1.1 Chemical Fundamentals updated 2023

Introduction

• Quote referencing an attractive external element: Br Na+

Chemical Fundamentals

• Significance of Molecules: Molecules such as lipids, nucleic acids, proteins, and carbohydrates are fundamental to cellular structures and processes, playing critical roles in metabolism, genetic information storage and transmission, and energy production. Mastery of the chemistry of these molecules is vital for understanding biology and is encapsulated in the field termed Biochemistry.

• Review of Previous Chemistry: A reminder of essential concepts from prior education (10th and 11th grades) that form the foundation for understanding complex biochemical processes.

Isotopes

• Definition: Isotopes are atoms that share the same number of protons and electrons but differ in the number of neutrons. This results in variations in atomic mass and, in some cases, nuclear stability.

• Example: Hydrogen isotopes include Protium (1 proton, 0 neutrons), Deuterium (1 proton, 1 neutron), and Tritium (1 proton, 2 neutrons), each with distinct physical and chemical properties.

Radioisotopes

• Characteristics of Radioisotopes: Radioisotopes possess unstable nuclei that decay spontaneously through a process known as radioactive decay. This decay results in the emission of radiation, which can be harnessed in various applications.

• Half-life: The half-life of a radioisotope is the time required for half of the radioactive nuclei in a sample to decay, which is crucial for dating archaeological finds and medical applications.

• Applications: A prime example is its application in radiometric dating, such as determining the age of the Pygmy Woolly Mammoth, estimated to be 1800 years old through isotope analysis.

Chemical Behavior

• Electron Movement: Electrons move around the nucleus at varying distances, with their distance being intrinsically related to their energy level. As energy increases, electrons can occupy higher energy levels.

• Energy Levels: Electrons exist in distinct energy levels, denoted as n=1, n=2, etc., with each level associated with specific energies and locations.

• Valence Electrons: The outermost electrons, known as valence electrons, are key determiners of an atom's chemical behavior, influencing bonding preferences and reactivity.

Chemical Bonding

• Types of Bonds:

  • Intramolecular Forces (within molecules): These include ionic bonds, which result from the transfer of electrons, and covalent bonds, which involve sharing electrons.

  • Intermolecular Forces (between molecules): These include weaker forces like hydrogen bonds, hydrophobic interactions, and van der Waals forces, which are crucial in determining the properties of substances.

Ionic Bonds

• Definition: Ionic bonds form from the electrostatic attraction between oppositely charged ions (cation and anion).

• Example: In sodium chloride (NaCl), sodium (Na+) donates an electron to chlorine (Cl-) resulting in a strong ionic bond. In aqueous solutions, NaCl dissociates into free ions (Na+ and Cl-), contributing to its conductivity and solubility properties.

Covalent Bonds

• Formation: Covalent bonds occur when two atoms share pairs of valence electrons, achieving greater stability. The sharing can be equal (nonpolar covalent) or unequal (polar covalent).

• Example: A water molecule (H2O) is formed from the covalent bonds between two hydrogen atoms and one oxygen atom, resulting in a bent molecular shape and polar characteristics.

Polarity in Covalent Bonds

• Polar Covalent Bonds: These result from the unequal sharing of electrons, leading to partial positive (δ+) and negative (δ-) charges on atoms, driven by differences in electronegativity. This polarity is crucial for the interactions between different molecules.

Molecular Shapes

• VSEPR Model: The Valence Shell Electron Repulsion model provides a visual representation of molecular shapes based on repulsions between valence electron pairs.

• Examples of Shapes:

  • Methane (CH4): Tetrahedral geometry with a bond angle of approximately 109.5°.

  • Ammonia (NH3): Pyramidal shape with bond angles close to 107°, influenced by one lone pair of electrons.

  • Water (H2O): Angular shape with bond angle around 104°, significant for its polarity.

  • Hydrogen Chloride (HCl): Linear molecule due to the single bond between hydrogen and chlorine.

Molecular Polarity Criteria

• Conditions for Polarity: A molecule must possess polar covalent bonds and have an asymmetrical arrangement of those bonds to be classified as polar, enabling interactions with other polar substances.

Properties of Water

• Water as a Polar Molecule: Water exhibits unique properties due to its polarity, making it an excellent solvent for ionic and other polar compounds. Its polar nature leads to hydrogen bonding effects that significantly influence biochemical processes.

• Hydrogen Bonding: One of the strongest intermolecular forces, hydrogen bonding is critical for the unique physical properties of water, such as high specific heat, surface tension, and cohesion, all necessary for supporting life.

Density of Water

• Ice as an Insulator: Ice's lower density compared to liquid water allows it to float, providing an insulating layer that keeps water below from freezing, thus protecting aquatic life.

• Habitat for Organisms: Bodies of water serve as habitats for a range of organisms, both unicellular and multicellular, playing crucial ecological roles.

Importance of Water Solubility

• Key Role: Water is essential for life; it facilitates biochemical reactions and helps maintain cell structure by operating within watery cellular environments.

• Polar Molecules: Polar compounds (e.g., glucose) tend to be soluble in water, which is vital for cellular metabolism and nutrient transport.

• Non-Polar Molecules: Non-polar substances (e.g., octane) exhibit hydrophobic properties, making them insoluble in water, which has implications for their behavior in biological systems.

Intermolecular Forces - Hydrogen Bonding

• Definition: The hydrogen bonding phenomenon refers to weak attractions between slightly positive hydrogen atoms and slightly negative atoms (such as O, N, or F) in neighboring molecules, which significantly impact the physical properties of substances.

Intermolecular Forces - Hydrophobic Interactions

• Hydrophobic Effect: Non-polar molecules tend to cluster together in an aqueous environment due to their mutual repulsion from water, influencing the structure of biological membranes.

• Plasma Membrane: The arrangement of molecules in the plasma membrane is influenced by hydrophobic interactions, dictating the membrane's permeability and fluidity, which are essential for cellular function.

Biochemical Reactions Overview

• Types of reactions covered:

  • Neutralization reactions (covered in coming slides)

  • Condensation reactions (to be discussed in the next lesson)

  • Hydrolysis reactions (to be discussed in the next lesson)

  • Redox reactions (to be explored in Unit two).

Acids

• Definition: Acids are substances that, when dissolved in water, release hydrogen ions (H+), resulting in the formation of hydronium ions (H3O+).

• Example: When hydrogen chloride (HCl) dissolves in water, it ionizes to produce hydrochloric acid, a strong acid commonly found in gastric juice.

Bases

• Definition: Bases are substances that generate hydroxide ions (OH-) upon dissolution in water.

• Example: Sodium hydroxide (NaOH) dissociates in water to yield Na+ and OH- ions, functioning as a strong base used in various industrial applications.

Neutralization Reaction

• General Reaction: The typical reaction between an acid and a base produces water and a salt, a fundamental aspect of acid-base chemistry.

• Example: The reaction between HCl(aq) and NaOH(aq) follows the equation: HCl + NaOH → H2O + NaCl, demonstrating the formation of common table salt.

Acid-Base Buffers

• Buffers: Buffers are crucial substances that help in moderating fluctuations in pH levels, maintaining homeostasis in biological systems.

• Example: Human blood has a normal pH operating around 7.4 (acceptable range: 7.35 to 7.45), with buffers helping to stabilize this pH against dietary or metabolic fluctuations.

• Natural Buffers: These substances neutralize excess acids or bases; the bicarbonate system is a key buffer mechanism.

• Key Buffer System: The Carbonic Acid - Bicarbonate system plays a significant role in pH regulation in biological fluids.

Carbonic Acid - Bicarbonate Buffer System

• Components: The equilibrium exists between carbon dioxide (CO2), water (H2O), carbonic acid (H2CO3), bicarbonate (HCO3-), and hydrogen ions (H+), all pivotal in maintaining pH homeostasis in biological systems.

• Functions: This system actively regulates pH balance in blood through reversible reactions based on metabolic activities and respiratory performance.

Buffer Response to Basic Blood

• Reaction to High pH: In response to an elevated pH level, the reaction shifts as follows: CO2 + H2O ↔ H2CO3 → HCO3- + H+, thereby lowering the pH to restore balance.

Buffer Response to Acidic Blood

• Reaction to Low pH: When the blood pH drops, excess hydrogen ions (H+) combine with bicarbonate (HCO3-) to form additional carbonic acid (H2CO3), effectively increasing the pH and re-establishing homeostasis.

Introduction

• Quote referencing an attractive external element: Br Na+

Chemical Fundamentals

• Significance of Molecules: Molecules such as lipids, nucleic acids, proteins, and carbohydrates are fundamental to cellular structures and processes, playing critical roles in metabolism, genetic information storage and transmission, and energy production. Mastery of the chemistry of these molecules is vital for understanding biology and is encapsulated in the field termed Biochemistry.

• Review of Previous Chemistry: A reminder of essential concepts from prior education (10th and 11th grades) that form the foundation for understanding complex biochemical processes.

Isotopes

• Definition: Isotopes are atoms that share the same number of protons and electrons but differ in the number of neutrons. This results in variations in atomic mass and, in some cases, nuclear stability.

• Example: Hydrogen isotopes include Protium (1 proton, 0 neutrons), Deuterium (1 proton, 1 neutron), and Tritium (1 proton, 2 neutrons), each with distinct physical and chemical properties.

Radioisotopes

• Characteristics of Radioisotopes: Radioisotopes possess unstable nuclei that decay spontaneously through a process known as radioactive decay. This decay results in the emission of radiation, which can be harnessed in various applications.

• Half-life: The half-life of a radioisotope is the time required for half of the radioactive nuclei in a sample to decay, which is crucial for dating archaeological finds and medical applications.

• Applications: A prime example is its application in radiometric dating, such as determining the age of the Pygmy Woolly Mammoth, estimated to be 1800 years old through isotope analysis.

Chemical Behavior

• Electron Movement: Electrons move around the nucleus at varying distances, with their distance being intrinsically related to their energy level. As energy increases, electrons can occupy higher energy levels.

• Energy Levels: Electrons exist in distinct energy levels, denoted as n=1, n=2, etc., with each level associated with specific energies and locations.

• Valence Electrons: The outermost electrons, known as valence electrons, are key determiners of an atom's chemical behavior, influencing bonding preferences and reactivity.

Chemical Bonding

• Types of Bonds:

  • Intramolecular Forces (within molecules): These include ionic bonds, which result from the transfer of electrons, and covalent bonds, which involve sharing electrons.

  • Intermolecular Forces (between molecules): These include weaker forces like hydrogen bonds, hydrophobic interactions, and van der Waals forces, which are crucial in determining the properties of substances.

Ionic Bonds

• Definition: Ionic bonds form from the electrostatic attraction between oppositely charged ions (cation and anion).

• Example: In sodium chloride (NaCl), sodium (Na+) donates an electron to chlorine (Cl-) resulting in a strong ionic bond. In aqueous solutions, NaCl dissociates into free ions (Na+ and Cl-), contributing to its conductivity and solubility properties.

Covalent Bonds

• Formation: Covalent bonds occur when two atoms share pairs of valence electrons, achieving greater stability. The sharing can be equal (nonpolar covalent) or unequal (polar covalent).

• Example: A water molecule (H2O) is formed from the covalent bonds between two hydrogen atoms and one oxygen atom, resulting in a bent molecular shape and polar characteristics.

Polarity in Covalent Bonds

• Polar Covalent Bonds: These result from the unequal sharing of electrons, leading to partial positive (δ+) and negative (δ-) charges on atoms, driven by differences in electronegativity. This polarity is crucial for the interactions between different molecules.

Molecular Shapes

• VSEPR Model: The Valence Shell Electron Repulsion model provides a visual representation of molecular shapes based on repulsions between valence electron pairs.

• Examples of Shapes:

  • Methane (CH4): Tetrahedral geometry with a bond angle of approximately 109.5°.

  • Ammonia (NH3): Pyramidal shape with bond angles close to 107°, influenced by one lone pair of electrons.

  • Water (H2O): Angular shape with bond angle around 104°, significant for its polarity.

  • Hydrogen Chloride (HCl): Linear molecule due to the single bond between hydrogen and chlorine.

Molecular Polarity Criteria

• Conditions for Polarity: A molecule must possess polar covalent bonds and have an asymmetrical arrangement of those bonds to be classified as polar, enabling interactions with other polar substances.

Properties of Water

• Water as a Polar Molecule: Water exhibits unique properties due to its polarity, making it an excellent solvent for ionic and other polar compounds. Its polar nature leads to hydrogen bonding effects that significantly influence biochemical processes.

• Hydrogen Bonding: One of the strongest intermolecular forces, hydrogen bonding is critical for the unique physical properties of water, such as high specific heat, surface tension, and cohesion, all necessary for supporting life.

Density of Water

• Ice as an Insulator: Ice's lower density compared to liquid water allows it to float, providing an insulating layer that keeps water below from freezing, thus protecting aquatic life.

• Habitat for Organisms: Bodies of water serve as habitats for a range of organisms, both unicellular and multicellular, playing crucial ecological roles.

Importance of Water Solubility

• Key Role: Water is essential for life; it facilitates biochemical reactions and helps maintain cell structure by operating within watery cellular environments.

• Polar Molecules: Polar compounds (e.g., glucose) tend to be soluble in water, which is vital for cellular metabolism and nutrient transport.

• Non-Polar Molecules: Non-polar substances (e.g., octane) exhibit hydrophobic properties, making them insoluble in water, which has implications for their behavior in biological systems.

Intermolecular Forces - Hydrogen Bonding

• Definition: The hydrogen bonding phenomenon refers to weak attractions between slightly positive hydrogen atoms and slightly negative atoms (such as O, N, or F) in neighboring molecules, which significantly impact the physical properties of substances.

Intermolecular Forces - Hydrophobic Interactions

• Hydrophobic Effect: Non-polar molecules tend to cluster together in an aqueous environment due to their mutual repulsion from water, influencing the structure of biological membranes.

• Plasma Membrane: The arrangement of molecules in the plasma membrane is influenced by hydrophobic interactions, dictating the membrane's permeability and fluidity, which are essential for cellular function.

Biochemical Reactions Overview

• Types of reactions covered:

  • Neutralization reactions (covered in coming slides)

  • Condensation reactions (to be discussed in the next lesson)

  • Hydrolysis reactions (to be discussed in the next lesson)

  • Redox reactions (to be explored in Unit two).

Acids

• Definition: Acids are substances that, when dissolved in water, release hydrogen ions (H+), resulting in the formation of hydronium ions (H3O+).

• Example: When hydrogen chloride (HCl) dissolves in water, it ionizes to produce hydrochloric acid, a strong acid commonly found in gastric juice.

Bases

• Definition: Bases are substances that generate hydroxide ions (OH-) upon dissolution in water.

• Example: Sodium hydroxide (NaOH) dissociates in water to yield Na+ and OH- ions, functioning as a strong base used in various industrial applications.

Neutralization Reaction

• General Reaction: The typical reaction between an acid and a base produces water and a salt, a fundamental aspect of acid-base chemistry.

• Example: The reaction between HCl(aq) and NaOH(aq) follows the equation: HCl + NaOH → H2O + NaCl, demonstrating the formation of common table salt.

Acid-Base Buffers

• Buffers: Buffers are crucial substances that help in moderating fluctuations in pH levels, maintaining homeostasis in biological systems.

• Example: Human blood has a normal pH operating around 7.4 (acceptable range: 7.35 to 7.45), with buffers helping to stabilize this pH against dietary or metabolic fluctuations.

• Natural Buffers: These substances neutralize excess acids or bases; the bicarbonate system is a key buffer mechanism.

• Key Buffer System: The Carbonic Acid - Bicarbonate system plays a significant role in pH regulation in biological fluids.

Carbonic Acid - Bicarbonate Buffer System

• Components: The equilibrium exists between carbon dioxide (CO2), water (H2O), carbonic acid (H2CO3), bicarbonate (HCO3-), and hydrogen ions (H+), all pivotal in maintaining pH homeostasis in biological systems.

• Functions: This system actively regulates pH balance in blood through reversible reactions based on metabolic activities and respiratory performance.

Buffer Response to Basic Blood

• Reaction to High pH: In response to an elevated pH level, the reaction shifts as follows: CO2 + H2O ↔ H2CO3 → HCO3- + H+, thereby lowering the pH to restore balance.

Buffer Response to Acidic Blood

• Reaction to Low pH: When the blood pH drops, excess hydrogen ions (H+) combine with bicarbonate (HCO3-) to form additional carbonic acid (H2CO3), effectively increasing the pH and re-establishing homeostasis.