Difficulty: Introductory to Intermediate | Prerequisites: Comfort with the layout of the periodic table. Familiarity with basic metric units helps for ionic radius comparisons.
Tags: atomic structure, protons, neutrons, electrons, mass number, isotopes, ions, periodic table, groups, periods, electron configuration, noble gas abbreviation, valence electrons, core electrons, electronegativity, ionic radius, diatomic molecules, CHM 11100, Purdue general chemistry
Everything in chemistry comes back to atoms and how their electrons are arranged. This set of topics covers the internal structure of atoms (protons, neutrons, electrons), how the periodic table organises elements by those structures, and how electron configurations determine an element's chemical behaviour. If you understand this material well, nomenclature, bonding and reactivity will make far more sense. This is the conceptual foundation of the entire course.
Atoms are defined by their proton count (atomic number). Neutrons affect mass but not identity. Electrons determine chemical behaviour and are arranged in shells and subshells following specific filling rules. The periodic table is organised so that elements in the same group share the same number of valence electrons and similar chemistry.
Atomic number (Z)
The number of protons in the nucleus of an atom. This defines the element. All carbon atoms have 6 protons; all nitrogen atoms have 7. In simple terms, the atomic number is the element's identity card.
Mass number (A)
The total number of protons plus neutrons in the nucleus. Electrons are not included because their mass is negligible. Think of it as the "heavyweight count" of the atom.
Isotopes
Atoms of the same element (same number of protons) that differ in the number of neutrons. Carbon-12 and carbon-14 are both carbon, but carbon-14 has two extra neutrons. In simple terms, isotopes are siblings with different weights.
Ion
An atom or group of atoms with a net electrical charge because it has gained or lost electrons. Cations are positively charged (lost electrons); anions are negatively charged (gained electrons).
Isoelectronic
Species (atoms or ions) that have the same number of electrons and therefore the same electron configuration. For example, Sr²⁺ and Kr both have 36 electrons.
Valence electrons
The electrons in the outermost (highest-energy) shell of an atom. These are the electrons involved in bonding and chemical reactions. The group number (for main-group elements) tells you the count directly.
Core electrons
All the electrons that are not valence electrons. These sit in filled inner shells and do not participate in bonding. Think of them as the stable, uninvolved inner layers.
Electron configuration
A notation showing how electrons are distributed among an atom's orbitals. Written in order of filling: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, and so on. The superscript gives the number of electrons in each subshell.
Noble gas (abbreviated) electron configuration
A shorthand that replaces all the inner-shell electrons with the symbol of the nearest preceding noble gas in square brackets. For example, silicon's full configuration is 1s²2s²2p⁶3s²3p²; the abbreviated form is [Ne]3s²3p².
Electronegativity
A measure of how strongly an atom attracts bonding electrons towards itself. Fluorine is the most electronegative element. The trend increases going up and to the right on the periodic table (excluding noble gases).
Ionic radius
The effective size of an ion. Cations are smaller than their parent atoms (lost electrons, same nuclear charge pulling fewer electrons in). Anions are larger (gained electrons, same nuclear charge spread over more electrons).
Diatomic molecule
A molecule made of exactly two atoms. Seven elements naturally exist as diatomic molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂. The mnemonic is "Have No Fear Of Ice Cold Beer" or simply "HOFBrINCl."
Protons: positive charge, located in the nucleus, mass ≈ 1 amu. The number of protons defines the element.
Neutrons: no charge, located in the nucleus, mass ≈ 1 amu. Changing the neutron count gives a different isotope but the same element.
Electrons: negative charge, located in orbitals around the nucleus, mass ≈ 0 amu (negligible). Responsible for chemical behaviour and bonding.
Mass number = protons + neutrons. Electrons are not counted.
Example: an atom with 6 protons, 6 neutrons and 6 electrons has mass number = 6 + 6 = 12.
The proton determines both the identity and the chemical properties of an element.
Same element, different number of neutrons
Notation: element symbol followed by mass number, e.g. Cl-37 means chlorine with mass number 37
To find neutrons: neutrons = mass number − atomic number
Cl-37: atomic number of Cl = 17, so neutrons = 37 − 17 = 20
Cs-134: atomic number of Cs = 55, so neutrons = 134 − 55 = 79
Ca-24: atomic number of Ca = 20, so neutrons = 24 − 20 = 4
Co-59: atomic number of Co = 27, so neutrons = 59 − 27 = 32
Of these, Cs-134 has the most neutrons (79).
Metals lose electrons to form cations (positive ions). Sodium (Na) loses 1 electron to become Na⁺.
Non-metals gain electrons to form anions (negative ions). Chlorine gains 1 electron to become Cl⁻.
The nitride ion N³⁻ has 7 protons (nitrogen's atomic number) and 10 electrons (7 + 3). Protons do not change when an ion forms.
When sodium (1s²2s²2p⁶3s¹) loses one electron, it becomes Na⁺ with configuration 1s²2s²2p⁶, which is the same as neon. This is why Na forms a +1 ion.
Sr²⁺ has 38 − 2 = 36 electrons
Full configuration of Sr²⁺: 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶ (same as krypton)
Abbreviated: [Kr] (but since it has lost the 5s² electrons, [Kr] alone, or more precisely the configuration that ends at 4p⁶)
The neutral element isoelectronic with Sr²⁺ is krypton (Kr), which also has 36 electrons
Periods (rows): elements in the same period have the same number of electron shells
Groups (columns): elements in the same group have the same number of valence electrons and similar chemical properties
Main-group families and their valence electron counts:
Group 1A (alkali metals): 1 valence electron, form +1 ions
Group 2A (alkaline earth metals): 2 valence electrons, form +2 ions
Groups 3A through 8A: valence electrons = group number
Group 4A: 4 valence electrons (carbon, silicon, germanium, etc.)
Group 5A: 5 valence electrons (nitrogen, phosphorus, etc.)
Group 7A (halogens): 7 valence electrons, form −1 ions
Group 8A (noble gases): 8 valence electrons (except helium with 2), very stable, rarely react
Transition metals occupy the d-block (groups 3B through 2B, or groups 3 through 12). They sit in the middle of the periodic table.
Mn (manganese), Mo (molybdenum) and W (tungsten) are all transition metals
Mg (magnesium) is not a transition metal; it is an alkaline earth metal in group 2A
Alkaline earth metals are in group 2A (column B on many simplified tables)
On a blank periodic table, the transition metals span the wide middle block (typically labelled column C in simplified diagrams)
Helium = He (correct)
Boron = B (not "Bo")
Magnesium = Mg (not "Mn," which is manganese)
Silicon = Si (not "S," which is sulfur)
Seven elements exist naturally as diatomic molecules: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂
Bromine (Br₂) is one of them. Lithium, phosphorus and boron are not diatomic.
Fill orbitals in order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p ...
Maximum electrons per subshell: s = 2, p = 6, d = 10, f = 14
Silicon (Z = 14): 1s²2s²2p⁶3s²3p² (total = 2+2+6+2+2 = 14 electrons)
Common wrong answers skip the 3s subshell or overfill the 2p. The 2p subshell holds a maximum of 6, never 8.
Abbreviated configuration uses the preceding noble gas in brackets:
[Ne]3s²3p¹ corresponds to aluminium (Al), because [Ne] accounts for 10 electrons plus 3 more = 13 = Al
[Ne]3s²3p² = silicon (Si, Z = 14)
Valence electrons are those in the highest principal energy level (outermost shell)
Core electrons are everything else
For 1s²2s²2p⁶3s²3p⁶4s² (calcium), the valence electrons are the 4s² (2 electrons). Core electrons = 20 − 2 = 18.
Be careful: the 3d subshell is empty here, so the highest principal quantum number is 4, and only 4s counts as valence.
Increases going up a group (smaller atoms hold electrons more tightly)
Increases going left to right across a period (more protons pull harder on bonding electrons)
Fluorine (F) is the most electronegative element
Francium (Fr), at the bottom-left corner of the table, is the least electronegative among common choices
Noble gases are typically excluded from electronegativity rankings because they rarely form bonds
Cations are smaller than their parent neutral atoms
Anions are larger than their parent neutral atoms
For isoelectronic species (same electron count), the ion with more protons is smaller because the greater nuclear charge pulls the electron cloud inward
Among the choices Al³⁺, Cl⁻, Rb⁺, Br⁻: Br⁻ has the largest ionic radius. It has 36 electrons but only 35 protons, giving it a large, diffuse electron cloud. Rb⁺ also has 36 electrons but 37 protons, pulling the cloud tighter.
Ordering Ca²⁺, K⁺, Cl⁻, S²⁻ from largest to smallest:
All four are isoelectronic with 18 electrons
S²⁻ has 16 protons pulling on 18 electrons (least pull per electron, largest)
Cl⁻ has 17 protons on 18 electrons
K⁺ has 19 protons on 18 electrons
Ca²⁺ has 20 protons on 18 electrons (most pull per electron, smallest)
Order: S²⁻ > Cl⁻ > K⁺ > Ca²⁺
Mass number (A) = number of protons + number of neutrons
Number of electrons in an ion = atomic number − charge (where a positive charge means fewer electrons)
Valence electrons for main-group elements = group number (using the A-group numbering)
Isotopes are used in medicine (radioactive tracers for imaging), archaeology (carbon-14 dating) and energy (uranium-235 in nuclear reactors). Electronegativity differences between atoms determine whether a bond is ionic, polar covalent or nonpolar covalent, which in turn controls everything from why water dissolves salt to how drugs bind to proteins.
Students often include electrons when calculating mass number. Electrons have negligible mass and are not counted. Mass number = protons + neutrons only.
Students confuse atomic number with mass number. The atomic number (protons only) defines the element. The mass number (protons + neutrons) identifies the isotope.
A common mistake is thinking that when an atom becomes an ion, the number of protons changes. Protons do not change. Only the electron count changes.
Students sometimes think the 3d subshell can hold 8 electrons (like 2p "holds 8"). The p subshell holds 6 at most; the d subshell holds 10.
⚠️ Questions about the number of protons and electrons in ions are very common. Remember: protons = atomic number (always), electrons = atomic number minus the charge.
⚠️ Electron configuration questions test whether you know the correct filling order. Watch out for the jump from 3p to 4s (not 3d).
⚠️ Ionic radius ordering for isoelectronic species is a classic exam question. More protons with the same number of electrons = smaller ion.
⚠️ Know which elements are transition metals and which are not. Mg is alkaline earth, not transition.
⚠️ Memorise the seven diatomic elements. This comes up repeatedly.
True or false: isotopes of the same element have different numbers of protons.
Fill in the blank: the mass number of an atom with 8 protons and 9 neutrons is ______.
True or false: Na⁺ has more electrons than neutral Na.
Fill in the blank: phosphorus has ______ valence electrons.
True or false: Cl⁻ is larger than K⁺, even though both have 18 electrons.
Answers: 1. False (they have different numbers of neutrons). 2. 17. 3. False (it has one fewer). 4. 5. 5. True (Cl⁻ has fewer protons pulling on those 18 electrons).
Q: What is the mass number of an atom with 6 protons, 6 neutrons and 6 electrons?
A: 12. Mass number = protons + neutrons = 6 + 6. Electrons are not included.
Q: The nitride ion (N³⁻) has how many protons and how many electrons?
A: 7 protons and 10 electrons. Nitrogen's atomic number is 7 (always 7 protons). The 3− charge means 3 extra electrons: 7 + 3 = 10.
Q: Which of the following has the most neutrons: Cl-37, Cs-134, Ca-24 or Co-59?
A: Cs-134. Neutrons = 134 − 55 = 79.
Q: Which family of elements has 2 valence electrons?
A: The alkaline earth metals (group 2A).
Q: What is the correct electron configuration for silicon?
A: 1s²2s²2p⁶3s²3p². Total of 14 electrons, matching silicon's atomic number.
Q: What element has the abbreviated electron configuration [Ne]3s²3p¹?
A: Aluminium (Al). [Ne] = 10 electrons, plus 3 more = 13 = Al.
Q: Sodium (1s²2s²2p⁶3s¹) forms what type of ion, and what is the resulting configuration?
A: It forms a +1 ion (Na⁺) by losing the 3s¹ electron. The resulting configuration is 1s²2s²2p⁶.
Q: How many core electrons are in an atom with configuration 1s²2s²2p⁶3s²3p⁶4s²?
A: 18. The valence electrons are 4s² (2 electrons). Core = 20 − 2 = 18.
Q: Which of the following is the least electronegative: N, H, F, Fr or At?
A: Fr (francium). It sits at the bottom-left of the periodic table, the position of lowest electronegativity.
Q: Place Ca²⁺, K⁺, Cl⁻ and S²⁻ in order from largest to smallest ionic radius.
A: S²⁻ > Cl⁻ > K⁺ > Ca²⁺. All are isoelectronic with 18 electrons. More protons means a smaller radius.
Electron configuration directly determines how atoms bond (covered in the bonding and Lewis structures unit). Valence electron counts feed into nomenclature rules for ionic compounds, because you need to know the charge an element will carry. Electronegativity differences determine bond polarity, which leads to molecular polarity and intermolecular forces later in the course.
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