Basic and Inorganic Chemistry Notes

Matter and Atomic Structure

• Definition of Matter:

  • Matter is defined as anything that occupies space and has mass (weight).

  • It constitutes the physical (living and non-living) "stuff" of the universe.

  • Matter exists in three primary states: solid, liquid, or gas.

• Mass and Weight:

  • Mass is the quantification of the amount of a substance.

  • Under the influence of gravity on the surface of the Earth, mass is equal to the familiar term weight.

• Composition of Matter:

  • Elements:

  • Elements are fundamental units of matter that cannot be broken down into other substances by chemical means.

  • Approximately $96\%$ of living organisms is composed of four primary elements:

    • Hydrogen ($H$)

    • Oxygen ($O$)

    • Nitrogen ($N$)

    • Carbon ($C$)

  • Atoms:

  • Atoms serve as the building blocks of elements.

  • Atomic Structure:

    • Nucleus: Located at the center of the atom, containing:

    • Protons ($p^+$): Positively charged subatomic particles.

    • Neutrons ($n^0$): Uncharged (neutral) subatomic particles.

    • Outside the Nucleus:

    • Electrons ($e^-$): Negatively charged subatomic particles orbiting the nucleus.

Identifying Elements and Isotopes

• Distinguishing Elements:

  • Elements differ from one another based on the specific number of subatomic particles contained within their constituent atoms.

  • Each element is identified by:

  • Chemical Symbol: A standard one- or two-letter abbreviation (e.g., $H$, $O$, $N$, $C$, $Na$, $Cl$).

  • Atomic Number: Equal to the total number of protons ($p^+$) present in the atom's nucleus.

  • Mass Number / Atomic Mass: Equal to the combined sum of protons and neutrons ($p^+ + n^0$) in the nucleus.

• Subatomic Profiles of Key Elements:

  • Carbon ($^{12}C$): $6\,p^+$, $6\,n^0$, $6\,e^-$ (Mass Number = $12$).

  • Nitrogen ($^{14}N$): $7\,p^+$, $7\,n^0$, $7\,e^-$ (Mass Number = $14$).

  • Sodium ($^{23}Na$): $11\,p^+$, $12\,n^0$, $11\,e^-$ (Mass Number = $23$).

  • Chlorine ($^{35}Cl$): $17\,p^+$, $18\,n^0$, $17\,e^-$ (Mass Number = $35$).

• Isotopes:

  • Isotopes are structural variations of an element that contain the same number of protons ($p^+$) and electrons ($e^-$), but vary in the number of neutrons ($n^0$).

  • Examples of Carbon Isotopes:

  • Carbon-12 ($^{12}C$): $6\,p^+$, $6\,n^0$, $6\,e^-$.

  • Carbon-13 ($^{13}C$): $6\,p^+$, $7\,n^0$, $6\,e^-$.

  • Carbon-14 ($^{14}C$): $6\,p^+$, $8\,n^0$, $6\,e^-$.

• Radioisotopes and Radioactivity:

  • Radioisotope: A heavy isotope that is unstable and spontaneously decomposes into a more stable atomic form.

  • Radioactivity: The process of spontaneous atomic decay.

  • As unstable isotopes reorganize into stable forms, they emit measurable energy termed radiation.

  • Medical Diagnostic Applications (Thyroid Scan):

  • Low levels of radiation are safely utilized in diagnostic medicine.

  • Organizational Hierarchy context: Chemical $\rightarrow$ Cell $\rightarrow$ Tissue $\rightarrow$ Organ $\rightarrow$ Organ System.

  • Normal physiological mechanism: Epithelial tissue cells in the thyroid gland (an organ) selectively take up dietary iodine ($I$) to synthesize thyroid hormones.

  • Diagnostic procedure: Radioactive iodine is administered orally to the patient. As thyroid cells absorb the radioactive iodine, emitted radiation is captured by a machine (scanned).

  • The resulting diagnostic image reveals whether the thyroid gland displays cancerous tumors, an underactive state, an overactive state, or normal function.

Electrons and Chemical Bonding

• Electron Shells and Energy Levels:

  • Electrons orbit the nucleus within defined energy levels known as electron shells.

  • Electrons located closest to the nucleus experience the strongest attractive forces.

  • Electron shells fill sequentially from the innermost shell outward.

• Rules of Shell Capacity (Rule of 8s / Octet Rule):

  • Shell 1 (innermost): Holds a maximum of $2\,e^-$.

  • Shell 2: Holds a maximum of $8\,e^-$.

  • Shell 3: Holds a maximum of $8\,e^-$ (for chemical bonding purposes).

• Valence Shell Dynamics:

  • Chemical bonding interactions exclusively involve electrons in the outermost shell, known as the valence shell.

  • Inert Elements:

  • Atoms that possess complete, fully filled valence shells.

  • They are chemically stable and non-reactive.

  • Examples: Helium ($He$), Neon ($Ne$).

  • Reactive Elements:

  • Atoms whose valence shells are incomplete (fewer than $8\,e^-$, or fewer than $2\,e^-$ for Shell 1).

  • They are chemically unstable and actively seek to gain, lose, or share electrons.

  • Electron rearrangement allows atoms to achieve stable valence shells, leading to chemical bond formation.

Types of Chemical Bonds

• Ionic Bonds:

  • Ionic bonds form via attractive electrostatic forces between oppositely charged ions.

  • Ions: Charged particles or elements formed when atoms gain or lose electrons.

  • Cations: Positively charged ions resulting from electron loss (e.g., $H^+$, $K^+$, $Na^+$).

    • Sodium Atom ($11\,p^+$, $11\,e^-$) loses $1\,e^-$ to become a Sodium Cation ($Na^+$) with $10\,e^-$.

  • Anions: Negatively charged ions resulting from electron gain (e.g., $HCO_3^-$, $OH^-$, $Cl^-$).

    • Chlorine Atom ($17\,p^+$, $17\,e^-$) gains $1\,e^-$ to become a Chloride Anion ($Cl^-$) with $18\,e^-$.

  • Salts and Electrolytes:

  • Substances formed via ionic bonds are categorized as salts or electrolytes.

  • In aqueous environments, ionic compounds dissociate into individual positive and negative ions:

$NaCl \rightarrow Na^+ + Cl^-$

• Covalent Bonds:

  • Covalent bonds occur when atoms attain stability by directly sharing electrons.

  • Single Covalent Bond: Sharing of one pair of electrons (1 electron contributed by each atom).

  • Double Covalent Bond: Sharing of two pairs of electrons (2 electrons contributed by each atom).

  • Examples: Methane ($CH_4$), Molecular Oxygen ($O_2$).

  • Molecular Polarity in Covalent Structures:

  • Nonpolar Covalent Molecules:

    • Electrons are shared equally between atoms.

    • The molecule remains electrically neutral overall.

    • Example: Carbon Dioxide ($CO_2$) with structure $O=C=O$.

  • Polar Covalent Molecules:

    • Electrons are shared unequally due to differing electronegativities.

    • Produces partial positive ($\delta^+$) and partial negative ($\delta^-$) poles on the molecule.

    • Example: Water ($H_2O$), where oxygen carries a partial negative charge and hydrogen carries a partial positive charge.

• Hydrogen Bonds:

  • Hydrogen bonds are weak chemical attractions rather than true intramolecular bonds.

  • Formed when a partially positive hydrogen atom ($\delta^+$) in a polar covalent molecule is attracted to the partially negative region ($\delta^-$) of another polar molecule.

  • Provides important intermolecular forces that hold separate molecules together.

Molecules, Compounds, Solutions, and Concentrations

• Molecules vs. Compounds:

  • Molecules: Two or more atoms of the same or different elements joined together by chemical bonds (e.g., $O_2$, $N_2$, $H_2O$).

  • Compounds: Substances composed specifically of two or more different elements combined chemically (e.g., $H_2O$, $CH_4$, $NaCl$, $C_6H_{12}O_6$).

  • Terminology Exception: Because electrolytes like sodium chloride ($NaCl$) are held together by electrostatic ionic attraction rather than covalent sharing, the term "molecule" is technically not applied to $NaCl$

• Solutions:

  • Solutions consist of two or more components physically intermixed without being chemically bound.

  • Physiological Examples: Saline solution ($NaCl$ dissolved in water), blood plasma, interstitial fluid, urine.

  • Solvent:

  • The dissolving medium present in the greatest abundance.

  • Water is the human body's primary universal solvent.

  • Solute:

  • The substance dissolved in the solvent, present in smaller amounts.

  • Examples: $NaCl$, glucose, $O_2$, $CO_2$, $Ca^{2+}$.

• Solute Concentration Units:

  • Percent: Concentration expressed as parts of solute per 100 parts of total solution.

  • Milligrams per Deciliter ($mg\,dL^{-1}$): Mass of solute per deciliter of solution ($1\,dL = 100\,mL$).

  • Molarity ($M$): Expression of moles of solute per liter of total solution.

Chemical Reactions and Metabolism

• Reaction Terminology:

  • Reactants: The starting substances entering into a chemical reaction.

  • Products: The final substances produced by a chemical reaction.

  • Metabolism: The sum total of all chemical reactions occurring within the body.

  • Example Reaction:

$4H + C \rightarrow CH_4$

• Three Major Types of Chemical Reactions:

  • Synthesis Reactions (Anabolism):

  • Atoms, molecules, ions, or simpler compounds combine to form larger, more complex molecules.

  • Always involves new chemical bond formation.

  • General Formula:

$A + B \rightarrow AB$

* Examples: Synthesis of human muscle tissue; joining amino acids to form complex proteins.

• Decomposition Reactions (Catabolism):

  • Larger, complex molecules are broken down into smaller components or individual atoms.

  • Chemical bonds are broken.

  • General Formula:

$AB \rightarrow A + B$

* Examples: Digestive breakdown of ingested animal protein; breakdown of stored glycogen into individual glucose molecules.

• Reversible Reactions:

  • Reactions that proceed in both forward and reverse directions to establish chemical equilibrium.

  • Directionality is denoted by opposing reaction arrows.

  • Carbonic Acid Buffering System (crucial for physiological $CO_2$ balance and pH regulation):

$CO_2 + H_2O \rightleftharpoons H_2CO_3 \rightleftharpoons H^+ + HCO_3^-$

Inorganic Compounds and Biological Significance

• General Characteristics of Inorganic Compounds:

  • Lack carbon atoms (with essential exceptions such as carbon dioxide $CO_2$ and bicarbonate $HCO_3^-$).

  • Tend to be smaller, simpler molecules compared to complex organic molecules (e.g., $H_2O$ and $NaCl$ vs. $C_6H_{12}O_6$).

• Key Inorganic Compounds in Human Physiology:

  • Water ($H_2O$):

  • Most abundant inorganic compound in living organisms, constituting $60\% - 80\%$ of total body weight.

  • Vital Properties of Water:

    • High Heat Capacity: Absorbs and releases large amounts of heat energy before significantly changing temperature, thereby stabilizing body temperature.

    • High Heat of Vaporization: Requires extensive heat energy to transform from liquid to gas, providing an efficient cooling mechanism (sweating).

    • Polar Solvent Properties: Functions as the body's universal solvent. Biological molecules cannot react chemically unless dissolved in solution. Forms hydration layers around charged particles to shield charges and serves as a vital transport medium (e.g., blood plasma).

    • Chemical Reactivity: Serves as an essential reactant in hydrolysis reactions and is removed during dehydration synthesis.

    • Cushioning: Protects delicate body organs (such as brain tissue and fetus) against physical trauma.

    • Macromolecular Stabilization: Maintains structural integrity of macromolecules via hydrophobic interactions.

  • Oxygen ($O_2$):

  • Comprises approximately $20\%$ of atmospheric air.

  • Essential for cellular respiration to extract energy from nutrients.

  • Cellular deprivation of $O_2$ rapidly depletes cellular energy ($ATP$), resulting in cell death.

  • Carbon Dioxide ($CO_2$):

  • Waste product generated when long carbon-chain nutrient molecules are broken down for energy extraction.

  • Must be continuously removed from the body via respiration.

  • Salts:

  • Ionic compounds containing cations other than $H^+$ and anions other than $OH^-$ (e.g., $NaCl$, $CaCO_3$, $KCl$).

  • Readily dissociate into constituent ions in the presence of water, forming electrolytes capable of conducting electrical currents.

  • Electrolyte ions are critical for physiological functions, including nerve impulse transmission and muscle contraction.

  • Disruption of ionic balance (regulated primarily by the kidneys) impairs essential physiological processes, leading to widespread system failure and rapid death.

  • Pharmacological Application: Lithium Chloride ($LiCl$), an antidepressant, dissociates completely in water:

$LiCl \rightarrow Li^+ + Cl^-$

Acids, Bases, and pH Homeostasis

• Electrolyte Behavior of Acids and Bases:

  • Like salts, acids and bases function as electrolytes that ionize and dissociate in water to conduct electrical currents.

• Acids:

  • Defined as proton donors because they release hydrogen ions ($H^+$) in solution.

  • A free hydrogen ion ($H^+$) is functionally an isolated proton.

  • Weak Acids:

  • Incompletely/partially dissociate in water.

  • Example: Carbonic acid ($H_2CO_3$

$H_2CO_3 \rightleftharpoons H^+ + HCO_3^-$

* The body utilizes the decomposition phase of this reversible system to generate $CO_2$ for pulmonary elimination.

• Strong Acids:

  • Completely dissociate in water to release maximum $H^+$ concentrations.

  • Example: Hydrochloric acid ($HCl$

$HCl \rightarrow H^+ + Cl^-$

* The stomach utilizes complete $HCl$ dissociation to create an intensely acidic digestive environment.

• Bases:

  • Defined as proton acceptors because they bind and neutralize free hydrogen ions ($H^+$) in water.

  • Physiological Base Reaction Examples:

$NH_3 + H^+ \rightarrow NH_4^+$

$H^+ + HCO_3^- \rightarrow H_2CO_3$

• The kidneys utilize the synthesis reaction ($H^+ + HCO_3^- \rightarrow H_2CO_3$) to bind and eliminate excess metabolic acid ($H^+$) from the body.

  • The pH Scale and Hydrogen Ion Concentration:

• The pH scale quantifies the concentration of free hydrogen ions ($[H^+]$) in a solution.

• Scale ranges strictly from $0$ to $14$:

  • Neutral Solution ($\text{pH} = 7$): Equal concentrations of $H^+$ and $OH^-$ ions.

  • Acidic Solution ($\text{pH} < 7$): Higher $[H^+]$ concentration (values closer to $0$ represent stronger acidity).

  • Basic / Alkaline Solution ($\text{pH} > 7$): Lower $[H^+]$ concentration (values closer to $14$ represent stronger alkalinity).

• Inverse Relationship: As $[H^+]$ increases, pH decreases (more acidic); as $[H^+]$ decreases, pH increases (more basic).

  • Physiological pH and Homeostatic Buffering:

• Normal physiological pH range for arterial blood in human beings is strictly maintained between $7.35$ and $7.45$

• Acidosis: Occurs when arterial blood pH drops below $7.35$

• Alkalosis: Occurs when arterial blood pH rises above $7.45$

• Buffers:

  • Specialized chemical compounds that stabilize pH by accepting $H^+$ when solutions become too acidic, or releasing $H^+$ when solutions become too basic.

  • Primary homeostatic blood buffer: Carbonic acid-bicarbonate buffer system: