Chemistry Notes: Molecules, Ions, and Formulas

Molecules: Definition and Types

A molecule is two or more atoms bonded together by chemical forces.
Molecules can be composed of the same kind of atoms (homoatomic) or different kinds (heteroatomic).
- Examples:
- Homoatomic diatomic molecules:
  ext{O}2 ext{, Cl}2, ext{H}_2
  ight) as diatomic molecules.
- Heteroatomic diatomic molecule: $ext{HCl}$ (hydrogen chloride).
A diatomic molecule contains exactly two atoms. It can be the same or different types; the number of atoms is two.
A polyatomic molecule contains more than two atoms (e.g., water: $ext{H}2 ext{O}$ ; ammonia: $ext{NH}3$ ; carbon dioxide: $ext{CO}2$ ; methane: $ext{CH}4$ ).
Diatomic elements (seven) exist naturally as diatomic molecules when in pure elemental form (mostly gases, bromine is an exception as a liquid):
- $ext{H}2, ext{N}2, ext{O}2, ext{F}2, ext{Cl}2, ext{Br}2, ext{I}_2$
Mnemonic to remember the seven diatomic elements (as taught): "Have no fear of ice cold here" or a similar variant; there are multiple mnemonics.
Diatomic elements are especially stable in their diatomic form in nature when they are pure elements.
Special case: Bromine (Br) is liquid at room temperature, whereas the others are gases (mostly).
Notation reminders:
- H, N, O, F, Cl, Br, I are the seven diatomic elements; when in diatomic form, they appear as $ext{X}_2$ .
- A single element can exist as a diatomic molecule or as an atomic species depending on the element and state.

Ions: Charged Species

An ion is an atom or a group of atoms with a net positive or negative charge.
Two main types by charge:
- Cation: positively charged ion (formed by losing one or more electrons).
- Anion: negatively charged ion (formed by gaining one or more electrons).
Rules for charge formation:
- Protons are positively charged and stay in the nucleus; electrons carry negative charge and can be gained or lost.
- A neutral atom has the same number of protons and electrons.
- Net charge Q = (number of protons) − (number of electrons).
- If electrons are lost: the atom becomes a cation (positive charge).
- If electrons are gained: the atom becomes an anion (negative charge).
Size changes with ion formation:
- Cations are typically smaller than their neutral atoms (loss of electrons reduces electron-electron repulsion and often the effective radius).
- Anions are typically larger than their neutral atoms (extra electrons increase repulsion and size).
Examples:
- Sodium (Na) loses one electron to form Na⁺: ext{Na}
  ightarrow ext{Na}^+ + e^-
- Chlorine (Cl) gains one electron to form Cl⁻: ext{Cl} + e^-
  ightarrow ext{Cl}^-
- Naming:
- Positive ion: sodium cation or sodium ion (Na⁺).
- Negative ion: chloride ion (Cl⁻) – note the -ide ending in many anions applied to the element name.
Periodic-table trend (brief):
- Group 1A metals form +1 ions (lose 1 electron).
- Group 2A metals form +2 ions (lose 2 electrons).
- Group 17 nonmetals form −1 ions (gain 1 electron).
- Group 16 nonmetals form −2 ions (gain 2 electrons).
- Group 15 nonmetals form −3 ions (gain 3 electrons).
- Transition metals can form several possible positive charges (oxidation states).
Polyatomic ions: ions composed of two or more atoms with an overall charge; examples include ammonium, hydroxide, nitrate, sulfate, phosphate, cyanide, etc.
- Monoatomic ions: ions consisting of a single atom (e.g., Na⁺, Cl⁻, Ca²⁺).
- Polyatomic ions: groups of atoms bonded together that carry a net charge (e.g., ammonium NH₄⁺, hydroxide OH⁻, nitrate NO₃⁻, sulfate SO₄²⁻, phosphate PO₄³⁻, cyanide CN⁻).
Common polyatomic ions (names and formulas):
- Ammonium: ext{NH}_4^+
Hydroxide: ext{OH}^-
- Nitrate: ext{NO}_3^-
- Sulfate: ext{SO}_4^{2-}
- Sulfide: ext{S}^{2-}
- Phosphate: ext{PO}_4^{3-}
- Cyanide: ext{CN}^-
Rationale for polyatomic ion stability: grouped atoms share electrons covalently to minimize energy; the resulting group has a net charge that remains bonded as a unit.

Ions in the Periodic Table and Formula Formation

Ionic compounds are typically formed from metal cations and nonmetal anions; the classic example is a compound formed by transfer of electrons from metal to nonmetal (e.g., LiCl, NaF).
The most favorable ionic interactions occur when the metal can donate electrons to form a stable cation and the nonmetal can accept electrons to form a stable anion, often achieving noble-gas-like electron configurations.
Example: Group 1A metal and Group 17 nonmetal form simple ionic salts such as LiF, NaCl, KCl, etc., with a 1:1 ratio in many cases because the metal donates 1 electron and the nonmetal accepts 1 electron.
When a nonmetal requires more electrons (e.g., Group 16 needing 2 electrons) and the metal can only donate one electron, multiple metal atoms may be needed to satisfy the nonmetal’s electron requirement (e.g., Na₂O, K₂O, etc.). This ensures whole-number ratios in compounds, in line with Dalton’s atomic theory which requires whole-number ratios of atoms.
Concept recap: ionic compounds result from electron transfer; the overall formula reflects the smallest whole-number ratio that balances charges.

Worked Examples: Protons, Neutrons, and Electrons in Ions

Notation: ^{A}_{Z} ext{X} where Z is the atomic (proton) number, A is the mass number, N = A − Z is the neutron number, and E is the number of electrons.
Francium (Fr) example from class notes:
- Given: ^{223}_{87} ext{Fr}
- Protons: Z = 87
- Neutrons: N = A - Z = 223 - 87 = 136
- Electrons (neutral): E = Z = 87
Iron example (Fe) in a +2 oxidation state (Fe²⁺) with common isotope A = 56:
- Protons: Z = 26
- Neutrons: N = A - Z = 56 - 26 = 30
- Electrons in neutral Fe: E = Z = 26
- Electrons in Fe²⁺: E = Z - 2 = 24
Potassium example (K) with A = 39 and +1 charge (K⁺):
- Protons: Z = 19
- Neutrons: N = A - Z = 39 - 19 = 20
- Electrons in neutral K: E = Z = 19
- Electrons in K⁺: E = Z - 1 = 18
Indium example (In) with a common ion In³⁺ and isotope A = 111:
- Protons: Z = 49
- Neutrons: N = A - Z = 111 - 49 = 62
- Electrons in neutral In: E = Z = 49
- Electrons in In³⁺: E = Z - 3 = 46
Takeaway: Given Z, A, and simple ion charge, you can compute N and E using the formulas:
- Neutrons: N = A - Z
- Electrons in ion: E = Z - ext{charge} where charge is positive for cations and negative for anions.

Monoatomic vs Polyatomic Ions: Quick Clarifications

Monoatomic ions: single-atom ions such as ext{Na}^+, ext{Cl}^-, ext{Ca}^{2+}.
Polyatomic ions: a fixed group of atoms with a net charge, e.g., ext{NH}4^+, ext{OH}^-, ext{NO}3^-, ext{SO}4^{2-}, ext{PO}4^{3-}, etc.
Naming conventions:
- Cation: often named by the element name (e.g., sodium ion, calcium ion).
- Anion: the element name often changes to end with -ide (e.g., chloride ion, oxide ion, sulfide ion, nitride ion).
Visual cue: metals generally form cations; nonmetals generally form anions; electron transfer leads to stable ionic compounds.

Polar Covalent and Ionic Compounds: Formula and Nomenclature

Chemical formula: shorthand representation of a compound using symbols and subscripts; e.g.,
- Water: ext{H}_2 ext{O}
- Sodium chloride: ext{NaCl}
Nomenclature: systematic naming of compounds so scientists worldwide can communicate unambiguously.
- Example: water is commonly used, but its systematic name is dihydrogen monoxide (H₂O).
- Sodium chloride is named sodium chloride; the second part uses the -ide form for the anion (chloride).
- The difference between formula and nomenclature:
- Formula gives the composition (symbols and subscripts).
- Nomenclature gives the type of compound (and the names of the participating ions/elements).
Foundation roles:
- Formulas and names underpin balanced chemical equations, understanding molecular structure, and lab safety (e.g., sodium cyanide is a poison while table salt is edible).
Types of chemical formulas:
- Molecular formula: shows exact numbers of each type of atom in a molecule (e.g., ext{C}6 ext{H}{12} ext{O}_6 for glucose).
- Empirical formula: simplest whole-number ratio of atoms; for glucose, the empirical formula is ext{CH}_2 ext{O} (divide each subscript by 6).
- In this course, mean focus is often on molecular formulas, though empirical formulas are introduced.

Representations of Molecules: From Formulas to Models

Molecular formula: shows which atoms and how many are in a molecule (e.g., ext{H}2 ext{O}, ext{NH}3, ext{CH}_4).
Structural formula: shows how the atoms are connected (e.g., H–O–H for water; N with three H attached in ammonia; C–H bonds in methane).
Ball-and-stick model: atoms as spheres (balls) connected by sticks (bonds); helps visualize bonds and connectivity.
Space-filling model: shows the relative sizes of atoms and how they occupy space; provides a more realistic sense of molecular volume.
Shapes and geometry remarks:
- Water is a bent (V-shaped) molecule due to bond angles and lone pairs on oxygen.
- Methane is tetrahedral in shape.

Practice: Subscripts vs Coefficients; Counting Atoms and Molecules

Subscript: small number placed after an element symbol within a formula; indicates how many atoms of that element are in a single molecule.
- Example: ext{H}_2 ext{O} has two hydrogen atoms and one oxygen atom per molecule.
Coefficient: number placed in front of a formula; indicates how many molecules are present in the sample.
- Example: 2 ext{H}_2 ext{O} means two molecules of water (total of 4 H and 2 O atoms).
Combined: both a coefficient and subscripts can appear (e.g., 2 ext{H}_2 ext{O} has two molecules of water, totaling 4 H and 2 O atoms).
Quick application (counting atoms and molecules):
- Example results from class discussion:
- 1) 8 atoms present; 4 molecules
- 2) 6 atoms present
- 3) 9 atoms present
- 4) 10 atoms present
Distinguishing elements from compounds:
- A single element (e.g., H₂) is a diatomic molecule, not a compound.
- Compounds contain two or more elements (e.g., H₂O, CO₂).

Quick Reference: Common Concepts and Real-World Relevance

Why ions matter: ionic compounds are widespread (salts, electrolytes, etc.); their formation explains many laboratory and real-world phenomena.
The “energy minimization” principle: the most stable arrangement in polyatomic ions and molecular structures corresponds to the lowest potential energy configuration.
The periodic table as a guide: group trends help predict which elements tend to form cations or anions and typical charges.
The language of chemistry (formula vs nomenclature) is essential for clear communication, safety, and practical lab work.

Exam Prep and Course Context

You will have a study guide focusing on key concepts, with a forthcoming mock test to practice system-proctored exams.
Review topics covered:
- Atomic structure basics (protons, neutrons, electrons; mass number A; atomic number Z)
- Ions: cations vs anions; monoatomic vs polyatomic ions; common polyatomic ions
- Diatomic elements and the seven diatomic molecules
- Ionic bonding and formation of ionic compounds (one electron transfers, whole-number ratios)
- Molecular formulas, empirical formulas, and nomenclature basics
- Different ways to represent molecules (molecular formula, structural formula, ball-and-stick, space-filling models)
Final reminders:
- Practice with counting atoms and determining the number of molecules from formulas
- Use mnemonics and periodic table references to recall ion charges and common ions
- Understand the difference between subscripts (within a molecule) and coefficients (the number of molecules)
- Be comfortable with basic algebraic relationships: A = Z + N $and$ N = A - Z $, and the ion electron count$ E = Z - ext{charge}$$