Understanding Atoms
Everything in chemistry follows from atomic structure. This guide builds from subatomic particles through electron configuration, periodic trends, and ion formation — with the reasoning behind every pattern.
All matter is made of atoms. Every atom has a tiny, dense nucleus containing protons and neutrons, surrounded by electrons in shells. The number of protons is not just a label — it is the element. Swap the proton count and you have a completely different element with different chemistry.
The number of protons determines the nuclear charge, which sets how many electrons the neutral atom has, which determines every chemical property. Change one proton and the entire electron configuration restructures — you have a new element. Neutrons add mass and can affect nuclear stability (isotopes), but they barely change chemical behaviour because chemistry happens in the electron shells.
Z = atomic number = number of protons (identifies the element)
A = mass number = protons + neutrons
neutrons = A − Z
Isotopes = same Z, different A (different number of neutrons)
Cl-35: neutrons = 18. Cl-37: neutrons = 20.
= 26.52 + 8.97
= 35.48 u
The periodic table shows Cl = 35.45 u — a decimal, never a whole number, because it is a weighted average of isotopes.
b) Two atoms belong to the same element but have different mass numbers. What specifically differs between them?
c) Why do isotopes have nearly identical chemical properties?
a) Z=8 is Oxygen (O). Neutrons = A − Z = 16 − 8 = 8. Electrons = Z = 8.
b) They have different numbers of neutrons (and therefore different mass numbers A). These are isotopes — same element, same proton count, different neutron count.
c) Chemical properties are determined by electron configuration, which depends on the number of electrons (= number of protons = Z). Isotopes have the same Z and therefore the same electron arrangement and virtually identical chemical behaviour. Neutrons affect nuclear stability and mass, not chemistry.
Electrons do not randomly occupy space around the nucleus. They fill orbitals in a precise order governed by energy. The outermost electrons (valence electrons) determine how an atom interacts with other atoms — bonding, reactivity, and ion formation all stem from the valence electron count.
Systems naturally minimise energy. An electron in a higher-energy orbital is less stable. Left to find its own level, it will drop to the lowest available orbital. The rules (Aufbau, Pauli, Hund) are just the systematic description of electrons doing exactly this. The one exception — 4s filling before 3d — happens because 4s is actually lower in energy than 3d for most atoms.
Aufbau: fill lowest energy orbitals first. Order: 1s 2s 2p 3s 3p 4s 3d 4p 5s...
Pauli exclusion: max 2 electrons per orbital; they must have opposite spins.
Hund's rule: when orbitals of equal energy exist (e.g. three 2p orbitals), place one electron in each before pairing any.
Total so far: 2+2+6+2 = 12 ✓
Outermost shell is n=3. There are 2 electrons in 3s. Mg has 2 valence electrons — hence it is in Group 2.
[Ar]3d&sup6;
[Ar]3d&sup5;
[Ar]3d&sup5; is a half-filled d subshell (one electron in each of 5 orbitals). Half-filled subshells have extra stability because electrons are evenly spread, minimising repulsion. Fe³+ with [Ar]3d&sup5; is actually more stable than Fe²+ with [Ar]3d&sup6;.
b) Write the shorthand (noble gas) configuration for calcium (Z=20).
c) Fe³+ has configuration [Ar]3d&sup5; rather than [Ar]3d&sup6;. Why does Fe lose its 4s electrons before losing 3d electrons?
a) P (Z=15): 1s² 2s² 2p&sup6; 3s² 3p³ — total = 2+2+6+2+3 = 15 ✓. Note: the three 3p electrons each occupy a separate orbital (Hund's rule).
b) Ca (Z=20): [Ar]4s². Ar has 18 electrons, then 4s fills with 2 more.
c) When forming ions, electrons are removed from the outermost shell first (highest principal quantum number). For Fe, the 4s electrons (n=4) are in the outermost shell even though 4s filled before 3d. Once you are removing electrons, the 4s is the most exposed and least tightly held, so they go first.
Mendeleev noticed that when elements are arranged by properties, patterns repeat at regular intervals — this is periodicity. Today we know the repeating pattern exists because elements in the same group have the same number of valence electrons, and valence electrons dictate chemical behaviour.
Every element in Group 1 has 1 valence electron. Every element in Group 17 has 7 valence electrons. They all behave similarly because their outermost electron count is the same. As you move down a group, you add electron shells, but the valence count stays constant — so properties stay similar. Periodicity is just electron configuration made visual.
Period (row) = all atoms have the same number of electron shells (n = period number)
Group (column) = all atoms have the same number of valence electrons
Group number = valence electron count (for main-group elements)
Alkali metal: very reactive (only 1 e− to lose), forms +1 ion (Na⁺), similar properties to Li (above) and K (below).
As sits on the metalloid staircase — it is a metalloid (semiconductor-like properties). Group 15 with 5 valence electrons: can form −3 ion or share 3 electrons in covalent bonds.
As shares chemical properties with nitrogen (Z=7) and phosphorus (Z=15) — same group, same valence count.
b) Which group of the periodic table always has exactly 2 valence electrons?
c) Why do elements in the same group have similar chemical properties?
a) 2 shells = Period 2. 7 valence electrons = Group 17. That is Fluorine (F), the most electronegative element.
b) Group 2 (alkaline earth metals) — always 2 valence electrons in ns² configuration.
c) Elements in the same group have the same number of valence electrons. Since chemical reactions involve gaining, losing, or sharing valence electrons, elements with identical valence counts react in the same ways and form the same types of compounds. The inner electron shells differ (more shells as you go down), but they are shielded from chemistry by the outer shell.
Every periodic trend can be explained by the same two competing forces: nuclear charge (the positive protons pulling electrons inward) and electron shielding (inner electrons blocking some of that pull). The direction of each trend depends on which force wins in a given direction.
Going from Na to Cl across Period 3, the number of protons increases from 11 to 17. All those electrons remain in the same shell (n=3). The extra protons pull the electron cloud inward with more force, but the shielding barely changes because no new inner shells are added. Result: the atom shrinks, electrons are held tighter (higher IE), and the atom attracts bonding electrons more strongly (higher EN).
Why shielding wins down a group:
Going from F to Cl to Br in Group 17, each step adds a completely new electron shell between the nucleus and the valence electrons. Those inner shells shield the valence electrons from the nuclear charge. Even though the nucleus gets more protons, the valence electrons “feel” less pull. Result: the atom grows, electrons are easier to remove, and EN decreases.
Atomic radius: increases down, decreases across (opposite of IE and EN)
Ionization energy: decreases down, increases across
Electronegativity: decreases down, increases across
Highest EN: F = 4.0 Lowest EN: Fr ≈ 0.7
(More protons in Cl pull electrons in tighter)
(Cl holds its electrons tighter; Cl is a much stronger electron attractor)
Ca: Period 4, Group 2
Br: Period 4, Group 17
K has only 1 valence electron, loosely held by 19 protons — easiest to remove. Ca has 2 valence electrons, tighter grip. Br has 17 protons pulling on 7 valence electrons in the same shell — hardest to remove of the three.
b) Arrange F, Cl, and Br in order of decreasing electronegativity.
c) Why does ionization energy generally increase across a period?
a) K has a larger atomic radius. Both are in Group 1 (same number of valence electrons), but K is in Period 4 and has 4 electron shells compared to Na's 3. The extra shell pushes the outermost electrons farther from the nucleus, more than offsetting the greater nuclear charge.
b) F > Cl > Br (decreasing EN down Group 17). F is the most electronegative element on the table (4.0). Going down the group, more electron shells shield the nucleus from bonding electrons, so each element attracts shared electrons less strongly.
c) Across a period, the number of protons increases while electrons are added to the same shell (no new inner shells, so shielding barely changes). The increased nuclear charge pulls valence electrons more tightly, making them harder to remove — hence higher ionization energy.
Atoms form ions to achieve the electron configuration of the nearest noble gas — a full outer shell. This is the most stable arrangement because filled shells have symmetrical electron distribution and strong shielding. Metals lose electrons to reach the configuration of the noble gas before them; nonmetals gain electrons to reach the configuration of the noble gas after them.
Metals (left side) have 1–3 valence electrons. It costs relatively little energy to remove a small number of electrons. Once lost, the atom reaches a full inner shell — stable. Nonmetals (right side) have 5–7 valence electrons. They are so close to a full shell that gaining 1–3 electrons is energetically favourable — the nucleus pulls additional electrons in readily. The driving force is always: reach the nearest noble gas configuration.
Cation = atom loses e− → positive charge (metals, Groups 1–13)
Anion = atom gains e− → negative charge (nonmetals, Groups 15–17)
Cations are smaller than the parent atom (fewer electrons, same nuclear charge)
Anions are larger than the parent atom (more electrons, same nuclear charge — more repulsion)
No. Na+ has 10 electrons (like Ne). Cl− has 18 electrons (like Ar). They have different electron counts, so they are NOT isoelectronic with each other — though each is isoelectronic with its respective noble gas.
A half-filled d subshell ([Ar]3d&sup5;) has one electron in each of the five d orbitals. This symmetric, evenly-spread arrangement minimises electron-electron repulsion and has extra stability (similar to why half-filled p subshells are slightly more stable). Fe³+ is therefore unusually stable for a +3 ion.
b) Na+ and Mg²+ both have 10 electrons. Which ion is smaller, and why?
c) Why do Group 17 elements form −1 ions but Group 16 elements form −2 ions?
a) Ca (Z=20): [Ar]4s². Ca²+ loses both 4s electrons: [Ar] — same configuration as argon, 18 electrons.
b) Mg²+ is smaller. Both ions have 10 electrons. However, Mg²+ has 12 protons and Na+ has only 11. More protons pulling the same number of electrons inward means the electron cloud is contracted more tightly in Mg²+. In isoelectronic series, the ion with more protons is always the smallest.
c) Group 17 elements have 7 valence electrons — they need to gain just 1 more electron to complete the octet (full outer shell). Group 16 elements have 6 valence electrons — they need to gain 2 more electrons to complete the octet. The charge on the ion reflects exactly how many electrons were gained to achieve the nearest noble gas configuration.