical organelle in the nucleus of a cell that produces ribosomes.
- Ribosomal RNA synthesis
Plasma membrane
- Transports protein
- Provides protection for a cell
Ribosome
- Translates the genetic code from messenger RNA (mRNA) into a sequence of
amino acids, thereby building proteins within a cell;
Endoplasmic Reticulum
- A cellular factory for synthesizing proteins and lipids, playing a critical role in
protein folding, modification, and transport, as well as storing calcium ions within
the cell
Smooth ER
- Synthesize lipids (like phospholipids and cholesterol), metabolize carbohydrates,
and detoxify drugs and other harmful substances, particularly in the liver; it also
plays a crucial role in storing and releasing calcium ions in muscle cells.
Rough ER
- Primarily functions to synthesize and process proteins that are destined to be
secreted from the cell or become part of the cell membrane, achieved by the
ribosomes attached to its surface which actively produce proteins; essentially, it
is the site of protein synthesis and initial modification within a cell
