Chapter 2 Part 1 Notes
Atomic Structure
Nucleus contains Protons (p⁺) and Neutrons (n⁰); Protons determine the Atomic Number (Z) and chemical identity; Neutrons contribute to Atomic Mass and isotope type.
Outside nucleus: Electrons (e⁻) form the Electron Cloud and determine chemical properties.
In a neutral atom, the number of protons equals the number of electrons: Z = N_e.
Electrical attraction between positively charged protons and negatively charged electrons holds the atom together.
Identifying Elements
Atomic Number (Z): number of protons (also equal to number of electrons in a neutral atom).
Average Atomic Mass: weighted average of all isotopes' masses for the element.
Mass Number (A): Z + N, where N is the number of neutrons.
Isotopes differ in N but have the same chemical properties; isotopes have different mass numbers.
Example relationships: Mass Number A = Z + N and Neutrons N = A - Z.
Isotopes and Radioactivity
Isotopes: atoms with same Z but different N; may be stable or unstable (radioactive).
Radioisotopes tend to decay spontaneously to more stable forms, releasing particles/energy.
Diagnostic and therapeutic uses: small amounts of radioisotopes can be used in medicine.
Electron Energy Levels
Electrons occupy energy levels (shells) around the nucleus.
Outer (valence) electrons determine reactivity; inner electrons are less reactive.
Maximum electrons per level: Level 1 = 2, Level 2 = 8, Level 3 = 18, Level 4 = 32.
Electron configuration fills from inner to outer levels (e.g., 1s² 2s² 2p⁶ …).
An energy level that is not completely filled is less reactive unless a stable octet is achieved (often 8 electrons in outer level).
Example fill pattern: octet rule for stability in many main-group elements.
Chemical Bonds & Reactions
Atoms bond by sharing, gaining, or losing electrons to form molecules and compounds.
Bonds trap energy within molecules.
Three major bond types: Ionic, Covalent, Hydrogen.
Bonding and breaking govern metabolism and body function.
Molecules and Compounds
Molecule: two or more atoms bonded together; shape determined by bond arrangement.
Compound: two or more different atoms bonded; may involve shared or transferred electrons, producing ions.
Compounds have new chemical properties different from their constituent elements.
Ion Formation
When atoms gain or lose electrons, they form ions.
Ions carry positive or negative charges: Cations (positive), Anions (negative).
Ionic Compounds and Salts
Ionic compounds (salts) do not contain H⁺ or OH⁻ exclusively.
In solution, they dissociate into ions (e.g., Na⁺, Cl⁻).
Salts form electrolytes in the body; electrolytes conduct current across membranes and must be in specific ranges for homeostasis.
Ionic Bonds
Ionic bonds form when electrons are transferred from one atom to another.
Resulting ions attract to form salts; bonds are relatively strong.
Covalent Bonds
Covalent bonds form when atoms share electrons.
Most common bonds in the body; essential for building biological structures.
Single covalent bonds share one pair of electrons; relatively less energy trapped; easier to break.
Carbon can form up to four covalent bonds, enabling great molecular diversity.
Covalent Bonds & Polarity
Electron movement in covalent bonds can create non-polar or polar molecules.
Non-polar: electrically neutral with even e⁻ distribution.
Polar: uneven distribution of electrons; results in partial positive and negative regions; polar molecules attract other polar molecules and water (like dissolves like).
Hydrogen Bonds
Occur between the hydrogen of one molecule and N, O, or F of another molecule.
Important for protein 3D structure and DNA double helix; they are weaker than ionic or covalent bonds.
Chemical Reactions
A chemical reaction occurs when bonds are broken and/or new bonds form: Reactants → Products.
Products generally reside in a lower-energy state.
Metabolism: sum of all chemical reactions in the body; transfers energy to carry out functions.
Patterns of Chemical Reactions
Synthesis (A + B → AB): atoms/ions/molecules combine; often involves dehydration synthesis; energy absorbed; Anabolism; growth/repair.
Decomposition (AB → A + B): molecules broken down; often hydrolysis; energy released; Catabolism; digestion.
Single & Double Replacement (Exchange) Reactions:
AB + C → AC + B
AB + CD → AD + CB
Involve rearrangement of parts; example: HCl + NaOH → NaCl + H₂O; relevant to ATP ↔ ADP cycle.
Reversible Reactions & Equilibrium
Reactions can be reversible: A + B
ightleftharpoons C + D.At equilibrium, forward and reverse rates are equal; adding more reactants can shift the balance toward product formation (7 Le Chatelier principle in simple terms).
Rate of Chemical Reactions
Rate is influenced by:
Concentration of reactants: higher concentration can increase rate (within limits).
Temperature: higher temperature generally increases rate (within limits).
Catalyst: speeds up the reaction without being consumed.