Chemistry Notes: Periodic Table Trends, Group Properties, Bonding and Nomenclature
Study Environment and Exam Prep
Instructor-student support: the instructor and a peer tutor are available during office hours to answer questions, provide study help, or offer another perspective on problem-solving.
Extra study time: use the optional hour per week to practice chemistry or discuss approaches to problems; there is more than one way to solve these questions.
Exam logistics: the course uses a lockdown/web browser for exams; you’ll need to use a browser to access the exam securely; there is mention of an alternate workflow in-browser (the term Axe on the browser) which can cause confusion.
Purpose of the session: build understanding of content and problem-solving approaches beyond a single method.
Periodic Table Structure and Core Concepts
Periodic table layout:
Periods: horizontal rows on the periodic table.
Groups: vertical columns on the periodic table.
There is a progression in properties from left to right across a period and a change in properties down a group.
Metals vs nonmetals vs metalloids:
Metals (left/molded area): good conductors of heat and electricity; typically ductile; mostly solids.
Nonmetals (right side): insulators; not good conductors of heat or electricity; can be gases at room temperature; brittle; not ductile.
Metalloids: intermediate properties (between metals and nonmetals).
General organization (as discussed): left to right you see a gradual change from metals to metalloids to nonmetals.
Notable groups and regions (names and general properties):
Group 1: Alkali metals (alkali metals are highly reactive; form salts with halogens; e.g., Na forms NaCl with Cl).
Group 2: Alkaline earth metals (more stable than Group 1 but still reactive; form salt-like compounds).
Group 17: Halogens (salt-forming elements; form one bond typically with alkali metals to give salts; e.g., NaCl; they form salts in large crystalline lattices and are highly reactive).
Group 18: Noble gases (chemically inert/inactive; little tendency to form bonds).
Inner transition metals: lie in the middle region; rare or radioactive; limited interaction in introductory courses; include the lanthanides and actinides.
Transition metals: middle block; not all share the same properties, so they’re grouped rather than named by simple column numbers.
Practical takeaway: the periodic table is designed so that similar chemical properties cluster together (akin to Mendel’s grouping idea, but organized by periodicity).
Atomic Size (Atomic Radius) Trends
Across a period (left to right): atomic radius generally decreases.
Represented as: as you move across a period, the size “gets smaller.”
Described numerically as a trend in radii across the period.
Down a group (top to bottom): atomic radius generally increases.
Observational notes:
For the first three periods, the trend is clear and decreases across the period.
In the 4th and 5th periods, the trend is less visually clear due to more elements and cross-period variations; trend deficiencies exist and are typically discussed in more advanced courses.
Quick rules of thumb for comparing sizes:
In the same period, the one farther to the left is smaller.
In the same group, the one closer to the top is smaller.
Practice question insights:
The smallest element is found toward the top-right of the table; the largest is toward the bottom-left.
Example exercise (five elements): Be, Mg, Ne, K, Rb
The smallest element among these is Be (beryllium) because it is high up and toward the left in its period.
The ordering of size across these elements follows the general across-down trend discussed above.
Periodic Table Segments and Group Names (Context and Practice)
Periods: 1 through 7 (horizontal rows).
Groups: vertical columns used as a coordinating system to identify element similarities.
Common named groups and rough content (from the lecture):
Group 1: Alkaline metals (alkali metals) — highly reactive, especially with water.
Group 2: Alkaline earth metals — more stable than Group 1 but still reactive.
Group 16 (interpreted from the lecture as “groups six”) and Group 17 — show distinct bonding patterns; halogens (Group 17) form one bond with metals in simple binary salts.
Group 17 (Halogens): form salts broadly; halogen + alkali metal gives a salt (e.g.,
).Group 18: Noble gases — chemically inactive; generally do not form bonds.
Bottom-line concept: three notable trends drive chemical behavior: metallic character, bonding tendencies, and reactivity patterns across and down the table.
Inner transition metals: rare, often radioactive; less emphasis in introductory chapters; not a heavy focus in early exams; positioned between main blocks.
Practice reminder: be able to locate period numbers (1–7) and group numbers on the periodic table to answer basic questions about element identity and placement.
Chemical Formulas, Bonding, and Reactions
Basic idea: chemical formulas represent the types and numbers of atoms in a compound.
Example bonding patterns:
A simple metal + nonmetal compound (ionic): one metal atom combines with a nonmetal to form a compound with overall charge balance (e.g.,
, sodium chloride).Example provided: if you replace any element in the same group in a formula, the overall formula can remain the same for counting purposes, though the properties will change (e.g.,
vs.
are both salts with similar bonding patterns but different identities and properties).
Specific bonding chemistry concepts given:
Group 1 elements (alkali metals) tend to make one bond in simple salts (e.g., Na forms NaCl with Cl).
Group 2 elements (alkaline earth metals) tend to form salts with the formula
(e.g.,
).Group 16 elements (e.g., oxygen, sulfur) tend to form two bonds in typical compounds (e.g.,
in oxides or
in sulfides).
Specific example discussion:
Sodium reacts with chlorine gas to form crystalline sodium chloride, a stable ionic lattice:
The crystalline form of NaCl can be visually described as a large cubic lattice.
Na metal reacting with water is a classic demonstration of chemical reactivity (a chemical property) while the melting point of sodium is a physical property:
Sodium metal (Na) has a melting point of , a physical property, but it reacts violently with water, a chemical property.
Key lesson: when atoms combine, new chemical and physical properties emerge that differ from the original elements; compounds contain two or more different elements.
Notion of salts: halogens form salts with metals; salts are typically large crystalline substances with ionic bonding.
Nomenclature (IUPAC) for Naming Compounds
Core rule for metal + nonmetal compounds (Rule 1):
Name the metal first (the cation).
Name the nonmetal second, changing its ending to -ide.
Hydrogen is ignored in this naming rule when forming the base compound from a metal and a nonmetal.
Rough heuristic to apply naming: metals are found on the left side of the periodic table, nonmetals on the right side; this helps decide if you’re dealing with a metal+nonmetal compound.
Special cases and scope:
The lecture emphasizes memorizing the first two groups (alkali metals and alkaline earth metals) and the first 20 elements (to cover the nonmetals commonly encountered in early chemistry), so you can name typical metal + nonmetal compounds.
For transition metals, the naming scheme can differ and can involve other conventions not deeply covered in this lecture.
Example spellings and common pitfalls:
The correct spelling is fluoride (not flower I) when naming fluoride-containing compounds.
The second element’s name should end in -ide (e.g., chlorine → chloride, oxygen → oxide).
Prefix use with two nonmetals (binary covalent compounds):
Prefixes such as di- and tri- are used to indicate the number of atoms when two nonmetals are involved (e.g., NO is nitrogen monoxide; NO₂ is nitrogen dioxide; CO is carbon monoxide; CO₂ is carbon dioxide).
The prefixes are generally used when two nonmetals form the compound; they do not apply to metal-nonmetal ionic compounds (which follow Rule 1).
Practical naming practice:
Given a compound consisting of a metal and a nonmetal, you should be able to name it by identifying the metal first and the nonmetal second with -ide ending (e.g., NaCl → sodium chloride).
For two nonmetals together, apply di-, tri- as needed (e.g.,
).
Teaching tips for naming practice:
Begin with the element symbols and names to build familiarity with those basics.
Then practice with compounds formed by metals in Groups 1 and 2 and nonmetals on the right side of the table.
Use the left-right line on the periodic table as a quick mental cue for metal vs nonmetal status when naming.
Memorization and Exam Preparation
Key memorization targets:
The first 20 elements and their chemical symbols/names:
Hydrogen (H), Helium (He), Lithium (Li), Beryllium (Be), Boron (B), Carbon (C), Nitrogen (N), Oxygen (O), Fluorine (F), Neon (Ne), Sodium (Na), Magnesium (Mg), Aluminium (Al), Silicon (Si), Phosphorus (P), Sulfur (S), Chlorine (Cl), Argon (Ar), Potassium (K), Calcium (Ca).
The first two groups: Group 1 (alkali metals) and Group 2 (alkaline earth metals).
The nonmetals section excluding noble gases (i.e., the common nonmetals highlighted in the session).
exam-oriented approach:
You should be able to identify the period number for a given element (1–7) and the group number for that element (when provided) to infer properties and typical behavior.
You should be able to determine the relative size of elements using the period/group position and explain why Be is smaller than Mg, why Ne is smaller than Na within their respective contexts, and so on.
You should be able to name simple binary metal-nonmetal compounds using Rule 1 (e.g., NaCl as sodium chloride) and use prefixes di- or tri- for two-nonmetal compounds (e.g., NO₂ as nitrogen dioxide).
Additional practice resources (mentioned in the session):
A short five-minute video linked in the slides demonstrates some reactivity of first-row metals with water; it’s optional but illustrative of reactivity trends (for example, sodium reacting violently with water) and is useful for real-world understanding.
Real-World Relevance and Scientific Context
Reactivity and safety: alkali metals are highly reactive; some reactions are dramatic (as described with water interactions) and are useful for understanding trends in reactivity and safety considerations in handling reactive metals.
Inert environments: noble gases like argon are used to create inert atmospheres to prevent unwanted reactions (e.g., keeping reactive sodium from contacting air).
Materials science: salts formed from metals and halogens crystallize into salt-like lattices, illustrating ionic bonding and crystal structure concepts.
Foundational principles: the periodic table’s structure and the trends discussed underpin predictive chemistry, including bonding types, formula formation, and naming conventions.
Quick Reference: Key Formulas and Examples
Salts formed from metal + halogen:
A basic salt formed from metal + nonmetal with balanced formula:
from Na + Cl₂ → 2 NaCl (illustrative stoichiometry in balanced reactions)Binary compound example with two nonmetals (prefix use):
Atomic radius trend indicators (qualitative):
Across a period:
Down a group: