Difficulty: Introductory to Intermediate | Prerequisites: Part 1 notes (periodic table structure, metals/nonmetals/metalloids).
This set of notes covers the trends that emerge when you move across a period or down a group on the periodic table: effective nuclear charge, atomic and ionic radii, ionization energy, electron affinity and electronegativity. These trends are the reason the periodic table is useful rather than just a catalogue. Once you can predict how a property changes with position, you can make educated guesses about an element's behaviour without memorising every data point. You should already understand what periods, groups and valence electrons are before tackling this material.
Five major properties trend predictably across the periodic table. Effective nuclear charge, ionization energy, electron affinity and electronegativity all increase from left to right across a period (and generally from bottom to top in a group). Atomic radius does the opposite, increasing down a group and to the left across a period. Understanding these trends lets you predict how any element will behave in bonding and reactions.
Effective nuclear charge (Z_eff)
The net positive charge experienced by a valence electron after accounting for the shielding effect of inner electrons. Think of it as how strongly the nucleus "pulls" on the outermost electrons once you subtract the screening from the electrons in between.
Principal quantum number (n)
The number that identifies the main energy level (shell) an electron occupies. As n increases, the electron is farther from the nucleus. In simple terms, it tells you which floor of the atom the electron lives on.
Atomic radius
One half of the distance between the nuclei of two identical atoms that are bonded or in contact. Think of it as the "size" of an atom.
Ionic radius
The effective radius of an ion in a crystal lattice. Cations (positive ions) are smaller than their parent atom because they have lost electrons. Anions (negative ions) are larger because they have gained electrons.
Ionization energy (IE)
The energy required to remove an electron from a gaseous atom or ion. It is an endothermic process (delta H is positive). Think of it as the cost of pulling an electron free.
Electron affinity (EA)
The energy released when a gaseous atom gains an electron. It is an exothermic process (delta H is negative). Think of it as how badly an atom wants an extra electron.
Electronegativity (EN)
A measure of how strongly an atom attracts the electrons in a chemical bond toward itself. Fluorine has the highest electronegativity of any element. In simple terms, it describes how "greedy" an atom is for shared electrons.
Z_eff measures the electrostatic attraction between the nucleus and the valence electrons
Across a period (left to right): Z_eff increases because protons are added to the nucleus while electrons are added to the same shell, so shielding does not increase proportionally
Down a group: Z_eff stays roughly constant for valence electrons because each new shell adds both protons and a full set of inner shielding electrons
The principal quantum number (n) increases by one each time you move down a group, meaning valence electrons are in a higher, more distant shell
Higher n means the valence electrons are farther from the nucleus and held less tightly, even though Z_eff may be similar
Atomic radius trends
Increases moving down a group (higher n, electrons in more distant shells)
Increases moving to the left across a period (fewer protons pulling on the same shell, lower Z_eff)
The largest atom in any period is in Group IA (alkali metals)
The largest atom in any group is at the bottom of the table
Atomic radius is inversely proportional to effective nuclear charge: higher Z_eff pulls the electron cloud in tighter
Ionic radius trends
Cations (metals losing electrons): ionic radius is smaller than the neutral atom. The atom loses its outermost electrons, leaving a smaller ion with fewer electron-electron repulsions.
Metal ionic radius becomes smaller as you move toward the metalloid staircase (higher charge, more electrons lost)
Anions (nonmetals gaining electrons): ionic radius is larger than the neutral atom. Extra electrons increase electron-electron repulsion and expand the cloud.
Nonmetal ionic radius becomes larger as you move toward the metalloid staircase
General rule: cations < neutral atoms < anions (for the same element or isoelectronic series)
The energy required to remove an electron from a gaseous atom or ion
Endothermic process: requires heat input, delta H is positive
Across a period (left to right): IE increases because Z_eff increases and atoms are smaller, so electrons are held more tightly
Down a group: IE decreases because valence electrons are farther from the nucleus (higher n) and easier to remove
Greater Z_eff = more tightly bound electrons = higher IE
First vs. second ionization energy
First ionization energy (IE₁): energy to remove the first electron from a neutral atom
Second ionization energy (IE₂): energy to remove an electron from the resulting +1 cation to form a +2 cation
IE₂ is always greater than IE₁ because you are removing an electron from a species that already has a net positive charge
Special cases
Active metals (e.g. Li, Be, Na): very low ionization energies, always found in ionic compounds rather than in elemental form
Noble gases: highest ionization energies in their period because their electron configuration is already full and stable
The energy released (dissipated) when a gaseous atom gains an electron
Exothermic process: delta H is negative (energy comes out)
Across a period (left to right): EA generally increases (atoms want electrons more)
Down a group (bottom to top): EA generally increases (smaller atoms hold the extra electron more tightly)
Halogens have the highest electron affinities because they need just one more electron to complete their octet
Noble gases have effectively zero electron affinity: their octet is already full, so forcing an extra electron in would require energy rather than release it
Metals have low electron affinity: they prefer to lose electrons to reach a full or half-full shell rather than gain them
The measure of how strongly an atom attracts shared electrons in a chemical bond
Directly proportional to Z_eff: higher effective nuclear charge means a stronger pull on bonding electrons
Across a period (left to right): EN increases
Down a group (bottom to top): EN increases
Fluorine is the most electronegative element (EN approximately 4.0)
Of the noble gases, only krypton and xenon have measurable electronegativity values (because they can form certain compounds)
Unlike ionization energy and electron affinity, electronegativity applies only to atoms in bonds, not isolated atoms
Common electronegativity values worth memorising
H approximately 2.0, C approximately 2.5, N approximately 3.0, O approximately 3.5, F approximately 4.0. These round figures are close enough for quick bond-polarity predictions and come up frequently on exams.
Ionization energy is why sodium street lamps produce that distinctive yellow glow: the energy needed to excite and de-excite sodium's outer electron corresponds to yellow-light wavelengths. Electronegativity differences determine whether a bond is ionic, polar covalent or nonpolar covalent, which in turn controls whether a substance dissolves in water, conducts electricity or forms a crystal. The atomic-radius trend explains why lithium-ion batteries use lithium: it is the smallest, lightest metal, so lithium ions can move quickly through the battery's electrolyte.
Students often think Z_eff increases as you go down a group. It does not: added inner shells screen the added nuclear charge, so Z_eff for the valence electrons stays roughly constant.
Students sometimes confuse ionization energy with electron affinity. Ionization energy is the cost of removing an electron (endothermic). Electron affinity is the energy released when adding one (exothermic). One is about losing, the other about gaining.
Students frequently assume that electron affinity and electronegativity are the same thing. Electron affinity is measured for an isolated gaseous atom. Electronegativity describes behaviour within a bond.
A common error is to say that atomic radius increases to the right across a period. The opposite is true: additional protons pull the same shell inward.
⚠️ Be prepared to rank a set of elements by any periodic property (IE, EA, EN, atomic radius) given only their positions on the table.
⚠️ Know the direction of every trend: left/right across a period and up/down a group.
⚠️ Understand why IE₂ > IE₁ and be able to explain it in terms of increased effective nuclear charge on the cation.
⚠️ Be ready to explain why cations are smaller than their parent atoms and anions are larger.
⚠️ The approximate electronegativity values for H, C, N, O and F are commonly tested.
⚠️ Expect questions linking Z_eff to at least one other property (radius, IE or EN).
True or false: ionization energy increases as you move down a group.
Fill in the blank: the most electronegative element is ______.
True or false: a cation is larger than its parent neutral atom.
Fill in the blank: electron affinity is an ______ (exothermic/endothermic) process.
True or false: Z_eff increases from left to right across a period.
Answers: 1. False (it decreases). 2. Fluorine. 3. False (smaller). 4. Exothermic. 5. True.
Q: Arrange the following in order of increasing atomic radius: Na, Mg, Al. Explain your reasoning.
A: Al < Mg < Na. All three are in Period 3. Moving left to right, Z_eff increases and pulls the electron cloud inward, so atomic radius decreases from Na to Al.
Q: Why is the second ionization energy of sodium dramatically higher than its first?
A: Sodium's first electron is removed from the 3s subshell (valence shell). The second electron must come from the 2p subshell (a core electron), which is much closer to the nucleus and far more tightly held. The jump crosses a principal energy level.
Q: Which has a larger ionic radius, F⁻ or Na⁺? Both are isoelectronic (10 electrons). Explain.
A: F⁻ is larger. Both have 10 electrons, but fluorine has only 9 protons pulling on them while sodium has 11. Fewer protons mean less inward pull, so the fluoride ion's electron cloud is more spread out.
Q: Why do noble gases have essentially zero electron affinity?
A: Noble gases have a full valence shell (complete octet). Adding an electron would mean placing it in a higher, less stable energy level, which requires energy input rather than releasing it.
Q: Oxygen and sulphur are in the same group. Which has the higher electronegativity, and why?
A: Oxygen. It is higher in the group (smaller atom, higher Z_eff on the valence electrons), so it attracts bonding electrons more strongly.
Periodic trends feed directly into chemical bonding: the electronegativity difference between two atoms determines whether a bond is ionic, polar covalent or nonpolar covalent. Ionization energy and electron affinity together predict the lattice energy of ionic compounds (covered in thermochemistry). Atomic and ionic radii are essential for understanding crystal structures and coordination chemistry in later courses.
Effective nuclear charge, Z_eff, shielding, screening, principal quantum number, atomic radius, ionic radius, cation, anion, ionization energy, first ionization energy, second ionization energy, electron affinity, electronegativity, Pauling scale, periodic trends, trend direction, active metals, noble gas stability, isoelectronic, core electrons, valence shell