Metals and Nonmetals: Bonding Fundamentals
Metals and Nonmetals: Focus and Overview
Goal: build a foundation to understand everyday materials and their bonding behaviors.
Everyday Uses and Examples of Metals
Everyday objects that rely on metals include:
Spoon (hardness and durability)
Jewelry (luster, malleability, ductility)
Kettle (heat conduction, durability)
Coins and door handles (conductivity, durability, metal content: copper, iron, etc.)
Water bottles (metal variants)
Phones and charging devices (conductivity, housings)
Keys and watches (durability, aesthetics)
Metals appear in a wide variety of applications due to their combination of properties.
Key Properties of Metals
Hardness: ability to withstand deformation; important for durability of objects like spoons, doorknobs, etc.
Conductivity: metals conduct electricity and heat; enables wires in charging devices, kettles, etc.
Malleability: can be hammered into thin sheets (e.g., aluminum foil) and molded.
Ductility: can be drawn into wires (e.g., electrical cables).
Luster: shiny appearance (e.g., faucets, doorknobs).
Metals generally have solid, dense structures and exhibit a combination of these properties, making them versatile for everyday use.
Periodic Table Overview: Metals, Nonmetals, and Metalloids
The periodic table is organized by atomic number; elements colored yellow are metals, blue are nonmetals, and green elements are metalloids (in-between metals and nonmetals).
There is a white staircase line on the periodic table used to separate metals from nonmetals; elements just above/below this staircase tend to be metal, while elements on the other side tend to be nonmetal.
Hydrogen is a special case: it is a nonmetal but sits on the left side of the table; it’s the only nonmetal in that left region.
The colorized version of the periodic table is not shown here, but the staircase line helps identify metals vs nonmetals.
The lecture emphasizes understanding the metals and nonmetals through bonding, not through memorizing color schemes alone.
Valence Electrons and Electron Shells: Basic Concepts
Atoms have energy levels (shells); outermost shell is the valence shell.
Valence electrons are the electrons in the outermost shell and are the ones involved in bonding.
Core electrons are all electrons in inner shells (not outermost).
Lithium (Li) example:
Li has 3 total electrons: distribution is 2 in shell 1, and 1 in shell 2.
Valence electrons in Li: 1 (outermost shell).
Core electrons in Li: 2 (in shells 1 and any inner shells).
Question to consider: How does valence electron count relate to chemical reactivity? Valence electrons are the ones most likely to be lost or gained during bonding; core electrons are held more tightly by the nucleus.
Pattern of valence electrons across a period (left to right): the valence electron count increases by 1 as you move across the row.
Pattern down a column (same group): elements have the same number of valence electrons.
Exceptions and special notes:
Helium (He) is in column 18 but has only 2 electrons total, so it cannot have 8 valence electrons. It is an exception to the typical "8 valence electrons" rule for most elements.
The table also notes special group names: Column 1 = alkaline metals, Column 2 = alkaline earth metals, Column 17 = halogens, Column 18 = noble gases.
Hydrogen is not an alkaline metal; it is a nonmetal and sits on the left side of the table.
Transition metals (in the middle, columns 3–12) do not follow the simple valence electron trends described here; they are not the focus for these basic trends.
Quick rule summaries from the lecture:
Elements in the same period: valence electrons increase by 1 left to right.
Elements in the same column: same number of valence electrons.
Helium: exception to the eight-electron valence rule; only 2 electrons total.
Hydrogen: nonmetal on the left, unique placement among the left-side elements.
Quick exercise patterns (as described in class):
If asked to identify elements with four valence electrons by scanning the periodic table, candidates are found in column 14 (excluding transition metals in columns 3–12).
In a larger sense, use the periodic table to infer valence electron counts for non-transition metals by their column position.
The Octet Rule and Bonding Motivation
Octet rule: most elements strive to have a full outer (valence) shell with eight electrons. For helium, the rule is different: its valence shell can hold only two electrons.
Motivation for bonding: atoms bond to achieve a full outer shell (stable electronic configuration).
Sodium (Na, 11 protons) and chlorine (Cl, 17 protons) example:
Sodium has one valence electron in its outer shell, chlorine has seven valence electrons in its outer shell.
It is easier for sodium to lose one electron than for sodium to gain seven; it is easier for chlorine to gain one electron than to gain seven.
Result: sodium donates an electron to chlorine, forming an ionic bond via electron transfer.
Ionic bond overview:
An ionic bond is the electrostatic attraction between opposite charges created by transfer of electrons.
It occurs typically between metals (which donate electrons) and nonmetals (which accept electrons).
Sodium (Na) becomes a cation (Na⁺) after losing an electron; chlorine becomes an anion (Cl⁻) after gaining an electron.
Notation and terminology:
Cation: positively charged ion formed when an atom loses electrons. Example: Na → Na⁺ + e⁻.
Anion: negatively charged ion formed when an atom gains electrons. Example: Cl + e⁻ → Cl⁻.
Ionic bond: attraction between Na⁺ and Cl⁻ in sodium chloride, NaCl.
In simple ionic bonding, the metal is the donor (cation) and the nonmetal is the receiver (anion).
Transition metals can participate in ionic bonds as well, but these lectures focus on alkali/alkaline metals in the context of simple ionic bonds.
Examples of ionic bonds and charges (conceptual):
NaCl (sodium chloride) forms from Na⁺ and Cl⁻ via electron transfer.
The charge on the metal side corresponds to the number of electrons donated; the charge on the nonmetal side corresponds to the number of electrons gained.
The bond is an electrostatic attraction between the ions, not a sharing of electrons as in covalent bonds.
Crystals, dissolution, and what happens when salt dissolves:
In a crystal of salt, each ionic bond ties together Na⁺ and Cl⁻ in a lattice.
If you break the salt crystal into smaller pieces, the individual NaCl units (ions) remain bound in the lattice within each piece.
When salt is dissolved in water, the ionic bonds break, and the ions separate and disperse in the solution (ions are solvated by water molecules).
Additional ionic bonding example: calcium sulfide and multiple-electron transfer
Calcium and sulfur example: Ca has two valence electrons to donate; sulfur needs two electrons to complete its octet.
Transfer: Ca donates two electrons to S, resulting in Ca²⁺ and S²⁻.
The resulting compound: CaS (calcium sulfide).
Another example with a multielectron transfer: lithium oxide (Li₂O)
Oxygen needs two electrons to achieve its octet; each lithium can donate one electron.
Reaction steps:
Two Li atoms donate one electron each: 2 ext{Li}
ightarrow 2 ext{Li}^{+} + 2e^{-}Oxygen accepts two electrons: ext{O} + 2e^{-}
ightarrow ext{O}^{2-}Combined product: 2 ext{Li}^{+} + ext{O}^{2-}
ightarrow ext{Li}_{2} ext{O}
Naming of ionic compounds (basic convention):
The compound name reflects the ions involved, but you do not usually specify the exact number of each ion in the name.
Examples:
Sodium chloride → NaCl
Calcium sulfide → CaS
Lithium fluoride → LiF
Lithium oxide → Li₂O
Formulas convey the number of ions needed to balance charges, but the name focuses on the identities of the ions (not the numeral counts).
Summary of ionic bonding rules from the lecture:
A metal (alkali or alkaline earth metal) tends to lose electrons → becomes a cation.
A nonmetal (in the appropriate columns, e.g., halogens) tends to gain electrons → becomes an anion.
The combination aims to achieve full octets for both species.
The resulting compound forms a crystal lattice with ionic bonds; dissolution in water separates into ions.
Naming conventions reflect the ions involved, not the exact counts of each ion in every molecule.
Quick conceptual recap (connection to prior topics and real-world relevance):
The tendency of metals to lose electrons and nonmetals to gain electrons underpins many materials and salts used in everyday life (table salt, battery materials, corrosion processes).
Understanding valence electrons helps explain why bonding occurs and why certain compounds form with specific ratios.
The octet rule provides a simplified framework for predicting bonding; real chemistry includes exceptions and more complex bonding in transition metals and polyatomic ions, which are explored in more advanced topics.
Note for study strategy (as discussed in class):
Use the periodic table to infer valence electron counts by group (with caution for transition metals).
Remember the common ion names and charges for alkali/alkaline earth metals and halogens when predicting ionic bonds.
Practice with examples like NaCl, LiF, CaS, and Li₂O to reinforce electron transfer concepts and naming conventions.