Comprehensive notes on isotopes, atomic structure, periodic trends, ions, and naming conventions
Subatomic Particles, Isotopes, and Nuclear Composition
Isotopes are variants of the same element that differ in the number of neutrons in the nucleus. The key idea is that isotopes have the same number of protons (and thus the same atomic number) but different neutron counts, which changes the mass but not the identity of the element. In general, the mass number A of an atom is the sum of its protons and neutrons: where Z is the atomic number (the number of protons) and N is the number of neutrons. The nucleus contains protons and neutrons, collectively called nucleons, and the electrons form a cloud surrounding the nucleus. Protons and neutrons have mass and reside in the nucleus, while electrons have negligible mass by comparison and surround the nucleus in the electron cloud. This layout is central to understanding Rutherford’s alpha-particle scattering experiment: when alpha particles (which are positively charged) were fired at thin metal foils, most went through the foil with little deflection, indicating that atoms are mostly empty space with a small, dense, positively charged nucleus at the center. If atoms were a uniform dense sphere, the alpha particles would have all bounced back. Instead, the results pointed to a nucleus containing protons (and neutrons) and most of the atom’s mass concentrated there.
Atomic Structure, Charge Balance, and Neutrons/Protons
In neutral atoms, the number of electrons equals the number of protons, so the overall charge is zero. If electrons are gained or lost, the atom becomes an ion. A sodium atom (atomic number 11) with a loss of one electron becomes Na⁺, a positively charged cation. A chlorine atom (atomic number 17) that gains one electron becomes Cl⁻, a negatively charged anion. The relative sizes of ions change with electron gain/loss: a cation (positive ion) generally has a smaller ionic radius than the neutral atom, while an anion (negative ion) has a larger ionic radius than the neutral atom. Conceptually, you can imagine a tug-of-war: when electrons are lost, protons pull more strongly on the remaining electrons, shrinking the electron cloud; when electrons are gained, the electron cloud expands because there are more electrons repelling one another.
In chemistry, we distinguish between monoatomic ions (single atoms with a charge) and polyatomic ions (groups of atoms carrying an overall charge). The most common monoatomic cations are alkali and alkaline-earth metals (e.g., Na⁺, Mg²⁺); polyatomic cations include ammonium, NH₄⁺. The periodic table helps predict typical charges: alkali metals tend to form +1 ions, alkaline earth metals +2, and halogens typically form -1 ions. Transition metals can form a range of charges, necessitating special notation when those metals form ions.
Periodic Table Structure and Key Trends
The periodic table is organized into periods (horizontal rows) and groups (vertical columns). A period represents a repeating set of properties as you move across the table, while a group contains elements with similar chemical behavior. Hydrogen is often placed with nonmetals at the top left, though some tables place it with the alkali metals; this placement can vary by table.
Elements are broadly categorized as metals, nonmetals, and metalloids. In the common color-oriented schematics: metals occupy a big block on the left and center (green in the slide), nonmetals (blue) occupy the upper-right portion and include hydrogen, and metalloids (purple) lie along the border between metals and nonmetals. These groups matter because they influence bond formation:
Metals tend to lose electrons and form cations.
Nonmetals tend to gain or share electrons and form anions or covalent bonds.
Metalloids typically show mixed properties and can participate in various bonding types.
Important global groups include the alkali metals (Group 1), the alkaline earth metals (Group 2), the transition metals (middle block), the halogens (Group 17), and the noble gases (Group 18). The transition metals are notable for their variable oxidation states, which makes their chemistry rich and useful in materials science. The halogens typically form -1 anions, and the noble gases are largely inert. As you move down a group, reactivity can increase for many metals, and some properties shift (e.g., metallic character, ionization energy).
Within this table, the shape of bonding can often be anticipated: metals and nonmetals bonded together tend to form ionic compounds through electron transfer (leading to lattice solids), while nonmetals bonded to nonmetals tend to form covalent bonds (molecular compounds).
Molecular and Ionic Compounds; Diatomic and Polyatomic Molecules
Molecules are discrete assemblies of two or more atoms held together by bonds. When a molecule contains exactly two atoms, we call it a diatomic molecule (prefix di- means two). Examples include CO (carbon monoxide) and O₂ (dioxygen). In contrast, molecules with more than two atoms, such as H₂O (water) or CH₄ (methane), are polyatomic molecules.
A significant observation from the lecture is that many of the drawn molecular structures focus on nonmetals, with diverse examples such as CH₄, H₂O, CO, CO₂, and O₃ (ozone). When atoms of metals and nonmetals combine, the result is typically an ionic compound characterized by a lattice structure in which cations and anions are arranged to neutralize overall charge. In contrast, molecular compounds (composed of nonmetals) form discrete molecules with shared electron pairs rather than a full lattice of ions.
Ions: Cations, Anions, and Ionic Radii
An ion is a charged species formed when electrons are transferred or shared in a way that yields a net charge. If an atom loses electrons, it forms a cation (positive charge). If it gains electrons, it forms an anion (negative charge). For example, sodium (Na, Z = 11) loses an electron to form Na⁺, while chlorine (Cl, Z = 17) gains an electron to form Cl⁻. The ionic radius is generally smaller than the atomic radius for cations and larger for anions, due to the changing balance of electrostatic attraction in the electron cloud.
In ionic compounds, the charges must balance to produce a neutral formula unit. For example, Na⁺ combines with Cl⁻ in a 1:1 ratio to give NaCl; Fe²⁺ with Cl⁻ yields FeCl₂; Fe³⁺ with Cl⁻ yields FeCl₃. When a polyatomic anion is involved, the total negative charge must balance the positive charges of the cation(s). Common polyatomic anions include carbonate (CO₃²⁻), sulfate (SO₄²⁻), nitrate (NO₃⁻), nitrite (NO₂⁻), phosphate (PO₄³⁻), and hydroxide (OH⁻). Ammonium (NH₄⁺) is a well-known polyatomic cation.
Empirical vs Molecular Formulas; Balanced Formulas for Ionic Compounds
A molecular formula shows the exact number of atoms of each element in a discrete molecule (e.g., H₂O, CH₄, CO₂). An empirical formula shows the simplest whole-number ratio of the elements in a substance (e.g., glucose has empirical formula CH₂O because its molecular formula is C₆H₁₂O₆). For some substances, the molecular and empirical formulas are the same (e.g., H₂O). For others like glucose, they differ, illustrating the underlying composition.
To determine the empirical formula from a molecular formula, divide all subscripts by their greatest common divisor. For example, C₆H₁₂O₆ reduces to CH₂O. For NO₂₄? (as a hypothetical), dividing by the greatest common divisor would yield the simplest ratio if applicable; for some ionic substances the concept of an empirical formula is more fundamental because the formula unit represents the simplest whole-number ratio that yields electrical neutrality.
When constructing ionic formulas, you must ensure charge neutrality. If a cation with charge +m combines with an anion with charge -n, you choose integers a and b such that The empirical formula is then determined by the smallest whole-number ratio that balances the charges. Examples:
Sodium chloride: NaCl (Na⁺ with Cl⁻; 1:1).
Magnesium bromide: MgBr₂ (Mg²⁺ with Br⁻; need two Br⁻).
Iron(II) chloride vs Iron(III) chloride: FeCl₂ (Fe²⁺) and FeCl₃ (Fe³⁺). For transition metals with multiple oxidation states, Roman numerals denote the specific charge: iron(II) chloride vs iron(III) chloride.
Ammonium nitrate: NH₄NO₃ (NH₄⁺ and NO₃⁻).
For ionic compounds with polyatomic ions, the cation name is followed by the polyatomic anion name (e.g., sodium sulfate Na₂SO₄; calcium phosphate Ca₃(PO₄)₂). When a polyatomic ion carries a charge, that charge is part of the overall neutrality condition that determines the formula unit rather than being written in the name unless needed for disambiguation (e.g., transition metals).
Naming Ionic and Molecular Compounds
Naming conventions distinguish ionic compounds (metal + nonmetal or polyatomic ions) from molecular compounds (two or more nonmetals). For ionic compounds, the cation is named first (the metal), followed by the anion. If the metal is a transition metal with variable oxidation states, a Roman numeral in parentheses indicates the charge of the cation, e.g., FeCl₂ is iron(II) chloride and FeCl₃ is iron(III) chloride. For monoatomic anions, the suffix -ide is added to the element name (e.g., chloride, oxide, sulfide).
In polyatomic ions, the name of the ion is used in full (e.g., sulfate SO₄²⁻, nitrate NO₃⁻, hydroxide OH⁻). When salts incorporate a polyatomic anion with varying hydrogen content, systematic naming can involve hydrogen-containing forms (e.g., bicarbonate, HCO₃⁻, though common naming often uses hydrogen carbonate). For ammonium-containing salts, the cation is ammonium (NH₄⁺). A few examples illustrated in the material include:
Potassium nitrate: KNO₃ (K⁺ + NO₃⁻).
Sodium sulfate: Na₂SO₄ (Na⁺ and SO₄²⁻).
Calcium phosphate: Ca₃(PO₄)₂ (Ca²⁺ with PO₄³⁻).
For molecular compounds (nonmetals bonding to nonmetals), prefixes are used to indicate the number of atoms of each element, and the element farther to the left in the periodic table is named first (with the one lower in the table taking precedence if elements are at the same horizontal level). The second element is given an ending of -ide, with prefixes indicating the count (mono-, di-, tri-, tetra-, etc.). There are several nuanced rules observed in practice, such as omitting the prefix mono- for the first element and sometimes omitting the prefix on the first element depending on common usage. Examples mentioned include:
Carbon monoxide vs carbon dioxide (CO vs CO₂).
Nitrogen trichloride (NCl₃) and phosphorus pentachloride (PCl₅) as common examples.
The general method: name the element on the left first, then the second element with the appropriate prefix and -ide ending. If there is only one of the first element, the prefix mono- is generally not used (e.g., carbon monoxide, not monocarbon monoxide).
Some practical examples from the material:
Hydrogen iodide: HI (hydrogen + iodine).
Nitrogen trichloride: NCl₃.
Sulfur dioxide: SO₂.
Dinitrogen tetroxide (note the more common standard name is dinitrogen tetroxide, N₂O₄).
The warning example in the lecture about “dinitrogen hydroxide” reflects a non-standard or mistaken naming in that moment; the standard names use oxides and nitrates as shown above.
A few other naming conventions mentioned include:
For polyatomic ions, the ending -ide is used on the anion when it is a simple monoatomic ion, as in chloride (Cl⁻), oxide (O²⁻), sulfide (S²⁻), nitride (N³⁻).
For oxyanions (ions containing oxygen such as carbonate, sulfate, sulfite, phosphate, nitrate, nitrite), the standard names are carbonate (CO₃²⁻), sulfate (SO₄²⁻), sulfite (SO₃²⁻), phosphate (PO₄³⁻), nitrate (NO₃⁻), nitrite (NO₂⁻), and so on.
Hydride (H⁻) is the anion formed from hydrogen when it gains an electron; hydride is named hydride in salts such as sodium hydride, NaH.
In summary, the notes emphasize two broad categories: ionic compounds (metal + nonmetal or polyatomic ions) and molecular compounds (nonmetals only). Ionic compounds are named by listing the cation first and the anion second, with Roman numerals used for transition metals to indicate the charge. Molecular compounds are named using prefixes to indicate how many atoms of each element are present, with the element on the left in the periodic table named first and the second element named with a prefix and the -ide ending for the second element. These rules enable consistent naming across a wide range of compounds and help connect chemical formulas with their actual compositions and structures.