Bio 100: Introduction to Chemistry, Atoms, and Bonds

Foundations of Matter and Mass

• Definition of Matter:

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

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

• Definition of Mass:

  • Mass is a measure of the amount of material present in an object.

  • Comparative Example: A bowling ball compared to a beach ball. A bowling ball held in a person's hands is significantly heavier than a beach ball of similar volume. Consequently, the bowling ball possesses greater mass than the beach ball because it contains a greater amount of material.

The Periodic Table, Chemical Elements, and Compounds

• Chemical Elements:

  • All matter is composed of chemical elements.

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

  • Example: Carbon ($\text{C}$) is a chemical element and cannot be simplified into other substances via chemical reactions.

• Quantitative Breakdown of Elements:

  • There are $92$ naturally occurring elements.

  • Among these $92$ naturally occurring elements, $25$ are essential to human life.

  • Four of these $25$ essential elements make up $96\%$ of human total body weight.

• The Four Essential Body Elements:

  • Carbon ($\text{C}$)

  • Oxygen ($\text{O}$)

  • Hydrogen ($\text{H}$)

  • Nitrogen ($\text{N}$)

  • These four elements form the foundational building blocks for the four major classes of biological organic molecules:

    • Carbohydrates

    • Proteins

    • Nucleic acids

    • Lipids

• Trace Elements:

  • Trace elements are required by organisms for survival, but only in extremely small quantities.

  • Examples: Copper ($\text{Cu}$), Iron ($\text{Fe}$).

• Molecules and Compounds:

  • Molecule: A group of two or more atoms bonded together.

  • Compound: A molecule specifically composed of two or more different chemical elements.

  • Example: Combining sodium ($\text{Na}$) and chlorine ($\text{Cl}$) produces sodium chloride ($\text{NaCl}$), commonly known as table salt.

Structure of the Atom and Subatomic Particles

• Definition of an Atom:

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

• Subatomic Particles:

  • Atoms are composed of three primary subatomic particles: protons, neutrons, and electrons.

• Subatomic Particle Characteristics:

  • Protons:

    • Electrical Charge: Positive ($+$)

    • Mass: $1$

    • Location: Situated inside the atomic nucleus

  • Neutrons:

    • Electrical Charge: Neutral / No charge ($0$

    • Mass: $1$

    • Location: Situated inside the atomic nucleus

  • Electrons:

    • Electrical Charge: Negative ($-$)

    • Mass: $0$ (negligible mass)

    • Location: Floating/orbiting in energy shells outside the nucleus

    • Biological Significance: Electrons drive the formation of chemical bonds between atoms.

Atomic Mass, Atomic Number, and Isotopes

• Atomic Number:

  • The atomic number is defined as the unique number of protons contained within all atoms of a specific element.

  • Example: Carbon has $6$ protons; therefore, its atomic number is $6$.

• Atomic Mass:

  • Atomic mass is the sum total of the number of protons and neutrons within an atom's nucleus:         $\text{Atomic Mass} = \text{Protons} + \text{Neutrons}$

  • Example: Carbon has a total atomic mass of $12$. Since its proton count is $6$ (contributing a mass of $6$), the remaining mass of $6$ is supplied by $6$ neutrons ($6 + 6 = 12$).

• Isotopes:

  • Isotopes are variant forms of a single chemical element that possess the exact same number of protons but differ in their number of neutrons.

  • The proton count in isotopes remains entirely constant, while the neutron count varies.

  • Isotopes of Carbon:

    • Carbon-12 ($\text{C-}12$): Contains $6$ protons and $6$ neutrons ($\text{Mass} = 12$). This is the standard naturally occurring form.

    • Carbon-13 ($\text{C-}13$): Contains $6$ protons and $7$ neutrons ($\text{Mass} = 13$).

    • Carbon-14 ($\text{C-}14$): Contains $6$ protons and $8$ neutrons ($\text{Mass} = 14$).

Chemical Reactions and Energy Shells

• Chemical Reactions:

  • A chemical reaction is defined as a process in which one or more substances undergo changes resulting in either the sharing or transferring of electrons.

  • Electrons are the key subatomic particles driving chemical reactions.

• Planetary Model of Electron Distribution:

  • An atom can be modeled similarly to a planetary system: the atomic nucleus acts as the central Sun, and electrons act as planets orbiting at fixed energy shell levels (orbitals) situated at varying distances from the center.

  • Elements evaluated under this orbital arrangement include Hydrogen, Carbon, Nitrogen, Oxygen, Neon, and Phosphorus.

  • The energy shell situated furthest from the nucleus (valence shell) is the most critical for governing chemical bonding.

• Energy Shell Capacity Rules:

  • Lowest Energy Shell (Innermost Shell): Contains $1$ orbital and holds a maximum capacity of $2$ electrons.

  • Valence Shell (Outermost Shell): Governs chemical reactivity and bonding based on its number of electron vacancies.

• Electron Pairing Behavior and Stability:

  • Electrons naturally resist existing as single unpaired entities; they must be aligned in pairs to achieve stability.

  • Carbon Stability Profile: Carbon contains $2$ total energy shells. The inner shell holds $2$ electrons. The outer valence shell contains $4$ free electrons and $4$ vacancies. Because of these $4$ empty vacancies, an isolated carbon atom is highly unstable.

  • Atoms achieve stability by forming chemical bonds with other atoms to fill their valence shell vacancies.

• Chemical Synthesis Example — Methane ($\text{CH}_4$):

  • $1$ Carbon atom ($4$ valence vacancies) reacts with $4$ Hydrogen atoms (each containing $1$ valence electron and $1$ vacancy).

  • The single Carbon atom forms bonds with all $4$ separate Hydrogen atoms ($\text{A} + \text{B} \rightarrow \text{AB}$).

  • This combination fills all vacant electron slots, yielding a single molecule of methane gas ($\text{CH}_4$).

Chemical Bonding: Ionic and Covalent Bonds

• The Three Major Types of Chemical Bonds:

  1. Covalent bonds

  2. Ionic bonds

  3. Hydrogen bonds (critically important for water molecules)

• Ions:

  • An ion is defined as an atom or molecule that has become electrically charged as a direct result of either gaining or losing one or more electrons.

  • Anion: An ion formed when an atom gains electrons. Because electrons carry a negative charge, adding electrons produces a net negative charge ($-$).

  • Cation: An ion formed when an atom loses electrons. Because electrons carry a negative charge, losing electrons produces a net positive charge ($+$).

• Ionic Bonds:

  • An ionic bond is formed between two oppositely charged ions through the complete transfer of electrons from one atom to another (one atom completely takes or strips the electron from the other).

  • Sodium Chloride ($\text{NaCl}$) Example: Sodium ($\text{Na}$) possesses $1$ single electron in its outer valence shell, making it unstable. Chlorine ($\text{Cl}$) possesses $7$ valence electrons in its outer shell. Chlorine takes the single valence electron away from Sodium. This electron transfer creates a positively charged sodium cation ($\text{Na}^+$) and a negatively charged chloride anion ($\text{Cl}^-$). The resulting opposite electrical charges attract one another to form an ionic bond, generating sodium chloride ($\text{NaCl}$) table salt.

• Covalent Bonds:

  • A covalent bond is formed when two atoms share one or more pairs of electrons between them.

  • Example: Two Hydrogen atoms, each possessing $1$ electron in their outer shell, share their single electrons with each other to complete their valence shell and form a covalent bond.

• Fundamental Comparison:

  • Ionic Bonding: Involves the complete transfer or taking of electrons.

  • Covalent Bonding: Involves the mutual sharing of electron pairs.

Questions & Discussion

• Question regarding rounding atomic mass values:

  • Query: Should atomic mass values be rounded up to whole numbers?

  • Answer: Rounding is frequently done in formal chemistry courses. However, in introductory biology, the precise rounding protocol is secondary to understanding the core mathematical relationship ($\text{Atomic Mass} = \text{Protons} + \text{Neutrons}$).

• Interactive Problem regarding Nitrogen subatomic structure:

  • Query: The atomic weight of nitrogen is very near $14$, indicating that most nitrogen atoms have a mass number of $14$. How many neutrons does the average nitrogen atom have?

  • Answer: $7$ neutrons.

  • Explanation: Nitrogen has an atomic number of $7$, meaning it possesses $7$ protons. If the total atomic mass number is $14$, subtracting the $7$ protons from the total mass of $14$ yields exactly $7$ neutrons ($14 - 7 = 7$).

• Question regarding maximum energy shell electron capacities:

  • Query: Does electron shell capacity increase with each subsequent shell, going up to $12$ or $16$?

  • Answer: Electron capacity increases in higher energy shells, reaching capacities such as $12$ or up to $36$ electrons.