Difficulty: Intermediate | Prerequisites: Atomic structure, periodic table basics (Doc 2)
This set of notes covers how electrons arrange themselves in atoms (electron configuration), what happens when atoms gain or lose electrons (ions), how atoms bond together (ionic vs. covalent), and the systematic rules for naming the resulting compounds. If you can write an electron configuration, predict an ion's charge, draw a Lewis structure, and name a compound from its formula, you have the core of this material.
Electron configuration
A notation showing how electrons are distributed among an atom's orbitals (e.g. 1s2 2s2 2p6 3s1 for sodium).
In simple terms, this means: it is the address system for every electron in an atom.
Valence electrons
Electrons in the outermost shell of an atom. These are the electrons involved in bonding.
Think of it as: valence electrons are the ones doing all the chemistry. Inner electrons mostly sit there.
Core electrons
Electrons in the inner, filled shells. They shield the valence electrons from the full nuclear charge but are not directly involved in bonding.
Cation
A positively charged ion, formed when an atom loses one or more electrons.
In simple terms, this means: metals typically form cations. Sodium loses one electron to become Na+.
Anion
A negatively charged ion, formed when an atom gains one or more electrons.
In simple terms, this means: nonmetals typically form anions. Chlorine gains one electron to become Cl-.
Isoelectronic species
Different atoms or ions that have the same electron configuration (e.g. Na+, Ne, and F- all have 10 electrons arranged as 1s2 2s2 2p6).
Atomic radius
The size of an atom, measured as half the distance between the nuclei of two bonded atoms of the same element.
Effective nuclear charge (Zeff)
The net positive charge experienced by valence electrons after accounting for the shielding effect of core electrons.
Think of it as: the "pull" the nucleus exerts on the outermost electrons. More protons with the same shielding means a stronger pull.
Ionic bond
A bond formed by the electrostatic attraction between a cation and an anion, typically between a metal and a nonmetal.
Covalent bond
A bond formed when two atoms share one or more pairs of electrons, typically between two nonmetals.
Lewis dot symbol
A representation of an atom showing its valence electrons as dots around the element symbol.
Octet rule
Atoms tend to form bonds until they have eight electrons in their valence shell, achieving the electron configuration of a noble gas.
In simple terms, this means: atoms "want" a full outer shell, and they bond to get one.
Molecular formula
A formula showing the types and numbers of atoms in a molecule (e.g. H2O, C6H12O6).
Structural formula
A formula showing how atoms are bonded to each other, using lines to represent bonds.
Polyatomic ion
A charged group of two or more covalently bonded atoms that acts as a single ion (e.g. SO4 2-, NH4+).
Electrons fill orbitals in order of increasing energy. The filling order is:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p ...
Complete configuration lists every occupied orbital:
Sodium (Na, Z = 11): 1s2 2s2 2p6 3s1
Chlorine (Cl, Z = 17): 1s2 2s2 2p6 3s2 3p5
Abbreviated (noble gas core) configuration replaces the inner electrons with the nearest preceding noble gas in brackets:
Sodium: [Ne] 3s1
Chlorine: [Ne] 3s2 3p5
Valence electrons are in the outermost shell (highest principal energy level). For main-group elements, the number of valence electrons equals the group number (Groups 1 through 8, using the 1-18 numbering: Group 1 = 1, Group 2 = 2, Group 13 = 3, Group 14 = 4, and so on).
Core electrons are everything else. They sit in filled inner shells and shield the valence electrons from the nucleus.
For sodium ([Ne] 3s1): 1 valence electron, 10 core electrons.
When atoms gain or lose electrons, they become ions:
Metals lose electrons to form cations (positive ions). Sodium loses 1 electron to become Na+ (same electron configuration as Ne).
Nonmetals gain electrons to form anions (negative ions). Chlorine gains 1 electron to become Cl- (same electron configuration as Ar).
Ion charges for main-group elements are predictable from the group number:
Group 1: +1 ions
Group 2: +2 ions
Group 16: -2 ions
Group 17: -1 ions
Atomic radius: decreases across a period (left to right), because more protons pull the electron cloud in tighter. Increases down a group, because each new period adds a shell.
Ion size: cations are smaller than their parent atom (fewer electrons, same nuclear charge). Anions are larger (more electrons, same nuclear charge).
Effective nuclear charge (Zeff): increases across a period. More protons with roughly the same shielding means each valence electron feels a stronger pull.
Species with the same number of electrons and the same electron configuration. Na+, F-, Ne, Mg2+, and O2- are all isoelectronic (10 electrons, 1s2 2s2 2p6).
Among isoelectronic species, the one with more protons is smaller (stronger nuclear pull on the same number of electrons).
Some elements exist as isolated atoms in their natural state. The noble gases are the main examples: He, Ne, Ar, Kr, Xe, Rn.
Seven elements naturally exist as two-atom molecules. Memorise them:
H2, N2, O2, F2, Cl2, Br2, I2
A common mnemonic: "Have No Fear Of Ice Cold Beer" (H, N, F, O, I, Cl, Br).
When writing chemical equations involving these elements in their elemental form, you must write them as diatomic molecules, not single atoms. Oxygen gas is O2, not O.
Main-group metals lose electrons to form positive ions; nonmetals gain electrons to form negative ions. The ion charge is predictable from the periodic table position (see the ions section above).
Molecular formula: shows the number of each type of atom (H2O, C6H12O6).
Structural formula: shows how atoms are connected by bonds.
Ball-and-stick and space-filling models: 3D representations that illustrate geometry and relative atomic size.
Name the cation (the metal) first, then the anion (the nonmetal) with an -ide ending.
NaCl = sodium chloride
MgO = magnesium oxide
For transition metals with variable charges, use Roman numerals in parentheses to indicate the charge:
FeCl2 = iron(II) chloride
FeCl3 = iron(III) chloride
For compounds with polyatomic ions, name the ions directly:
Na2SO4 = sodium sulfate
Ca(NO3)2 = calcium nitrate
Use Greek prefixes to indicate the number of each atom:
mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6)
CO2 = carbon dioxide
N2O5 = dinitrogen pentoxide
The first element drops the mono- prefix (carbon monoxide is the exception: CO is carbon monoxide, not monocarbon monoxide)
Binary acids (no oxygen): hydro- + root + -ic acid. HCl = hydrochloric acid.
Oxoacids (contain oxygen): named from the polyatomic ion.
If the ion ends in -ate, the acid ends in -ic. SO4 2- is sulfate, so H2SO4 is sulfuric acid.
If the ion ends in -ite, the acid ends in -ous. SO3 2- is sulfite, so H2SO3 is sulfurous acid.
A Lewis dot symbol places dots around an element's symbol to represent its valence electrons. Carbon (4 valence electrons) gets four dots, one on each side of the C before pairing begins.
Lewis structures for molecules show bonding pairs (shared electrons, drawn as lines) and lone pairs (unshared electrons, drawn as dots).
Atoms tend to form bonds until they have eight electrons in their valence shell. Hydrogen is the exception: it aims for two (a duet).
The number of covalent bonds an atom typically forms can be predicted from its Lewis structure:
Carbon: 4 bonds
Nitrogen: 3 bonds (+ 1 lone pair)
Oxygen: 2 bonds (+ 2 lone pairs)
Halogens: 1 bond (+ 3 lone pairs)
Formed when electrons transfer from a metal to a nonmetal. The resulting cation and anion attract each other electrostatically.
Properties of ionic compounds: high melting points, conduct electricity when molten or dissolved in water, typically form crystalline solids.
Formed when two nonmetals share electrons. The shared pair is attracted to both nuclei simultaneously.
Properties of covalent compounds: varied melting points (generally lower than ionic), poor conductors of electricity, can exist as gases, liquids, or soft solids.
The difference in electronegativity between two bonded atoms determines the bond type:
Large difference (typically > 1.7): ionic
Small difference (typically < 1.7): covalent
Intermediate cases exist (polar covalent bonds), but at the CHEM101 level the key distinction is ionic vs. covalent.
Electron configuration determines every element's chemical behaviour, which is why the periodic table is organised the way it is. Semiconductor design, for example, depends entirely on understanding how silicon's four valence electrons form covalent bonds and how doping with elements from neighbouring groups changes conductivity.
Naming conventions are not academic exercises. Pharmacists, lab technicians, and chemical engineers rely on systematic naming to identify compounds unambiguously. Iron(II) chloride and iron(III) chloride are completely different substances with different properties and uses.
Students often think that electrons orbit the nucleus in neat circular paths, like planets around the sun. They do not. Electron clouds (orbitals) describe probability distributions, not fixed orbits.
"Valence electrons" and "outermost electrons" mean the same thing for main-group elements, but students sometimes confuse them with the highest-energy electrons in transition metals, where the situation is more nuanced.
Students frequently mix up cations and anions. Remember: a cat has paws (cat-ion is paws-itive). Or: cation = "cat" = positive, anion = "a negative ion."
The octet rule is a guideline, not a law. Hydrogen follows a duet rule, and some elements (especially those in Period 3 and below) can have expanded octets. At the CHEM101 level, focus on the exceptions your professor highlights.
⚠️ Writing electron configurations (full and abbreviated) for any element is a core exam skill. Practice until it is automatic.
⚠️ Predicting ion charges from periodic table position appears on every exam. Group 1 = +1, Group 2 = +2, Group 16 = -2, Group 17 = -1.
⚠️ Know the seven diatomic elements. If you write O instead of O2 in a chemical equation, the equation is wrong.
⚠️ Naming compounds (ionic, molecular, and acids) is heavily tested. Be able to go in both directions: formula to name and name to formula.
⚠️ Understand periodic trends (atomic radius, ion size, effective nuclear charge) well enough to compare any two elements or ions.
Fill in the blank: the abbreviated electron configuration for chlorine is [Ne] 3s2 ______. (3p5)
True or false: a cation is larger than its parent atom. (False; it is smaller.)
Fill in the blank: NaCl is named sodium ______. (chloride)
True or false: CO2 is an ionic compound. (False; it is a molecular/covalent compound.)
Name the seven diatomic elements. (H2, N2, O2, F2, Cl2, Br2, I2)
Q: Write the full and abbreviated electron configuration for potassium (K, Z = 19).
A: Full: 1s2 2s2 2p6 3s2 3p6 4s1. Abbreviated: [Ar] 4s1.
Q: Predict the charge of a magnesium ion and explain why.
A: Mg forms Mg2+. Magnesium is in Group 2, so it loses its two valence electrons to achieve the noble gas configuration of neon.
Q: Name the compound Fe2O3.
A: Iron(III) oxide. Each iron must be +3 to balance the three oxide ions (each -2): 2(+3) + 3(-2) = 0.
Q: What type of bond forms between sodium and chlorine? Between two oxygen atoms?
A: Sodium and chlorine form an ionic bond (metal + nonmetal, large electronegativity difference). Two oxygen atoms form a covalent bond (nonmetal + nonmetal, electrons are shared).
Q: Rank Na+, Mg2+, F-, and Ne in order of increasing size.
A: Mg2+ < Na+ < Ne < F-. All are isoelectronic (10 electrons). More protons means a stronger pull on the same electron cloud, so the ion is smaller. Mg2+ has 12 protons, Na+ has 11, Ne has 10, F- has 9.
Electron configuration is the foundation for understanding chemical bonding, molecular geometry (VSEPR theory, typically covered next in CHEM101), and intermolecular forces.
The naming rules introduced here apply throughout the rest of the course, including stoichiometry, solution chemistry, and acid-base reactions. If you can name a compound and write its formula, you can balance an equation involving it.
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