Introduction to Atomic Structure and Chemical Bonds

Course Overview and Logistics

• Birthday Cake Examples:

  • A double-layer cake designed for a $13$th birthday, celebrated as a diamond birthday because the child was born in the year $2013$. The bottom layer featured a silver and diamond theme, while the top layer featured a mermaid theme.

  • A cake designed for a child turning $7$ years old (moving from age $6$ to $7$), featuring a custom "6, 7" design drawn by the child and executed in frosting.

• Laboratory Protocols and Preparation:

  • Laboratory sessions begin next week.

  • Working with bacteria requires adherence to specific, enhanced safety protocols.

  • To prepare for laboratory sessions, read the introduction section of each lab manual assignment for critical background information and context.

  • Content in the lab introduction may overlap with lecture material, but it often explores topics from a different perspective or in greater depth.

• Quizzes and Assessment Schedule:

  • The first quiz takes place on Wednesday.

  • A sample quiz will be posted on the online course portal (Katie) on Monday to familiarize students with the question format and writing style.

  • The practice quiz is provided specifically for the first quiz to establish expectations; sample quizzes will not be provided for every subsequent quiz.

• Course Roadmap and Chemistry Prerequisites:

  • There is no chemistry prerequisite for this course.

  • Understanding basic chemical principles—specifically how atoms and molecules behave—is mandatory before studying cellular structure and function.

  • Topic Guide 1 covers fundamental chemical principles (Learning Objectives $1a$ and $1b$), followed by Learning Objective 2.

  • The study of pH is postponed to the end of the chemical principles section, as understanding molecular structure is required before analyzing pH.

• Student Code Name System:

  • Student code names are used to anonymously report grades for laboratory worksheets and study plans.

  • Rules for selecting a code name:

  • It cannot be your actual real name.

  • It should not be an existing nickname that others call you.

  • It should be unique, memorable, and personally meaningful or inspiring.

Atomic Structure and Subatomic Particles

• Fundamental Definition of an Atom:

  • An atom is the basic building block and smallest functional unit of an element.

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

• Protons:

  • Charge: Positively charged ($+$).

  • Location: Centered inside the atom within the nucleus.

  • Element Identity: The exact number of protons in the nucleus uniquely defines the identity of an element (its atomic number).

  • An atom containing exactly $1$ proton is always Hydrogen ($H$).

  • An atom containing exactly $6$ protons is always Carbon ($C$).

  • Altering or adding/removing a proton changes the element into a completely different element.

• Neutrons:

  • Charge: Neutral (no electrical charge, $0$).

  • Location: Centered inside the atom within the nucleus alongside protons.

  • Variation: The number of neutrons in a given element can vary without changing the element's fundamental identity.

  • Carbon can exist with $6$ neutrons or $7$ neutrons, but as long as it retains $6$ protons, it remains Carbon.

  • Variations in neutron number are utilized in advanced medical science and diagnostic imaging.

• Nucleus of an Atom:

  • The central compartment containing all protons and neutrons.

  • It is distinct from the biological nucleus of a eukaryotic cell.

  • There is no physical outer membrane holding the subatomic particles together inside the atomic nucleus; they are simply concentrated at the center.

• Electrons:

  • Charge: Negatively charged ($-$).

  • Location: Positioned outside the nucleus, constantly orbiting around it at designated distances.

  • Electrostatic Attraction: Electrons remain in orbit around the nucleus because opposite electrical charges attract (negatively charged electrons are attracted to positively charged protons in the nucleus).

  • Variation: The total number of electrons associated with an atom can vary without altering the element's identity.

• Periodic Table Organization:

  • The periodic table arranges chemical elements sequentially from left to right based on increasing proton count (atomic number).

  • On standard periodic reference charts, the proton number is listed at the bottom of each element block.

  • No two distinct elements share the same atomic number.

Electron Organization, Shells, and Orbitals

• Electron Shells:

  • Defined as specific concentric distances or energy levels moving outward from the nucleus.

  • Carbon possesses $2$ distinct electron shells: an inner shell closer to the nucleus and an outer shell further away.

  • Capacity Rules for Electron Shells:

  • First Shell (innermost): Can hold a maximum of $2$ electrons (occupancy ranges from $0$ to $2$).

  • Second Shell: Can hold a maximum of $8$ electrons.

  • Third Shell: Operates identically to the second shell in introductory microbiology, holding up to $8$ electrons.

• Electron Orbitals:

  • Within any given electron shell, electrons travel along specific three-dimensional pathways called orbitals.

  • Orbital Capacity and Distribution:

  • Each individual orbital can hold a maximum of $2$ paired electrons.

  • The First Shell contains $1$ orbital (holding up to $2$ electrons).

  • The Second Shell contains $4$ orbitals (holding up to $8$ electrons total).

  • The Third Shell contains $4$ orbitals (holding up to $8$ electrons total).

  • Electron Repulsion and Pairing Rules:

  • Because electrons are all negatively charged, they exert repulsive forces on one another.

  • When filling the $4$ orbitals of the second or third shells, place $1$ electron into each orbital first before pairing them up.

  • Once all $4$ orbitals contain $1$ single electron, any additional electrons are paired up into the existing orbitals (up to the $2$ electron limit per orbital).

• Step-by-Step Procedure for Drawing Atomic Structures:

  • Step 1: Represent the central nucleus as a simplified circle or central sphere.

  • Step 2: Draw the innermost electron shell around the nucleus. Fill it with up to $2$ electrons paired together in its single orbital.

  • Step 3: Draw the second electron shell. Distribute electrons singly across the $4$ orbital locations (top, bottom, left, right) before adding a second electron to any orbital.

  • Step 4: If necessary, draw the third shell following the same orbital distribution rules as the second shell.

• Structural Drawing Examples:

  • Phosphorus ($P$):

  • Total Electrons: $15$

  • First Shell: $2$ electrons (paired in $1$ orbital, completely full).

  • Second Shell: $8$ electrons ($4$ fully paired orbitals, completely full).

  • Third Shell: $5$ remaining electrons ($1$ paired orbital containing $2$ electrons, and $3$ single unpaired orbitals containing $1$ electron each).

  • Nitrogen ($N$):

  • Total Electrons: $7$

  • First Shell: $2$ electrons (paired in $1$ orbital, completely full).

  • Second Shell: $5$ remaining electrons ($1$ paired orbital containing $2$ electrons, and $3$ single unpaired orbitals containing $1$ electron each).

Valence Shells and Chemical Stability

• Definition of Valence Shell:

  • The outermost electron shell of an atom that contains electrons, located furthest from the nucleus.

• Conditions for Chemical Stability:

  • Full Valence Shell: The most chemically stable state for any atom is to possess a completely filled valence shell (a characteristic seen in noble gases on the far right of the periodic table).

  • Orbital Stability: Individual orbitals achieve stability when they contain either $0$ electrons (empty) or $2$ electrons (full). An orbital containing a single ($1$) unpaired electron is unstable and reactive.

  • Driving Force of Chemical Reactions: Atoms interact and form chemical bonds specifically to fill their unpaired orbitals and achieve a full valence shell.

Chemical Bonding Mechanics

• Primary Driving Force of Bond Formation:

  • Atoms undergo chemical bonding to increase their stability relative to their unbonded states by completing their valence shell configuration.

• Ionic Bonds:

  • Mechanism: Occurs when an electron is completely transferred from one atom's shell to another atom's shell.

  • Formation of Ions:

  • An ion is an atom carrying a net electrical charge due to an imbalance between total protons and total electrons.

  • Cations: Positively charged ions formed when an atom loses one or more electrons (protons outnumber electrons).

  • Anions: Negatively charged ions formed when an atom gains one or more electrons (electrons outnumber protons).

  • Bond Attraction: Opposites attract. The electrostatic attraction between a positively charged ion and a negatively charged ion draws them together to form an ionic bond.

  • Example - Table Salt (Sodium Chloride, $NaCl$):

  • Sodium ($Na$) has $1$ electron in its outer third shell. It transfers this single electron to Chlorine ($Cl$), which has $7$ electrons in its outer shell.

  • Sodium becomes a positively charged ion ($Na^+$), and Chlorine becomes a negatively charged chloride ion ($Cl^-$).

  • The electrostatic attraction between $Na^+$ and $Cl^-$ forms ionic bonds that aggregate into a crystalline salt lattice.

• Covalent Bonds:

  • Mechanism: Occurs when two atoms share one or more pairs of electrons rather than transferring them completely.

  • Orbital Overlap: The shared electrons travel in a shared orbital that orbits around the nuclei of both participating atoms.

  • Electrical Balance: No net ions are produced during covalent bonding. Because total proton counts equal total electron counts across the bonded molecule, the overall system remains electrically neutral on average.

  • Examples of Covalent Bonding:

  • Hydrogen Gas ($H_2$): Two Hydrogen atoms, each with $1$ unpaired electron in their first shell, come together to share $1$ pair of electrons, completing the single orbital of their first shell.

  • Water ($H_2O$): An Oxygen atom has $2$ unpaired electrons in its valence shell. It forms two separate single covalent bonds with two individual Hydrogen atoms.

  • Methane ($CH_4$): A Carbon atom has $4$ unpaired electrons in its outer shell. It forms four separate single covalent bonds with four individual Hydrogen atoms.

• Multiple Covalent Bonds:

  • Single Bond: A chemical bond where two atoms share $1$ pair of electrons ($1$ shared orbital).

  • Double Bond: A chemical bond where two atoms share $2$ pairs of electrons ($2$ shared orbitals).

  • Example - Oxygen Gas ($O_2$): Two Oxygen atoms share two pairs of electrons to complete both of their valence shells.

  • Triple Bond: A chemical bond where two atoms share $3$ pairs of electrons ($3$ shared orbitals).

  • Example - Nitrogen Gas ($N_2$): Two Nitrogen atoms share three pairs of electrons across three overlapping orbitals to achieve stability.