Periodic Table, Ionic/Molecular Compounds, and Nomenclature - Study Notes

Periodic Table Basics: structure and conventions

Periodic table organization
- Rows are called periods; each row corresponds to one period.
- Columns are called groups; each column corresponds to one group.
- Elements are ordered by increasing Z (protons).
Group naming and numbering conventions
- Two common systems:
- North American system: groups 1–18 (most widely used in teaching).
- Older IUPAC/European style: use 1A–8A and 1B–8B (letters a/b).
- In this course, the North American 1–18 system is used for consistency.
- Important note: in your book you may still see the older 1A–8A, etc.
Main-group vs. transition metals (classification within groups)
- Main-group elements: often denoted as groups 1–2 and 13–18 (A designations in some texts).
- Transition metals: denoted as groups 3–12 (often with a B designation).
- The main-group portion includes groups labeled from 1 to 8 (A), while transition metals are the B groups (3–8, in the lecture’s framing).
Periodic table layout and element types
- Metals: typically have metallic luster, are malleable, and good conductors of electricity and heat.
- Nonmetals: do not have metallic properties.
- Metalloids (semi-metals): have intermediate properties between metals and nonmetals; shown on the “staircase” line in many diagrams (often colored green).
- Color cues in the slide example: blue = metals, yellow = nonmetals, green staircase = metalloids.
General trends and positioning
- Most elements on the periodic table are metals.
- Most nonmetals are found toward the upper-right portion of the table; metals are extensive and can be found throughout, especially on the left and center.
- Important practical implication: metal vs. nonmetal behavior helps predict bonding and compound type (ionic vs covalent).
Element names and groups to know from a nomenclature and bonding perspective
- Group 1A and Group 2A elements: form common monatomic cations with charges equal to their group number (as a general rule for main-group elements).
- Example charges for main-group cations:
- Group 1A: +1 cation
- Group 2A: +2 cation
- Group 3A: +3 cation (e.g., Al^{3+})
- Halogens (Group 7A) and chalcogens (Group 6A) are particularly important for anions and coordination in compounds.
Period lengths to recall (overview)
- Period 1: 2 elements
- Period 2: 8 elements
- Period 3: 8 elements
- Period 4 and beyond: 18 elements (and later periods can have 32 in the modern table)

Chemical formulas: molecules, ions, and formula units

What a chemical formula represents
- Molecular formulas: show the exact number of atoms of each element in a molecule (e.g.,
- H$_2$O tells you two hydrogens and one oxygen per molecule).
Ionic substances vs molecular substances
- Ionic compounds: do not form discrete molecules; they are extended lattices of ions.
- For ionic compounds, the appropriate descriptor is a formula unit, which reflects the simplest ratio of ions that yields electrical neutrality.
- Molecular compounds (covalent): described by molecular formulas and often have structural formulas showing connectivity.
Charge balance and the idea of neutrality
- Metals tend to lose electrons to form cations; nonmetals tend to gain electrons to form anions.
- The overall compound must be electrically neutral.
- Example balancing concept: to balance Mg^{2+} with N^{3-}, the smallest whole-number ratio comes from the least common multiple of 2 and 3, which is 6. Use 3 Mg^{2+} and 2 N^{3-} to yield neutral Mg$3$N$2$.
- Notation rule demonstrated in the class: write the subscripts to reflect the correct ratio; charges are not written in the chemical formula itself.
Polyatomic ions and the use of parentheses
- Some ions are polyatomic (consisting of more than one atom) but carry a single overall charge (e.g.,
- Ammonium, NH$_4^+$, is a polyatomic cation.
- Mercury(I) ion, Hg$_2^{2+}$, is also polyatomic (two mercury atoms sharing a 2+ overall charge).
- When a polyatomic ion is part of a formula unit, parentheses are used so the subscript applies to the entire polyatomic ion (e.g., Ca$3$(PO$4$)$_2$).
Crystalline vs molecular: key takeaway
- Ionic compounds are typically crystalline solids and are described by a formula unit.
- Molecular compounds are composed of discrete molecules and described by molecular/structural formulae.

Nomenclature of ionic inorganic compounds

General rule for binary ionic compounds (monatomic ions)
- Cation name (the metal) + anion name (nonmetal stem + -ide).
- Example: NaCl is sodium chloride.
Role of charge and the stock system
- Many main-group metals form only one cation (e.g., Group 1 and Group 2 metals): no need to specify charge.
- Transition metals and some post-transition metals form cations with multiple possible charges; their charges are indicated with a Roman numeral in parentheses after the metal name (stock system).
- Example: FeCl$2$ vs FeCl$3$ correspond to iron(II) chloride and iron(III) chloride, respectively.
Older nomenclature with -ous/-ic endings
- The lower oxidation state uses -ous; the higher oxidation state uses -ic (e.g., Fe^{2+} = ferrous, Fe^{3+} = ferric).
- Copper can have +1 or +2; Cu^{+} is cuprous, Cu^{2+} is cupric.
- Mercury(I) and other ions can have older names as well (e.g., mercurous for Hg$_2^{2+}$, mercuric for Hg$^{2+}$).
- The stock (Roman numeral) system is the standard method in this course; the old system can be used historically but is not required.
Polyatomic ions: key examples to memorize
- Ammonium: NH$_4^+$ (polyatomic cation).
- Cyanide: CN$^-$ (monatomic anion; cyanide is one of the important polyatomic anions in teaching materials).
- Hydroxide: OH$^-$; Peroxide: O$2^{2-}$; Nitrate: NO$3^{-}$; Nitrite: NO$_2^{-}$.
- Carbonate: CO$_3^{2-}$;
- Hydrogen carbonate (bicarbonate): HCO$_3^{-}$;
- Dihydrogen phosphate: H$2$PO$4^{-}$;
- Monohydrogen phosphate: HPO$_4^{2-}$;
- Phosphate: PO$_4^{3-}$.
- Sulfate: SO$4^{2-}$; Hydrogen sulfate (bisulfate): HSO$4^{-}$.
Practical example problems (from the lecture)
- Example: Write the formula for the ionic compound formed by Mg^{2+} and N^{3-}.
- Charges: +2 and −3. Use the least common multiple 6; multiply Mg by 3 and N by 2.
- Formula: ext{Mg}3 ext{N}2
- Example: Write the formula for calcium cation with phosphate anion.
- Ca^{2+} and PO$4^{3-}$ balance with 3 Ca^{2+} and 2 PO$4^{3-}$.
- Formula: ext{Ca}3( ext{PO}4)_2
Notes on polyatomic cations and anions in practice
- Polyatomic cations include ammonium (NH$4^+$) and Hg$2^{2+}$ (the diatomic cation from mercury, historically called mercurous/mercury(I)).
- Many polyatomic anions require parentheses to apply the overall charge to the entire ion when forming compounds (e.g., Ca$3$(PO$4$)$_2$).
Special case: determining ion charges from the periodic table (quick cheat sheet)
- Main-group metals (typical non-transition metals) tend to form monatomic cations; their charge often equals their group number for groups 1A–3A (with some exceptions in group 13).
- Nonmetals form anions; their charge is given by the group number minus 8:
- Example: Cl is in Group 17 → charge = 17 − 8 = −1; O is in Group 16 → charge = 16 − 8 = −2; N is in Group 15 → charge = 15 − 8 = −3.
- Transition metals can form multiple cations (e.g., Fe^{2+}, Fe^{3+}; Cu^{+}, Cu^{2+}); charge is indicated via Roman numerals in the stock system.
Summary of naming rules for ionic compounds
- If the compound contains only one possible cation charge (most main-group metals), just use the element name for the cation and the -ide ending for the anion.
- If there are multiple possible charges for the metal, specify the charge with a Roman numeral after the metal name (stock system).
- For polyatomic ions, use the ion’s standard name; put parentheses around a polyatomic ion when a multiplier applies to all atoms within it (as in Ca$3$(PO$4$)$_2$).

Organic vs. Inorganic compounds; functional groups and reactivity

-Organic compounds (by definition in this course)

Contain carbon (and usually hydrogen); often include other elements like O, N, S.
hydrocarbons: compounds containing only carbon and hydrogen.
Approximately 60% of known compounds are organic; organic chemistry is central to biology, medicines, and materials.
Historical note: organics were once thought to come only from living organisms, but lab synthesis (e.g., urea from urine) demonstrated otherwise.
Reactivity is largely governed by functional groups (e.g., -OH, -COOH, -C=O, etc.).
- Inorganic compounds (three important carbon-containing exceptions in the lecture)
Not all carbon-containing compounds are organic; some carbon-containing inorganic compounds include carbon monoxide (CO) and carbon dioxide (CO$_2$), as well as carbonates and cyanides.
The three exceptions frequently highlighted: carbonate (CO$3^{2-}$), cyanide (CN$^{-}$), and carbon dioxide (CO$2$).
The three exceptions are part of the inorganic chemistry portion of the course despite containing carbon.
- Nomenclature focus for inorganic compounds (inorganic naming rules, molecular and ionic)
Molecular inorganic compounds: naming follows a fixed set of rules (tables in the textbook/lecture) rather than memorizing every compound name.
Ionic inorganic compounds: balance of charges between cations and anions; rules for naming are introduced as part of the lecture.
- Functional groups and reactivity (brief)
Functional groups are specific groupings of atoms within molecules that determine reactivity patterns and properties.
Examples mentioned: hydroxyl group (-OH), carboxyl group (-COOH), alcohols, etc.

Molecular structure representations and polymers

Structural representations of molecules
- Structural formulas show not only which atoms are present, but how they are bonded and arranged in space.
- Examples: H$_2$O (water) is not linear in 3D space; it is bent with a bond angle around 104.5°.
- NH$_3$ (ammonia) has a pyramidal geometry.
- Ethanol: a more extended structure with both C–C and C–O bonds.
Ball-and-stick models vs space-filling models
- Ball-and-stick: atoms are balls connected by sticks representing bonds; good for visualizing connectivity and angles.
- Space-filling models: emphasize the relative sizes and spatial occupation of atoms; useful for a sense of molecular surface but can be crowded for large molecules.
Polymers and repeating units
- Polymers are large molecules built from repeating subunits called monomers.
- Examples: nylon, Kevlar, wool (natural polymer), DNA (biopolymer).
- A polymer chain is essentially a long sequence of monomers joined together.
- The concept of a repeating unit explains why polymers can have vast diversity despite a relatively small set of monomer types.

Quick recap of key ideas to memorize

Periodic table basics: rows = periods; columns = groups; 1–18 vs A/B notation; main-group vs transition metals; metal vs nonmetal vs metalloid location.
Ionic vs covalent (molecular) compounds: ionic compounds are crystalline, form formula units, and are neutral overall; molecular compounds are discrete molecules.
Rules of ionic compound nomenclature: cation name + anion name (stem + -ide); Roman numerals for metals with multiple possible charges; polyatomic ions require parentheses when needed.
Common polyatomic ions and their charges (quick recall):
- Ammonium: NH$_4^+$
- Hydroxide: OH$^-$
- Cyanide: CN$^-$
- Carbonate: CO$3^{2-}$; Hydrogen carbonate (bicarbonate): HCO$3^{-}$; Dihydrogen phosphate: H$2$PO$4^{-}$; Monohydrogen phosphate: HPO$4^{2-}$; Phosphate: PO$4^{3-}$
- Sulfate: SO$4^{2-}$; Hydrogen sulfate (bisulfate): HSO$4^{-}$
Balancing charges to form neutral ionic compounds (illustrative):
- Magnesium nitride:
- Mg^{2+} and N^{3-} balance to Mg$3$N$2$ (3 × +2 = +6; 2 × −3 = −6).
- Formula: ext{Mg}3 ext{N}2
- Calcium phosphate: Ca$3$(PO$4$)$_2$ balances 3 × (+2) and 2 × (−3) to neutrality.
- Formula: ext{Ca}3( ext{PO}4)_2
The role of oxidation state designations in naming (stock system) and the old ic/ous conventions (historical, not required, but good to know):
- Iron(II) chloride: FeCl$2$; Iron(III) chloride: FeCl$3$.
- Copper(I) vs Copper(II): Cu$^+$ (cuprous) and Cu$^{2+}$ (cupric).
- Mercury(I) can be represented as Hg$_2^{2+}$ (polyatomic cation; historically mercurous).
Organic vs inorganic: carbon-containing organics vs carbon-containing inorganics; functional groups govern reactivity; examples and exceptions to carbon-containing inorganic compounds (CO, CO$2$, CO$3^{2-}$, CN$^{-}$, HCO$3^{-}$, H$2$PO$_4^{-}$, etc.).
Practical skills to develop
- Be able to identify whether a compound is ionic or molecular from its components (metal + nonmetal usually ionic; nonmetal + nonmetal usually covalent).
- Be able to write and name simple ionic compounds and to recognize when a Roman numeral is needed to denote the cation charge.
- Recognize and name common polyatomic ions and understand when parentheses are needed in formulas.
- Understand the basic difference between formulas (empirical/molecular) and formula units for ionic compounds.
Encouragement for study habits mentioned in class
- Read the relevant sections ahead of time; actively listen, write, and think during the lecture; practice with example problems on pages referenced (e.g., Example problems around writing and balancing ionic compound formulas).