Class 2

Fundamental Principles of Matter and Atomic Models

• Definitions of Matter and Chemical Elements:

  • All organisms and physical entities in the universe are composed of matter.

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

  • Matter consists of chemical elements in pure form as well as in combinations.

  • An element is a substance that cannot be broken down into other substances by chemical reactions.

  • Each element consists of a unique type of atom.

• The Atom:

  • An atom is the smallest unit of matter that retains the chemical properties of an element.

  • Atoms serve as the fundamental building blocks of all matter.

  • An atom is comprised of three primary subatomic particles:

  • Protons: Positively charged subatomic particles ($+$.

  • Neutrons: Electrically neutral subatomic particles (no charge).

  • Electrons: Negatively charged subatomic particles ($-$.

  • Protons and neutrons pack tightly together in the central core to form the atomic nucleus, while electrons move within the space surrounding the nucleus.

• Historical Progression of Atomic Models:

  • 1904 (Plum Pudding Model): Conceived by J.J. Thomson, modeling the atom as a dispersed sphere of positive charge containing embedded electrons.

  • 1911 (Rutherford Model): Introduced by Ernest Rutherford following alpha particle scattering experiments, demonstrating a dense, positively charged nucleus surrounded by orbiting electrons.

  • 1913 (Bohr Model): Proposed by Niels Bohr, describing electrons orbiting the central nucleus in defined concentric circles or electron shells at fixed distances and quantized energy levels. For general biological concepts in this course, the Bohr planetary conceptual understanding is sufficient.

  • 1926 (Quantum Mechanical Model): Developed by Erwin Schrödinger and contemporaries, treating electrons not as particles in rigid orbits, but as wave functions ($\Psi$) representing three-dimensional probability distributions (electron orbitals or electron clouds).

Atomic Identity, Isotopes, and Valence

• Atomic Identity and Atomic Number:

  • The identity of an element is determined exclusively by its number of protons (atomic number).

  • All atoms of a specific element share identical numbers of protons.

  • Example: An oxygen atom possesses an atomic number of $8$, signifying that it invariably contains $8$ protons.

  • In an uncharged, neutral atom, the number of protons equals the number of electrons.

• Isotopes:

  • Isotopes are distinct atomic forms of the same element that have identical numbers of protons but differ in their number of neutrons.

  • While isotopes possess different atomic masses, they exhibit identical chemical behaviors because chemical reactivity is governed by electron configuration.

• Electrons and Quantized Potential Energy:

  • Energy is defined as the capacity to cause change or do work.

  • Potential energy is the energy that matter possesses due to its spatial location or structural arrangement.

  • Matter has an intrinsic tendency to move toward the lowest achievable state of potential energy.

  • The electrons of an atom possess potential energy determined by their distance from the positively charged nucleus; electrostatic attraction between negatively charged electrons and positively charged protons pulls electrons toward the center.

  • Energy levels of electrons are quantized into discrete electron shells:

  • First Shell: Located closest to the nucleus; possesses the lowest potential energy state.

  • Second Shell: Possesses a higher energy level than the first shell.

  • Third Shell: Represents an even higher potential energy level.

  • Staircase Analogy: An electron can exist only at discrete energy levels, analogous to a ball bouncing on a flight of stairs that can rest stably on individual steps, but never suspended between steps.

  • Energy Transitions:

  • Absorbing energy: An electron moves outward to a shell further from the nucleus when it absorbs an amount of energy equal to the difference between the two shells.

  • Losing energy: An electron emits or loses energy (often as electromagnetic radiation or heat) when falling back to a shell closer to the nucleus.

• Valence Electrons and Bonding Capacity:

  • Valence electrons are electrons residing in the outermost electron shell, termed the valence shell.

  • Chemical behavior and reactivity of an atom are predominantly governed by the configuration and quantity of valence electrons.

  • Atoms with completed valence shells are chemically unreactive (inert), whereas atoms with incomplete valence shells interact with other atoms to complete their shells.

  • Valence (Bonding Capacity): Typically equals the number of unpaired electrons needed to fill the outermost valence shell:

  • Hydrogen ($H$): Valence = $1$ (has $1$ valence electron, requires $1$ to complete its single shell capacity of $2$).

  • Oxygen ($O$): Valence = $2$ (has $6$ valence electrons, requires $2$ to complete its shell capacity of $8$).

  • Nitrogen ($N$): Valence = $3$ (has $5$ valence electrons, requires $3$ to complete its shell capacity of $8$).

  • Carbon ($C$): Valence = $4$ (has $4$ valence electrons, requires $4$ to complete its shell capacity of $8$).

Chemical Bonding and Electronegativity

• Mechanisms of Chemical Interactions:

  • Biological systems function through chemical interactions between atoms.

  • Atoms with incomplete valence shells achieve stability by sharing, donating, or accepting valence electrons, holding atoms together via chemical bonds.

  • Electrostatic rules dictate interactions: opposite charges attract ($+$ and $-$), while like charges repel ($+$ and $+$, or $-$ and $-$).

• Electronegativity:

  • Electronegativity is defined as the capability of an atom to attract electrons toward itself within a chemical bond.

  • An atom with high electronegativity pulls electrons closer than an atom with low electronegativity.

  • Periodic Trends:

  • Moving across a period (from left to right): The nuclear charge increases (addition of protons) while electrons enter the same general energy level; this draws the outer electrons closer, reducing atomic radius ($r$).

  • Smaller atomic radius combined with increased nuclear charge creates a much stronger electrostatic attraction to bonding electrons; therefore, electronegativity increases across a period (e.g., from $Li$ to $Be$ to $F$).

  • Moving down a group: Additional electron shells increase atomic radius and electron shielding, generally decreasing electronegativity.

  • Measured Electronegativity Values (Pauling Scale):

  • Potassium ($K$): $0.82$

  • Sodium ($Na$): $0.93$

  • Lithium ($Li$): $0.98$

  • Calcium ($Ca$): $1.00$

  • Magnesium ($Mg$): $1.31$

  • Beryllium ($Be$): $1.57$

  • Aluminum ($Al$): $1.61$

  • Silicon ($Si$): $1.90$

  • Boron ($B$): $2.04$

  • Phosphorus ($P$): $2.19$

  • Hydrogen ($H$): $2.20$

  • Carbon ($C$): $2.55$

  • Sulfur ($S$): $2.58$

  • Nitrogen ($N$): $3.04$

  • Chlorine ($Cl$): $3.16$

  • Oxygen ($O$): $3.44$

  • Fluorine ($F$): $3.98$

  • Noble gases (Helium, Neon, Argon): Exhibit complete valence shells, therefore displaying no electronegativity data under standard parameters.

• Classification of Chemical Bonds Based on Electronegativity Difference ($\Delta\text{EN}$):

  • Nonpolar Covalent Bonds ($\Delta\text{EN} \le 0.4$ to $0.5$):

  • Electrons are shared equally between the two bonded atoms.

  • Occurs when bonding atoms possess identical or nearly identical electronegativities.

  • Examples: Molecular hydrogen ($H_2$, where two hydrogen atoms share their single electrons equally like an evenly matched tug-of-war) and diatomic chlorine ($Cl_2$).

  • Polar Covalent Bonds ($0.5 < \Delta\text{EN} \le 1.7$):

  • Electrons are shared unequally between atoms.

  • The shared electrons spend significantly more time orbiting the atom with higher electronegativity, generating an uneven distribution of charge.

  • Results in partial electrical charges designated by Greek delta symbols: a partial negative charge ($\delta-$) on the more electronegative atom, and a partial positive charge ($\delta+$) on the less electronegative atom.

  • Examples: Hydrogen chloride ($HCl$), where Chlorine ($3.16$) attracts electrons away from Hydrogen ($2.20$), and water ($H_2O$), where Oxygen ($3.44$) pulls electrons away from Hydrogen ($2.20$).

  • Ionic Bonds ($\Delta\text{EN} > 1.7$, up to $3.3$):

  • The difference in electronegativity is so large that the more electronegative atom completely strips (steals) valence electrons away from its bonding partner rather than sharing them.

  • Results in the formation of charged ions: a positively charged cation ($Na^+$) and a negatively charged anion ($Cl^-$).

  • The ionic bond is the electrostatic attraction between oppositely charged ions.

  • Example: Sodium chloride ($NaCl$), where Chlorine ($3.16$) strips an electron from Sodium ($0.93$).

Molecules, Compounds, and Emergent Properties

• Distinction Between Molecules and Compounds:

  • Molecule: Two or more atoms held together specifically by covalent bonds. Atoms may be of the identical element or different elements (e.g., $O_2$, $N_2$, $H_2$, $CO_2$, $H_2O$).

  • Compound: A substance composed of two or more different chemical elements combined in a fixed ratio (e.g., $NaCl$, $CO_2$, $H_2O$, formic acid $CH_2O_2$).

  • All compounds formed by covalent bonds are molecules, but molecules consisting of identical elements (such as $O_2$ or $H_2$) are not compounds.

  • In cellular systems, the specific three-dimensional spatial interactions among molecules govern cellular function; aberrant molecular interactions frequently induce human pathologies and disease states.

• Emergent Properties in Chemical Compounds:

  • A compound possesses emergent properties—chemical and physical characteristics distinct from those of its constituent elemental components.

  • Sodium Chloride ($NaCl$) Case Study:

  • Sodium ($Na$): A soft, silvery, metallic element that reacts violently and combusts vigorously upon contact with liquid water.

  • Chlorine ($Cl$): A pungent, poisonous, toxic pale-yellow gas.

  • Sodium Chloride ($NaCl$): When chemically joined in a $1:1$ ionic ratio, these toxic/reactive elements synthesize table salt, an edible, non-toxic, crystalline ionic compound necessary for mammalian life.

• Biological Application: Formic Acid in Wood Ants:

  • Wood ants (Formica rufa) utilize defensive chemistry by spraying formic acid ($CH_2O_2$) to deter and incapacitate predators.

  • Chemical constitution of formic acid:

  • Formic acid is composed of carbon ($C$), hydrogen ($H$), and oxygen ($O$).

  • The proton count defines each atom's elemental identity (e.g., oxygen contains $8$ protons).

  • The valence electron distribution dictates bonding: oxygen requires $2$ additional electrons to complete its valence octet, thus forming $2$ covalent bonds.

  • The physical and biological properties of the compound depend strictly on the identities and bond configurations of its constituent atoms: in formic acid, an electronegative oxygen atom draws electron density away from a bonded hydrogen atom, allowing the hydrogen to dissociate readily as a free hydronium/hydrogen ion ($H^+$). This proton release confers acidic properties that cause stinging pain.

Elemental Abundance in Biology and the Primacy of Oxygen

• Elemental Composition of the Human Body:

  • Major Four Elements ($96.3\%$ of Total Body Mass):

  • Oxygen ($O$): $65.0\%$

  • Carbon ($C$): $18.5\%$

  • Hydrogen ($H$): $9.5\%$

  • Nitrogen ($N$): $3.3\%$

  • Minor Elements ($3.7\%$ of Total Body Mass):

  • Calcium ($Ca$): $1.5\%$

  • Phosphorus ($P$): $1.0\%$

  • Potassium ($K$): $0.4\%$

  • Sulfur ($S$): $0.3\%$

  • Sodium ($Na$): $0.2\%$

  • Chlorine ($Cl$): $0.2\%$

  • Magnesium ($Mg$): $0.1\%$

  • Trace Elements (Less than $0.01\%$ of Total Body Mass each):

  • Boron ($B$), chromium ($Cr$), cobalt ($Co$), copper ($Cu$), fluorine ($F$), iodine ($I$), iron ($Fe$), manganese ($Mn$), molybdenum ($Mo$), selenium ($Se$), silicon ($Si$), tin ($Sn$), vanadium ($V$), and zinc ($Zn$).

• Comparison of Mass Abundance versus Atomic Abundance:

  • The human body consists of approximately $50\%$ to $60\%$ water ($H_2O$) by mass.

  • Oxygen accounts for the highest proportion of human body mass ($65.0\%$) because an oxygen atom has an atomic mass of approximately $16\text{ amu}$, which is roughly $16$ times heavier than a hydrogen atom (atomic mass approximately $1\text{ amu}$).

  • In terms of total absolute atom count, Hydrogen ($H$) is the most abundant element in the human organism because each water molecule contains two hydrogen atoms for every single oxygen atom ($2:1$ ratio).

• Biological Importance of Oxygen:

  • High electronegativity ($3.44$) allows oxygen to strongly attract shared electrons, establishing molecular polarity.

  • Polarity induced by oxygen is required for:

  • Water Structure: Enables extensive intermolecular hydrogen bonding, conferring cohesion, high specific heat, and solvent capability.

  • Biochemical Reactions: Governs hydrolysis and condensation/dehydration reactions required for assembling and catabolizing macromolecules.

  • Cellular Bioenergetics: Functions as the terminal electron acceptor in cellular respiration, facilitating ATP generation.

  • In the absence of oxygen's high electronegativity, water would lack polarity, biochemical reaction pathways could not operate, and life as configured could not persist.

Questions & Discussion

• Evaluation of Bond Polarity in Hydroxyl Linkages:

  • Prompt: When an oxygen atom is bonded to a hydrogen atom ($O-H$), what will the electric charge be on the oxygen atom?

  • A. No charge

  • B. Partial negative charge

  • C. Partial positive charge

  • Correct Answer: B. Partial negative charge ($\delta-$).

  • Underlying Rationale: Oxygen possesses a high electronegativity of $3.44$, whereas hydrogen possesses an electronegativity of $2.20$. The substantial difference ($\Delta\text{EN} = 1.24$) results in a polar covalent bond where bonding electrons are drawn substantially toward the oxygen nucleus, conferring a partial negative charge ($\delta-$) on oxygen and leaving the hydrogen atom with a partial positive charge ($\delta+$).