Chemistry Review: Atoms, Bonding, Water, and Solutions
Atomic Structure and Isotopes
Atoms are the basic units of matter; have a nucleus containing protons (positive charge) and neutrons (neutral), with electrons (negative charge) orbiting the nucleus in electron orbitals (energy shells).
Elements are pure forms of matter; arranged on the periodic table by atomic number; each element has a chemical symbol and an atomic mass (often shown on the periodic table square). In biology, the bottom-right number of the element’s square is the atomic mass used when calculating molecular weights and molarity.
About 98% of living matter is composed of a small set of elements; these are the ones most often seen in biological molecules.
Isotopes are different forms of the same element with the same number of protons but different numbers of neutrons. Examples (for hydrogen):
${}^{1}_{1}\mathrm{H}$ (1 proton, 0 neutrons)
${}^{2}_{1}\mathrm{H}$ (1 proton, 1 neutron)
(1 proton, 2 neutrons)
Isotopes are heavier forms and many are radioactive.
Uses of isotopes in biology and medicine:
Carbon dating: tracking decay over time to date fossils and ancient samples.
Imaging activity in cells: radioisotopes used in imaging techniques to measure uptake and processing rates (e.g., brain activity scans).
Medical applications: radioisotopes help image or treat disease; higher uptake can indicate faster proliferation in tumors.
Basics of electrons and shells:
Electrons travel in orbitals (electron shells/energy levels) around the nucleus.
First shell (orbit) can hold up to 2 electrons; each subsequent shell can hold up to 8 electrons.
The outermost shell determines an atom’s reactivity; atoms tend to make their outer shell contain 8 electrons (the octet rule).
Outer-shell/valence concepts:
Atoms “want” a full outer shell of 8 electrons (octet).
To achieve this, atoms can gain, lose, or share electrons.
Key terms:
Chemical bond: linking of atoms via interactions of electrons.
Molecule: two or more atoms bonded together.
Compound: a molecule composed of two different elements bonded together (not every molecule is a compound; e.g., O2 is a molecule but not a compound).
Molecular weight and molarity: to prepare solutions with correct molarity, you need the molecular weight of the solute, which comes from summing the atomic weights of constituent atoms.
Electron Configuration, Bonding, and Molecular Polarity
Electron shells and capacity (recap):
First shell capacity: 2 electrons
All subsequent shells capacity: 8 electrons
Outer-shell (valence) electrons determine reactivity; atoms gain, lose, or share electrons to fulfill the octet rule.
Covalent bonds:
Formed by sharing a pair (or pairs) of electrons between two atoms.
Covalent bonds are generally strong.
Bond types by electron sharing:
Single bond: shared one pair of electrons.
Double bond: shared two pairs of electrons.
Triple bond: shared three pairs of electrons.
Examples:
H2: two hydrogen atoms share one pair of electrons (single covalent bond).
CH4: carbon shares four electrons with four hydrogens to complete its outer shell.
When electrons are not shared equally (electronegativity differences), a polar covalent bond forms.
Electronegativity and polarity:
Electronegativity is the tendency of an atom to attract electrons.
If one atom in a covalent bond pulls electrons more strongly, the bond is polar, giving partial charges:
Atom with higher electronegativity becomes partially negative.
Atom with lower electronegativity becomes partially positive.
Polar molecules have a net dipole and are usually hydrophilic (water-loving).
Nonpolar molecules (e.g., lipids) have equal sharing and are hydrophobic (water-fearing).
Hydrogen bonds:
Weak bonds formed between molecules that have polar covalent bonds.
In hydrogen bonds, the partially positive hydrogen end of one molecule interacts with a partially negative atom (commonly oxygen, nitrogen, or fluorine) in another molecule.
Importance:
Hold the double helix of DNA together via multiple hydrogen bonds.
Stabilize protein structures and many macromolecular conformations.
Confer unique properties to water, enabling cohesion and particular solvent behavior.
Ionic bonds:
Formed by complete transfer of one or more electrons from one atom to another.
Resulting ions with opposite charges attract and form an ionic bond.
Ions:
Cation: positively charged ion (loses electrons).
Anion: negatively charged ion (gains electrons).
Example: Sodium chloride (NaCl)
Na tends to lose one electron to achieve a stable configuration; Cl accepts that electron, forming Na+ and Cl−, which are held together by ionic attraction.
Chemical reactions and balance:
Reactions can be depicted by chemical equations with reactants on the left and products on the right, connected by an arrow.
Matter is conserved; equations must be balanced (same number of each type of atom on both sides).
Example (combustion):
Reactants: propane and oxygen; Products: carbon dioxide, water, heat, and light.
Balanced equation:
Energy changes in chemical reactions:
Two main energy types:
Potential energy: stored energy (e.g., ATP stores energy in phosphate bonds).
Kinetic energy: energy in use (energy of motion and chemical processes in action).
ATP (adenosine triphosphate) stores energy in its phosphate bonds; hydrolysis releases energy useful for cellular work:
The release and/or transfer of energy from ATP powers many cellular processes.
Water, Solutions, and Their Biological Significance
Water as the universal solvent:
Many biological reactions occur in aqueous solutions; water’s properties govern solubility, transport, and reaction rates.
Aqueous solution: when water is the solvent; solute is the substance dissolved in the solvent.
Example: Kool-Aid sugar (solute) dissolved in water (solvent) to form a solution.
Hydrophilic vs hydrophobic:
Polar molecules are typically hydrophilic and dissolve in water.
Nonpolar molecules (like lipids) are hydrophobic and tend to aggregate in water.
Four unique properties of water (that make it special for life):
1) Ice’s density is lower than liquid water (ice floats):Expression:
ho{ ext{ice}} < ho{ ext{liquid water}}Consequence: aquatic life can survive under frozen surfaces; ice forms a protective insulating layer.
2) High specific heat capacity: water requires a lot of energy to change its temperature, helping to stabilize body and environmental temperatures.Concept:
Implication: water buffers temperature changes in organisms and ecosystems.
3) High heat of vaporization: water requires substantial energy to vaporize, contributing to cooling via evaporation (e.g., sweating).Concept: energy required to convert liquid water to vapor; related to A H_vap
Biological relevance: sweating and panting cool the body.
4) Cohesive strength and surface tension: strong hydrogen bonding between water molecules makes water cohesive and creates surface tension.Consequences: allows capillary movement in plants (transpiration) and stable droplets on surfaces; enables transport of water through soil and blood.
Capillary action and transpiration:
Cohesion (water-water H-bonds) plus adhesion (water to other substances) enable water to move upward in narrow tubes (capillaries) and through plant xylem.
Solubility terms and practical definitions:
Solvent: the dissolving medium (water in aqueous solutions).
Solute: the substance dissolved in the solvent.
Real-world implications:
Water’s properties underlie biological temperature regulation, nutrient transport, and molecular interactions that shape protein folding, DNA structure, and cell signaling.
Connections to prior principles and real-world relevance:
Understanding atomic structure and bonding helps explain how macromolecules form (proteins, nucleic acids).
Isotope use links to dating, imaging, and cancer diagnostics, illustrating how physics and chemistry intersect with biology.
Water properties explain why life depends on aqueous environments and why organisms adapt to temperature and hydration changes.