chapter 2a

Alice Ball: Diversity Spotlight

Alice Ball earned degrees in chemistry, pharmacy, and pharmaceutical chemistry from the University of Washington (her alma mater).
She was the first African American to graduate with a Master of Science in chemistry, a remarkable achievement given the systemic barriers of her era.
She was the first African American and the first woman to become a chemistry professor at the University of Hawaii, breaking significant racial and gender barriers in academia.
Major medical contribution: developed the first successful injectable treatment for leprosy using the oil of the kalawao tree native to Hawaii. This breakthrough was pivotal because it allowed for systemic treatment, overcoming the limitations of previous oral or topical applications that were often ineffective or poorly tolerated due to gastrointestinal irritation. Her method involved isolating the active esters from the crude chaulmoogra oil.
Historical context: there was a severe leprosy crisis on the island of Molokai during the late 1800s and early 1900s where people with leprosy were forcibly sent to Kalaupapa colony; this underscores the profound real-world, humanitarian impact and urgent need for her work to alleviate suffering and improve lives.
Her injectable treatment was groundbreaking because it was a stable, injectable therapy (specifically, ethyl esters of chaulmoogra fatty acids) that could be absorbed by the body, rather than previous ineffective topical treatments or highly irritating oral doses.
Ball tragically died at age 24 in 1916, likely from chlorine gas poisoning during a laboratory accident, before she could publish her groundbreaking findings and receive due recognition.
After her death, Arthur L. Dean, a chemist and future president of the University of Hawaii, continued her research. He initially published the findings under his name without crediting Ball. Later, due to advocacy, particularly from Dr. Harry T. Hollmann (a physician who worked with Ball and continued her clinical trials), Dean retracted his claim and publicly credited Ball for her pivotal work in 1922. The method based on her work is today known as the Ball method, restoring her rightful place in scientific history after considerable effort.
Ethical and cultural implications: Ball’s story is a clear and poignant example of a brilliant underrepresented person in STEM whose crucial contributions were initially forgotten or stolen. It vividly illustrates broader patterns of underrepresentation, systemic bias, and inequitable credit attribution in scientific discovery, particularly for women and scientists of color.
Reflection: Her story prompts vital discussion about the importance of recognizing and crediting the contributions of marginalized scientists throughout history and actively working to integrate diverse histories into STEM education and curricula to inspire future generations.

Matter and the Building Blocks of Chemistry

Matter- Matter is anything that takes up space and has mass; it makes up all substances on Earth and in the universe.
- It exists in different states: solid (definite shape and volume), liquid (definite volume, indefinite shape), gas (indefinite shape and volume), and plasma (ionized gas, common in stars).
- Living matter vs nonliving matter; matter encompasses everything around us, from macroscopic objects like tables and stars to microscopic entities like organisms and atoms.
Elements and compounds- Elements: pure substances that cannot be broken down into other simpler substances by ordinary chemical reactions. Each element is defined by its unique number of protons.
- Compounds: substances made from two or more different elements chemically bonded together in fixed ratios; compounds have distinct chemical and physical properties that are often very different from their constituent elements (e.g., $ext{H}2$ and $ext{O}2$ are gases, but $ext{H}_2 ext{O}$ is a liquid).
- Essential elements: elements that are vital to an organism’s growth, development, and function; approximately $20 ext{\%}$ to $25 ext{\%}$ of the naturally occurring elements are considered essential for life. Their absence can lead to deficiency diseases.
- Living matter composition: approximately $96 ext{\%}$ of living matter is composed of four major elements: hydrogen (H), carbon (C), nitrogen (N), and oxygen (O). The remaining $4 ext{\%}$ includes essential elements like calcium (Ca) for bones/signaling, phosphorus (P) for DNA/ATP, potassium (K) for nerve function, sulfur (S) for proteins, sodium (Na) for nerve impulses, chlorine (Cl) for osmotic balance, and magnesium (Mg) for enzyme cofactors.
- Trace elements: required by organisms in very minute amounts ( < 0.01 ext{\%} of body mass), but are still critical for proper physiological function (e.g., iron (Fe) for hemoglobin, iodine (I) for thyroid hormones, copper (Cu) for enzyme activity, zinc (Zn) for immune function).
Atoms and subatomic particles- Atom: the fundamental unit of matter; it is the smallest unit of an element that retains its chemical properties. While indivisible by chemical means, atoms are composed of subatomic particles.
- Subatomic particles:
  - Protons ( $p^+$ ): positively charged ( $+1$ ) particles located in the nucleus, defining the atomic number.
  - Neutrons ( $n^0$ ): neutrally charged (0) particles also located in the nucleus, contributing to atomic mass.
  - Electrons ( $e^-$ ): negatively charged ( $-1$ ) particles that orbit the nucleus in specific energy levels or shells. Electrons have a negligible mass (about $1/1836$ th that of a proton) but are crucial for chemical bonding.
- Nucleus: the dense, positively charged center of an atom, containing protons and neutrons. It accounts for almost all of the atom's mass.
- Electron cloud: the region around the nucleus where electrons are found in specific probability distributions called orbitals. These orbitals represent the most likely locations for electrons and define the atom's volume.
- Mass unit: atomic mass is typically measured in daltons (Da) or atomic mass units (amu), where $1 ext{ Da} hickapprox 1.66 imes 10^{-24} ext{ g}$ . One dalton is approximately the mass of a single proton or neutron.
Atomic structure terminology- Atom’s nucleus contains protons and neutrons; electrons move in specific energy levels, or electron shells, around the nucleus.
- Atomic number ( $Z$ ): explicitly defined as the number of protons in the nucleus. This number uniquely identifies an element and determines its position on the periodic table.
- Mass number ( $A$ ): the sum of protons and neutrons in the nucleus ( $A = Z + N$ where $N$ is the number of neutrons). It indicates the approximate total mass of the atom's nucleus.
- Atomic mass ≈ $A ext{ Da}$ (a rough approximation for a specific isotope). The atomic mass displayed on the periodic table is a weighted average of the naturally occurring isotopes of an element.
- Isotopes: atoms of the same element (same $Z$ ) that have different numbers of neutrons ( $N$ ), and thus different mass numbers ( $A$ ). Some isotopes are radioactive and useful in dating or medical imaging.
- Ions: atoms or molecules that have a net electrical charge due to the gain or loss of one or more electrons. Cations are positively charged (lost electrons), and anions are negatively charged (gained electrons).
- Electron shells and energy states: electrons exist in discrete energy levels designated as shells. Electrons can transition between these levels; energy is absorbed when an electron moves to a higher energy shell (excited state) and released (often as light) when it falls back to a lower energy shell (ground state).
- Valence shell: the outermost electron shell of an atom. The electrons in this shell, called valence electrons, are primarily responsible for an atom's chemical behavior and bonding interactions. The number of valence electrons determines an atom's reactivity.
- If a valence shell is full (e.g., 2 electrons for the first shell, 8 for subsequent shells, following the octet rule), the atom is chemically inert or unreactive (e.g., noble gases like Neon or Argon have full valence shells, making them very stable and reluctant to form bonds).
Chemical bonds and molecular structure- Chemical bonds are forces that hold atoms together to form molecules and compounds. Their properties dictate the structure and function of all matter.
- Bonding types:
 - Ionic bonds: formed by the complete transfer of one or more valence electrons from one atom to another, resulting in the formation of oppositely charged ions (cations and anions). These ions are then held together by strong electrostatic attraction. Ionic compounds typically form crystalline solids (salts) and often dissociate in water.
 - Covalent bonds: formed by the sharing of one or more pairs of valence electrons between two atoms. These are generally very strong bonds and result in the formation of molecules.
 - Can be single (one shared pair), double (two shared pairs), or triple (three shared pairs) bonds.
 - The sharing allows each participating atom to achieve a stable electron configuration, typically a full outer valence shell.
 - Hydrogen bonds: weaker, intermolecular forces (or sometimes intramolecular) arising from the attraction between a partially positive hydrogen atom (which is covalently bonded to a highly electronegative atom like O, N, or F) and a partially negative atom (often a lone pair on another electronegative atom like O or N) in a different molecule or part of the same molecule. Crucial for biological structures and water's unique properties, providing stability to large biomolecules like DNA and proteins.
- Electronegativity:
 - Electronegativity is an intrinsic property of an atom that describes its tendency to attract electrons towards itself when it is part of a chemical bond. Represented by scales like the Pauling scale.
 - In modern chemistry, electronegativity generally increases as you move from left to right across a period (due to increasing nuclear charge) and decreases as you move down a group on the periodic table (due to increased atomic radius and electron shielding). Fluorine is indeed the most electronegative element.
 - The difference in electronegativity between two bonding atoms determines the type of covalent bond (polar or nonpolar) or if an ionic bond will form.
- Polar vs nonpolar covalent bonds:
 - Polar covalent bonds: occur when electrons are not shared equally between two atoms because one atom is significantly more electronegative than the other. This unequal sharing creates partial positive ( $\delta^+$ ) and partial negative ( $\delta^-$ ) charges, resulting in a dipole moment across the bond (e.g., O-H bonds in water).
 - Nonpolar covalent bonds: occur when electrons are shared relatively equally between two atoms, usually because they have similar electronegativity values (e.g., C-H bonds, or bonds between identical atoms like $ext{O}2$ or $ext{H}2$ ). There is no significant dipole moment.
- Ionic vs covalent vs hydrogen bonding implications:
 - Ionic bonds form inorganic salts (e.g., NaCl) that can readily dissolve and dissociate in water into free ions (e.g., $ext{Na}^+$ and $ext{Cl}^-$ ), enabling electrical conductivity and participating in biological processes.
 - Covalent bonds form stable molecules (e.g., $ext{CH}4$ , $ext{CO}2$ , $ext{H}_2 ext{O}$ ) by achieving stable valence electron configurations through electron sharing, which is the basis of organic chemistry and all life.
 - Hydrogen bonds, while individually weaker than covalent or ionic bonds, are incredibly significant in biology. Their collective strength and directionality stabilize the double helix structure of DNA (linking complementary base pairs), fold proteins into their specific three-dimensional functional shapes, and profoundly influence the properties of water vital for life.
Water and the properties it confers to life- Water ( $ext{H}_2 ext{O}$ ) is the most abundant molecule on Earth's surface, a universal solvent, and absolutely essential for all known forms of life.
- Water as a molecule: it consists of two hydrogen atoms covalently bonded to one oxygen atom. Oxygen is significantly more electronegative than hydrogen, leading to unequal sharing of electrons and creating a highly polar molecule.
- Polarity and dipole moment:
  - The unequal sharing of electrons in the O–H bonds means the electron density is pulled more towards the oxygen atom. This creates a partial negative charge ( $\delta^-$ ) on the oxygen end and partial positive charges ( $\delta^+$ ) on the hydrogen ends.
  - This inherent polarity, combined with its bent molecular geometry, results in a significant overall dipole moment for the water molecule.
  - This polarity enables water molecules to form extensive hydrogen bonds with each other and to interact with (dissolve) other polar molecules and ions (hydrophilic substances).
- Hydrogen bonding in water:
  - Each water molecule can form up to four hydrogen bonds with neighboring water molecules. These bonds are relatively weak individually ( $hickapprox 5-10 ext{ kcal/mol}$ ) but are numerous and collectively strong, giving water many of its anomalous and life-sustaining properties.
  - Hydrogen bonds underlie water’s high cohesion and surface tension, its unusually high specific heat capacity, its high heat of vaporization, and its unique property of expanding upon freezing.
- Cohesion and surface tension:
  - Cohesion: the strong attraction between water molecules themselves, primarily via hydrogen bonds. This allows water molecules to stick to each other, enabling the formation of water droplets, transport in plants, and creating a continuous liquid medium.
  - Surface tension: a measure of the force required to break the surface of a liquid. The strong cohesive forces (hydrogen bonds) at the air-water interface create a tight, film-like surface layer, allowing lighter objects (like water striders) to rest on water without sinking.
- Adhesion:
  - Adhesion: water’s attraction to different substances or surfaces, typically those that are polar or charged. This property, combined with cohesion, enables phenomena like capillary action (e.g., water climbing up a narrow tube or the vascular tissue of a tree, defying gravity, where adhesion to xylem walls and cohesion among water molecules work together).
- Temperature regulation and heat properties:
  - Water can effectively moderate temperature due to its exceptionally high specific heat capacity: $c ext{ (water)} hickapprox 1 rac{ ext{cal}}{ ext{g} ext{ } {}^ extrm{o} ext{C}}$ ( $4.184 rac{ ext{J}}{ ext{g} ext{ } {}^ extrm{o} ext{C}}$ ). This means it takes a relatively large amount of thermal energy to raise the temperature of water by $1 ext{ } {}^ extrm{o} ext{C}$ per gram.
  - This high specific heat is directly attributable to the extensive hydrogen bonding network. When heat is absorbed, a significant portion of the energy is first used to break hydrogen bonds before the kinetic energy (and thus temperature) of the individual water molecules increases. Conversely, when water cools, hydrogen bonds form, releasing a substantial amount of energy before the temperature drops significantly. This molecular mechanism buffers temperature changes in both aquatic environments and living organisms.
  - Evaporation (vaporization) cools surfaces: water also has a high heat of vaporization. As water absorbs even more heat, molecules at the surface gain enough kinetic energy to break all their hydrogen bonds and escape into the gas phase. This process effectively removes a large amount of heat from the remaining liquid and the surrounding surface, explaining why sweating or panting are effective biological cooling mechanisms.
- Water’s phase changes and environmental relevance:
  - Water exists as solid (ice), liquid (water), or gas (water vapor); phase changes involve the breaking and forming of hydrogen bonds. Unlike most substances, water solidifies at $0 ext{ } {}^ extrm{o} ext{C}$ and boils at $100 ext{ } {}^ extrm{o} ext{C}$ at standard pressure.
  - Water expands upon freezing: As liquid water cools to $0 ext{ } {}^ extrm{o} ext{C}$ , the hydrogen bonds become more rigid and ordered, forcing molecules into a more open, crystalline lattice structure where molecules are spaced further apart than in the liquid state. This makes ice about $9 ext{\%}$ less dense than liquid water, which is why ice floats. This phenomenon is critical for aquatic life, as floating ice insulates the water below, preventing entire bodies of water from freezing solid and allowing organisms to survive winters.
  - Evaporation contributes significantly to biological cooling (sweating in mammals, transpiration in plants) and drives global weather patterns through the water cycle.
- Role in biology and chemistry demonstrations:
  - Fireworks and flame tests illustrate the principles of electron energy transitions between different energy states; specific elements emit light (photons) at characteristic wavelengths when their excited electrons return to lower energy states, producing distinct colors.
  - In biology, hydrogen bonding is fundamental for stabilizing the crucial double helix structure of DNA (through specific base-pairing: A-T forming two H-bonds, G-C forming three H-bonds) and for the correct three-dimensional folding of proteins (critical for secondary structures like alpha-helices and beta-sheets, and tertiary/quaternary stability, which dictates their biological function).
  - A cloud chamber can visualize the paths of energetic subatomic particles (e.g., alpha or beta particles), illustrating atomic-scale behavior and fundamental physics concepts discussed in atomic structure lectures.
Chemical reactions and the role of bonds- Chemical reactions are processes that involve the rearrangement of atoms and molecules due to the making or breaking of chemical bonds to convert reactants into new products.
- During a reaction, chemical energy stored in bonds is either released (exothermic reactions) or absorbed (endothermic reactions).
- Example: The formation of water from hydrogen and oxygen gas ( $2 ext{H}2 ext{(g)} + ext{O}2 ext{(g)} ightarrow 2 ext{H}_2 ext{O (l)}$ ) is a highly exothermic reaction, releasing a substantial amount of energy.
- Many reactions require an initial input of energy, known as activation energy, to overcome barriers and initiate bond breaking/forming processes.
- Catalysts (like enzymes in biological systems) increase the rate of chemical reactions by lowering the activation energy without being consumed in the reaction.
- Salts (ionic compounds) readily dissociate in water into their constituent ions (e.g., $ext{NaCl (s)} ightarrow ext{Na}^+ ext{(aq)} + ext{Cl}^- ext{(aq)}$ in solution). This dissociation is vital for electrolyte balance and signal transmission in biological systems.
- Denaturation of DNA: High heat (e.g., around $95 ext{ } {}^ extrm{o} ext{C}$ ) can provide enough energy to break the relatively weak hydrogen bonds between complementary base pairs (A-T and G-C) in a DNA double helix, causing the two strands to separate (melt). This process is reversible and is crucial for DNA replication and transcription.
- Many chemical reactions proceed to a state of chemical equilibrium, where the rates of the forward and reverse reactions are equal, and the net concentrations of reactants and products remain constant.
Quick notes on common numbers and units mentioned in the lecture- Natural elements: $92$ elements exist naturally on Earth; approximately $20 ext{\%}$ to $25 ext{\%}$ are considered essential for life.
- Major elements in living matter: $ext{C}, ext{H}, ext{O}, ext{N}$ account for roughly $96 ext{\%}$ of the total mass of living organisms.
- Trace elements and other essentials (examples): Ca, P, K, S, Na, Cl, Mg, Fe, I, Zn, Cu.
- Atom mass unit: Da (dalton) or amu (atomic mass unit), where $1 ext{ Da} hickapprox 1 ext{ mass of proton/neutron}$ .
- Electron shell concepts: The valence shell determines an atom's chemical reactivity; full valence shells imply chemical inertness (maximal stability).
- Water’s specific heat: $c hickapprox 1 rac{ ext{cal}}{ ext{g} ext{ } {}^ extrm{o} ext{C}}$ (or $4.184 rac{ ext{J}}{ ext{g} ext{ } {}^ extrm{o} ext{C}}$ ).
- Water formation ratio: $2:1$ ratio for $ext{H}2$ to $ext{O}2$ molecules when forming $ext{H}2 ext{O}$ ( $2 ext{H}2 + ext{O}2 ightarrow 2 ext{H}2 ext{O}$ ).
- Atomic number and mass number relations: $Z = ext{number of protons}, ext{ } A = Z + N ext{ (where } N = ext{number of neutrons)}, ext{ } ext{Atomic mass} hickapprox A ext{ Da (for a specific isotope)}$ .
- Melting/Boiling point of water: $0 ext{ } {}^ extrm{o} ext{C} / 100 ext{ } {}^ extrm{o} ext{C}$ at standard pressure.
Summary of key ideas and connections- Atomic theory provides the foundation for understanding matter: elements are defined by their unique atoms, and their interactions dictate all chemical and biological processes.
- The fundamental bonding types (ionic, covalent, hydrogen) explain the immense diversity of substances, their physical properties, and the intricate behavior of biological systems, from simple molecules to complex proteins and nucleic acids.
- Water’s truly unique and anomalous properties—derived directly from its molecular polarity and extensive hydrogen bonding—are absolutely essential for life. These properties underpin temperature regulation, its role as a universal solvent, and its participation in countless biological processes critical for Earth's climate and organismal survival.
- The story of Alice Ball not only highlights a significant scientific breakthrough in medical treatment but also serves as a critical case study illustrating the vital importance of crediting diverse scientists and actively recognizing the pervasive historical underrepresentation and biases in STEM, encouraging a more inclusive and accurate understanding of scientific progress.

Key Concepts and Formulas (Reference Sheet)

Atomic notation and relations-
- $Z = \text{number of protons (defines the element)}$
- $A = Z + N$ where $N = \text{number of neutrons (defines isotope)}$
- $\text{Atomic mass} \approx A \text{ Da (for a specific isotope)}$
- $\text{Atomic mass on periodic table} = \text{weighted average of isotopes}$
Bond types-
- Covalent bond: sharing of valence electrons; can be single, double, or triple. Characterized by electronegativity difference:
 - Nonpolar covalent: small/zero difference (e.g., < 0.5)
 - Polar covalent: moderate difference (e.g., $0.5 - 1.7$ )
- Ionic bond: complete transfer of electrons, creating oppositely charged ions that attract (large electronegativity difference, e.g., > 1.7). Forms salts.
- Hydrogen bond: intermolecular attraction between $\delta^+ \text{H}$ (bonded to O, N, F) and a lone pair on a $\delta^- \text{atom}$ (O, N, F). Weaker than covalent/ionic, but crucial for biological structure and water properties.
Water-specific properties-
- Polar covalent bonds in $ext{H}_2 ext{O}$ create a net molecular dipole moment: partial negative on O, partial positive on H.
- Extensive hydrogen bonding leads to:
  - High cohesion and surface tension
  - High specific heat capacity ( $\approx 1 \frac{\text{cal}}{\text{g} \cdot {}^{\circ}\mathrm{C}}$ ) enabling temperature moderation
  - High heat of vaporization, facilitating evaporative cooling
  - Expansion upon freezing, making ice less dense and allowing it to float
Biological and chemical relevance-
- Salts dissociate in water into ions, critical for biological electrolyte balance.
- DNA denaturation involves breaking hydrogen bonds with heat (approx. $95 \text{ } {}^{\circ}\mathrm{C}$ ) to separate strands, key for replication and transcription.
- Energy transitions of electrons (e.g., excited state to ground state) produce visible light (fireworks), underlying spectroscopy and fluorescence techniques used to study molecules.
- Enzymes (biological catalysts) lower activation energy to speed up biochemical reactions, which are fundamental to metabolism.
- Chemical equilibrium ensures dynamic balance in many biological and chemical systems.