periodic table: elements combining (chapter 3)
Bonding, Formulas, and Nomenclature — Notes
Is essential distinction: ionic compounds form extended crystal lattices; formula units represent the smallest repeating ratio of ions in the lattice. For NaCl, the ratio is 1:1, so we write the formula as NaCl. In any real salt sample, you’ll have many such units, not a single Na and a single Cl.
Empirical formula vs molecular formula:
Empirical formula = smallest whole-number ratio of atoms in a compound.
Molecular formula = actual number of each type of atom in a molecule.
For ionic compounds, we usually use empirical formulas or formula units; these often match the simplest ratio in the lattice.
For covalently bonded (molecular) compounds, the molecular formula shows the actual numbers in the molecule; sometimes the empirical and molecular formulas are the same (e.g.,
$\text{H}_2\text{O}$), but not always.
Electrostatic potential energy and Coulomb's law (conceptual):
The energy of interaction for two charges is proportional to the product of their charges divided by the distance between them:
The stronger the charges (larger |q|), the stronger the interaction; smaller distances (r) allow charges to come closer, increasing the interaction.
In an ionic bond, a negatively charged ion (anion) and a positively charged ion (cation) attract more strongly as they approach, reach a minimum energy at an equilibrium separation, and then experience repulsion if forced closer due to nuclear repulsion.
The potential energy curve for ionic interactions shows a minimum at the bond distance; beyond this, repulsion rises as atoms get too close.
Note: This class discussion presents Coulomb’s law conceptually; you won't be required to memorize the equation at this stage.
Ionic compounds vs covalent (molecular) compounds: quick distinctions
Ionic compounds: metal cations + nonmetal anions; form extended crystal lattices; empirical formula/ formula unit; examples include NaCl, MgO, CaS.
Covalent (molecular) compounds: nonmetals bonded by covalent (electron-sharing) bonds; form discrete molecules; molecular formulas describe the composition (e.g., $\mathrm{C}2\mathrm{H}6\mathrm{O}$ for ethanol).
Bonding curve analogy: for two nonmetal atoms at large separation, energy is effectively zero; as they approach, there is attraction due to shared electrons forming a covalent bond; a minimum energy occurs at the bond length, $r_{\text{bond}}$; pushing them closer leads to repulsion.
Example molecules and formulas
Sodium chloride, NaCl: ionic compound; formula unit/empirical formula is NaCl; real samples contain many units.
Ethanol, $\mathrm{C}2\mathrm{H}6\mathrm{O}$: covalent molecular compound; molecular formula shows the actual ratio; empirical formula may be the same as molecular in this case.
Water, $\mathrm{H}_2\mathrm{O}$: an example where empirical and molecular formulas coincide.
Carbon allotropes: carbon exists as different forms (allotropes) such as diamond and graphite; both are elemental carbon but with different bonding structures.
Oxygen forms two common elemental diatomic molecules: $\mathrm{O}2$ (oxygen gas) and $\mathrm{O}3$ (ozone).
Diatomic molecules in general: $\mathrm{N}2$, $\mathrm{O}2$, $\mathrm{H}_2$, and the halogen diatomics working in nature.
Allotropy concept: a single element exists in more than one form with different properties (e.g., carbon as diamond vs graphite; oxygen as $\mathrm{O}2$ vs $\mathrm{O}3$).
Elemental molecules vs atomic elements: some elements exist as discrete atoms in nature (atomic elements, e.g., noble gases), while others exist as diatomic or polyatomic molecules (molecular elements, e.g., $\mathrm{O}2$, $\mathrm{N}2$).
Polyatomic ions: charged molecular species that can form ionic compounds (e.g., nitrate, $\mathrm{NO3^-}$; ammonium, $\mathrm{NH4^+}$).
Classification of pure substances (overview)
Elements vs compounds
Elements can be:
Atomic elements (exist as single atoms, e.g., noble gases)
Molecular elements (exist as molecules, e.g., $\mathrm{O}2$, $\mathrm{N}2$, $\mathrm{H}_2$; diatomic halogens also fall here)
Compounds can be:
Molecular compounds (nonmetals bonded covalently; have molecular formulas)
Ionic compounds (metal + nonmetal; rely on ionic bonds; often described by empirical formulas)
Note on allotropes and elemental molecules: allotrope is a different form of the same element (e.g., carbon in diamond vs graphite). The terms $\mathrm{P}4$ and $\mathrm{S}8$ are examples of elemental molecules for phosphorus and sulfur.
Naming binary ionic compounds (nomenclature rules, overview)
Rules (as discussed):
The first word is the cation (the metal, the element from the left side of the periodic table in the compound).
The second word is the anion name, with the suffix changed to -ide (indicating a negatively charged ion).
Examples:
Magnesium chloride: $\mathrm{MgCl_2}$ → Magnesium (cation) + chloride (an anion with -ide suffix)
Potassium bromide: $\mathrm{KBr}$ → Potassium + bromide
Calcium sulfide: $\mathrm{CaS}$ → Calcium + sulfide
Cross-method for balancing charges (simplified explanation):
To ensure overall neutrality, balance the charges of the ions by using the smallest whole-number subscripts that make the total charge zero.
Example: magnesium nitride
Ions: $\mathrm{Mg^{2+}}$ and $\mathrm{N^{3-}}$
Balance by cross-multiplication: the subscripts become $\text{Mg}3\text{N}2$ (because $2$ charges of Mg must balance $3$ charges of N; smallest integers that balance: 3 Mg2+ = +6 and 2 N3- = -6)
The balanced formula is $\mathrm{Mg3N2}$.
A general approach: use the magnitudes of the charges to determine the subscripts, then reduce if possible (common factors).
Common, but sometimes imperfect, rules encountered in class discussions (to be aware of):
Group 1 alkali metals are typically +1.
Group 2 alkaline earth metals are typically +2 (note: the transcript contains a misstatement here as part of its explanation; the correct charge is +2 for Group 2).
Nonmetals such as nitrogen typically take -3, oxygen typically -2, and halogens typically -1 (the transcript describes some inaccuracies here; the standard charges follow the periodic table and the octet rule).
The above charge assignments are the basis for predicting formulas of binary ionic compounds; always verify with proper oxidation state rules when applying to real problems.
Practice problems (classification and basic naming in ionic context)
Task: Classify each substance as atomic element, molecular element, molecular compound, or ionic compound.
Example discussion points from the transcript:
Calcium oxide is an ionic compound (CaO) with Ca^2+ and O^2-; forms