Notes on Elements, Atoms, Compounds, and Chemical Bonds

Course logistics and focus

• Sunday: complete the student success assessment; this is the only due item for the week.
• Monday: possible full-class session; we may chart Chapter 3 (macromolecules).
• Wednesday: in-person quiz in class; on paper; 5–7 questions; mostly multiple choice, may include matching.
• Book sections: relevant to the quiz are the Chapter sections that cover chemical principles and, in particular, the properties of water.
• Overall goal for this portion of the lecture: grasp the basics of elements, atoms, compounds, and the foundational properties of water; connect to later topics on macromolecules and life processes.
• Encourage questions at any time.

Elements, matter, compounds, and life chemistry

• Matter in living organisms is composed of chemical structures built from atoms.
• Matter is made of chemical elements; an element is a substance that cannot be broken down further by ordinary chemical means.
• There are 92 natural elements.
• An element example: Oxygen (O) is essential for respiration; Carbon (C), Hydrogen (H), and Nitrogen (N) are also fundamental.
• Some elements exist in pure form only rarely; most are found in compounds within organisms.
• About 25 elements are present in living organisms, with the four most abundant elements contributing roughly 96% of body weight: O, C, H, and N.
• The four most important elements in the body (by abundance) are:
  • Oxygen: ≈ 65% of body mass
  • Carbon: significant portion
  • Hydrogen: significant portion
  • Nitrogen: significant portion
• Other essential elements (e.g., calcium, phosphorus, magnesium, sodium) are present in smaller amounts but are still crucial for function.
• The six elements that collectively account for about 99% of body mass are typically listed as: ext{O}, ext{C}, ext{H}, ext{N}, ext{Ca}, ext{P}.
• Life’s chemistry relies on the way elements combine to form compounds, which are substances made of two or more elements in a fixed ratio.
• Example: glucose (a carbohydrate) is a compound composed of carbon, hydrogen, and oxygen in a fixed ratio.
• Matter and chemistry also hinge on the distinction between elements and compounds: elements are pure substances, compounds are formed by chemical bonding of two or more elements.

The atom: subatomic particles and basic properties

• The atom is the smallest unit of an element that retains its properties.
• An atom consists of three main types of subatomic particles:
  • Protons: positively charged, located in the nucleus.
  • Neutrons: electrically neutral, located in the nucleus.
  • Electrons: negatively charged, orbiting the nucleus in electron shells.
• Charges of subatomic particles: protons (+), electrons (−), neutrons (0).
• The nucleus contains protons and neutrons; electrons form a cloud around the nucleus.
• In a neutral atom, the number of protons equals the number of electrons, balancing the charge.
• Key terms:
  • Atomic number, Z: the number of protons in the nucleus; defines the element.
  • Mass number, A: the sum of protons and neutrons, i.e. A = Z + N0 where N0 is the number of neutrons.
  • Atomic mass: the overall mass of the atom; closely approximates the mass number A because electrons contribute negligibly to mass.
• Isotopes: atoms of the same element (same Z) with different numbers of neutrons, yielding different mass numbers; e.g., carbon-12, carbon-13, carbon-14.
• Example discussion: Carbon has Z = 6 protons; common isotopes include ^{12}{6} ext{C}, ^{13}{6} ext{C}, ^{14}_{6} ext{C}.

Electron shells, valence, and periodic trends

• Electrons occupy electron shells around the nucleus; shells have capacity limits:
  • The first shell can hold up to 2 electrons.
  • The second shell can hold up to 8 electrons.
  • All subsequent shells can hold up to 8 electrons (in the simplified, commonly taught model).
• The outermost shell is the valence shell; electrons in this shell are valence electrons and largely determine chemical reactivity.
• Atoms tend to have a full valence shell; they interact with other atoms to gain, lose, or share electrons in order to achieve this stability.
• The periodic table provides a visual summary:
  • Each row corresponds to the number of electron shells (periods).
  • Each column reflects the number of electrons in the outer (valence) shell for the representative elements.
  • Noble gases (e.g., He, Ne, Ar) have full valence shells and are highly unreactive.
• Understanding atomic structure helps explain why elements form bonds and why some elements are more reactive than others.
• A common exercise: deducing atomic number from an electron configuration diagram; e.g., Lithium has 3 protons (Z = 3) because its neutral atom has 3 electrons; Magnesium has Z = 12 because its neutral atom has 12 electrons in the diagram.
• Electron distribution and shell filling influence chemical behavior and bonding potential.

Chemical bonds: covalent, ionic, and hydrogen bonds

Covalent bonds

• Covalent bonds involve sharing electrons between atoms to fill outer electron shells.
• Examples:
  • Hydrogen gas: ext{H}_2, formed by sharing one electron pair between two hydrogen atoms.
  • Methane: ext{CH}_4, carbon shares electrons with four hydrogen atoms to achieve full outer shells.
  • Water: ext{H}_2 ext{O}, covalent bonds between O and H; can be polar.
• Types of covalent bonds:
  • Nonpolar covalent bond: electrons shared equally; e.g., H–H in ext{H}_2, the C–H bonds in methane are often treated as largely nonpolar due to similar electronegativities.
  • Polar covalent bond: electrons shared unequally, leading to partial charges on atoms; e.g., in water, the electrons spend more time near oxygen, giving a partial negative charge on O and partial positive charges on Hs.
• Represented schematically by lines between atoms in structural formulas.
• Water is a key example of a polar covalent bond due to electronegativity differences between oxygen and hydrogen.

Ionic bonds

• Ionic bonds form when one atom donates one or more electrons to another, creating positively and negatively charged ions (cations and anions).
• The electrostatic attraction between oppositely charged ions stabilizes the bond.
• Common example: sodium chloride, NaCl.
  • Sodium (Na) donates one electron to chlorine (Cl).
  • After donation/acceptance: Na becomes Na⁺ (cation), Cl becomes Cl⁻ (anion).
  • The resulting ionic attraction holds the compound together as a salt.

Hydrogen bonds

• Hydrogen bonds are weak interactions that occur between a hydrogen atom covalently bonded to a highly electronegative atom (such as O, N, or F) and another electronegative atom in a nearby molecule.
• In water, a single water molecule can form hydrogen bonds with multiple neighboring water molecules:
  • The partially positive hydrogen of one water molecule is attracted to the partially negative oxygen of another water molecule, and vice versa for the other hydrogen.
• Individual hydrogen bonds are weak, but collectively they can be strong and drive many emergent properties of water and biological systems.
• Hydrogen bonds are not bonds that hold atoms within a molecule together (that is the role of covalent or ionic bonds); instead, they stabilize interactions between molecules (intermolecular interactions).

Water: polarity and emergent properties (preview for later sections)

• Water's polar covalent bonds create a dipole; the molecule has a partial negative charge on the oxygen and partial positive charges on the hydrogens.
• This polarity enables:
  • Hydrogen bonding between water molecules
  • Cohesion (water molecules sticking to each other)
  • Adhesion (water sticking to other surfaces)
  • Hydrophilic vs. hydrophobic interactions (water-loving vs. water-hating substances)
• These properties are foundational to life and are discussed in more depth in the next part of the course.

Chemical reactions: bonds, reactants, and products

• A chemical reaction involves breaking existing bonds and forming new bonds to rearrange atoms.
• Reactants are the starting substances; products are what is formed after the reaction.
• In a closed system, atoms are conserved; reactions rearrange electrons and bonds rather than creating or destroying atoms or mass.
• Example types mentioned:
  • Formation of water from hydrogen and oxygen via covalent bonding.
  • Formation of table salt (NaCl) via ionic bonding.
• Note: In the context of this course, understand that there are three major types of bonding (covalent, ionic, hydrogen) and how they contribute to the structure and properties of molecules.

Quick recap of key terms and concepts to study

• Elements and compounds: definitions and distinctions; 92 naturally occurring elements; fixed ratios in compounds.
• Matter and life chemistry: atoms are the building blocks of matter; cells are comprised of molecules.
• Atomic structure: nucleus (protons + neutrons) and electron cloud; charges; atomic number Z; mass number A; isotopes.
• Electron shells and valence: capacity 2, 8, 8, …; valence electrons determine bonding behavior.
• Periodic trends: rows denote electron shells; columns reflect outer-shell electron counts; noble gases are inert due to full shells.
• Covalent bonds: sharing electrons; nonpolar vs polar; H–H, C–H, O–H examples; methane vs water.
• Ionic bonds: electron transfer; formation of ions; NaCl example.
• Hydrogen bonds: weak inter-molecular attractions; crucial for water's properties and biological interactions.
• Chemical reactions: rearrangement of electrons and bonds; conservation of matter.

Short practice prompts (exam-style)

• Which bond uses shared electrons? Answer: Covalent bonds.
• In a simple electron-dot diagram, if an atom has 3 electrons in its outer shell, what is the atomic number (roughly) and what does that imply about bonding tendency? Hint: number of protons equals the atomic number; this atom tends to form bonds to fill its outer shell unless it is a noble gas.
• What is the maximum number of electrons in the first electron shell? Answer: 2 electrons.
• What is the maximum number of electrons in the second shell? Answer: 8 electrons.
• Name two common isotopes of carbon and their mass numbers. Answer: ^{12} ext{C} and ^{13} ext{C} (and ^{14} ext{C} as another isotope).
• Give an example of a polar covalent bond and explain why it is polar. Answer: in ext{H}_2 ext{O}, oxygen is more electronegative than hydrogen, so electrons are pulled toward O, creating partial charges on O and H.