Atoms, Electrons & Periodicity — Print · Physical Science: Chemistry, Grade 11

Atomic Models & Subatomic Particles

Learning goals

• I can trace how the model of the atom changed from Dalton through Thomson, Rutherford, Bohr, and the quantum-mechanical picture, and explain what experiment forced each revision.
• I can state the charge, relative mass, and location of protons, neutrons, and electrons.
• I can use atomic number and mass number (including isotope notation) to find the numbers of protons, neutrons, and electrons in an atom or ion.

Readings for this unit

• From a spectrum to an atomic model

The idea that matter is made of tiny indivisible pieces is ancient, but chemistry became a science when John Dalton turned that hunch into a testable theory around 1803. Dalton proposed that each element is made of identical atoms, that atoms of different elements differ in mass, and that they combine in simple whole-number ratios to form compounds. His model pictured the atom as a solid, featureless sphere — a tiny billiard ball. That picture was wrong in its details, but it was powerful: it explained why compounds always form in fixed proportions, and it gave scientists something concrete to test and eventually improve.

The billiard-ball atom cracked open in 1897, when J.J. Thomson studied cathode rays and discovered a particle far lighter than any atom — the electron. If atoms contained even smaller negative particles, they could not be indivisible. Because whole atoms are neutral, Thomson reasoned there must be balancing positive charge, and he proposed the 'plum-pudding' model: a diffuse ball of positive charge with electrons scattered through it like raisins in dough. It was the first model to put internal structure inside the atom.

In 1909 Ernest Rutherford and his students aimed a beam of positively charged alpha particles at an ultra-thin sheet of gold foil, expecting them all to sail straight through the 'pudding.' Most did — but a stunning few ricocheted at sharp angles, some bouncing almost straight back. Rutherford famously compared it to a cannonball rebounding off tissue paper. The only explanation was that an atom's positive charge and nearly all its mass are concentrated in a minuscule central nucleus, with the electrons and mostly empty space surrounding it. The nuclear atom was born.

Rutherford's model still could not explain why atoms emit only specific colors of light, so in 1913 Niels Bohr refined it: electrons travel in fixed circular orbits, or energy levels, at set distances from the nucleus, and they jump between levels by absorbing or releasing exact amounts of energy. Bohr's model brilliantly predicted the hydrogen spectrum. Yet electrons are not tidy planets. In the 1920 s Erwin Schrödinger and others built the quantum-mechanical model we still use: electrons occupy fuzzy, three-dimensional regions of probability called orbitals, where we can describe only the likelihood of finding an electron, not a precise path. Notice the pattern — each model kept what worked and repaired the one flaw a new experiment exposed. That is exactly how scientific models grow.

Underneath all these pictures sit just three particles. Protons carry a +1 charge and a mass of about 1 amu; neutrons carry no charge and also weigh about 1 amu; both are packed into the nucleus. Electrons carry a −1 charge but are almost weightless — roughly 1/1836 the mass of a proton — so essentially all of an atom's mass lives in that tiny nucleus while essentially all of its volume is electron cloud. The number of protons, the atomic number (Z), is an element's fingerprint: every atom with 6 protons is carbon, and nothing with 6 protons can be anything but carbon.

Two counts unlock every atom. The mass number (A) is protons plus neutrons, so neutrons = A − Z. In a neutral atom the electrons equal the protons, balancing the charge to zero. Chemists write this with isotope notation, stacking the mass number and atomic number to the left of the symbol — for example, the sodium symbol written with mass number 23 and atomic number 11 tells you 11 protons, 11 electrons, and 23 − 11 = 12 neutrons at a glance. When an atom gains or loses electrons it becomes an ion with a net charge, but its proton count — its identity — never changes. Mastering this bookkeeping is the foundation for everything that follows, from isotopes to electron configurations to the whole periodic table.

Model check: A shell diagram is an organizing model, not a literal path for an electron. Charge=17−18=−1; atomic number 17 identifies chlorine, so Cl⁻.

Number of Neutrons

neutrons = mass number (A) − atomic number (Z)

atomic number Z = number of protons = number of electrons in a neutral atom; mass number A = protons + neutrons.

Worked examples

Avoid these mistakes

• A shell diagram is an organizing model, not a literal path for an electron.

Words to know

atom

The smallest particle of an element that retains the chemical identity of that element.

proton

A positively charged (+1) subatomic particle in the nucleus, with a relative mass of about 1 amu.

neutron

An electrically neutral subatomic particle in the nucleus, with a relative mass of about 1 amu.

electron

A negatively charged (−1) subatomic particle occupying the space around the nucleus, with a mass roughly 1/1836 that of a proton.

nucleus

The tiny, dense, positively charged core of an atom that holds the protons and neutrons and nearly all of its mass.

atomic number (Z)

The number of protons in an atom's nucleus; it defines which element the atom is.

mass number (A)

The total count of protons plus neutrons in an atom's nucleus.

ion

An atom that has gained or lost electrons, giving it a net negative or positive charge.

atomic mass unit (amu)

The standard unit of atomic-scale mass, defined so that one carbon-12 atom has a mass of exactly 12 amu.

How Thomson caught a particle no one could see

Thomson never actually saw an electron — he inferred it. Firing cathode rays through electric and magnetic fields, he measured how far the beam bent and worked out the particles' charge-to-mass ratio, which came out thousands of times larger than any whole atom's. A decade later Robert Millikan floated charged oil droplets between electric plates and pinned down the electron's charge itself, about 1.6 × 10⁻¹⁹ coulombs. Together the two experiments proved a piece of the atom smaller than the atom — and ended the 'indivisible' idea for good.

Just how empty is an atom?

If you shrank a nucleus to the size of a pea and set it at the center of a large sports stadium, the electrons would be whirling somewhere up in the back rows, with nothing but empty space in between. Atoms are roughly 99.99999% empty, yet a desk still feels solid because the electron clouds of its atoms repel the electron clouds in your hand before they ever touch. You are not truly resting on a chair so much as hovering a hair above it on a cushion of electric repulsion.

Even protons have parts: the quark surprise

Dalton thought the atom was the end of the line, but we now know that even protons and neutrons are built from smaller particles called quarks, bound together by the strong nuclear force. A proton is two 'up' quarks and one 'down' quark (charges +⅔ + ⅔ − ⅓ = +1); a neutron is one up and two down (+⅔ − ⅓ − ⅓ = 0). The pattern of this whole unit repeats yet again — every time we think we have found the smallest piece, a sharper experiment finds structure hiding inside it.

Three ways to change an atom (only one changes the element)

Students blur these together constantly, so keep them apart. Change the neutrons and you get a different isotope of the same element; change the electrons and you get an ion with a charge; but change the protons and you have literally made a different element. That is why a carbon-14 atom is still carbon, and why a sodium ion (Na⁺) is still sodium — the proton count, the atomic number, is the single number that fixes identity.

Rutherford's alpha particles are in your hallway

The very same alpha particles Rutherford fired at gold foil are quietly working in most household smoke detectors right now. A microscopic amount of the element americium-241 steadily emits alpha particles that ionize the air inside the detector, letting a tiny current flow; when smoke drifts in, it disrupts that current and the alarm sounds. A hundred-year-old physics experiment stands guard over millions of homes every single night.

In the real world

Rutherford-style particle scattering is still how physicists probe matter today: giant accelerators fire particles at targets and read the deflections to map structures far too small to see.

Exit check

An ion has 17 protons and 18 electrons. Identify charge and element.

Isotopes & Average Atomic Mass

Learning goals

• I can define isotopes and explain why they share chemistry but differ in mass.
• I can calculate the average atomic mass of an element from its isotopes' masses and percent abundances.
• I can explain why the average atomic mass sits closest to the most abundant isotope.

Readings for this unit

• From a spectrum to an atomic model

Not every atom of an element is identical. Isotopes are atoms that share the same number of protons — so they are the same element — but carry different numbers of neutrons, and therefore different mass numbers. Hydrogen is the classic case: ordinary hydrogen has zero neutrons, deuterium has one, and tritium has two, yet all three are still hydrogen because each has exactly one proton. Carbon has three natural isotopes as well: carbon-12, carbon-13, and carbon-14, differing only in neutron count.

Isotopes of the same neutral element share electron count and broadly similar electronic structure, so many chemical properties are similar. Different nuclear masses can nevertheless change vibrational behavior, equilibrium and reaction rates through isotope effects. Nuclear stability can also differ. Separate these effects instead of claiming extra neutrons never influence chemistry or that every isotope is radioactive.

Look closely at any periodic table and you will notice the atomic masses are almost never tidy whole numbers: chlorine is listed as 35.45, copper as 63.55, boron as 10.81. That is a clue. A single chlorine atom always has a whole number of protons and neutrons, so where does 35.45 come from? The listed value is not the mass of any one atom — it is the average atomic mass, a weighted average of the masses of all the element's naturally occurring isotopes, weighted by how common each one is.

The word 'weighted' is what makes this different from an ordinary average. You do not simply add the isotope masses and divide by how many there are. Instead you multiply each isotope's mass by its fractional abundance (its percent abundance written as a decimal) and add the results: average atomic mass = (mass₁ × fraction₁) + (mass₂ × fraction₂) + … A rare isotope barely nudges the total, while an abundant one dominates it. Chlorine is about 76 percent chlorine-35 and only 24 percent chlorine-37, so the average lands at 35.45 — much closer to 35.

Average atomic mass is a weighted mean of the measured isotope masses. Multiply each isotope mass by its fractional abundance, then add all contributions; fractions must total 1. The mean lies between the smallest and largest masses present. For a mixture of two isotopes it lies closer to the more abundant isotope, but with three or more isotopes the nearest isotope is not necessarily the most abundant. Use the full weighted calculation rather than rounding the mean to guess composition.

Isotope masses can be expressed in unified atomic mass units, u, defined through one twelfth of the mass of a carbon-12 atom in its specified state. Mass spectrometry separates ions through a mass-to-charge relationship; instrument design, charge states and response calibration matter when inferring abundance. The resulting sample-weighted value need not be identical for every natural source or sample of an element.

Model check: Mass number counts nucleons; average atomic mass is a weighted value, not a nucleon count. 10(.20)+11(.80)=10.8u. The result lies between the two masses and closer to the more abundant one.

Average Atomic Mass (weighted average)

average atomic mass = (mass₁ × fraction₁) + (mass₂ × fraction₂) + …

each mass is an isotope's mass in amu; each fraction is that isotope's percent abundance ÷ 100; the fractional abundances add up to 1.

Worked examples

Avoid these mistakes

• Mass number counts nucleons; average atomic mass is a weighted value, not a nucleon count.

Words to know

isotopes

Atoms of the same element (same number of protons) that have different numbers of neutrons, and therefore different mass numbers.

percent abundance

The percentage of an element's atoms found in nature that are a particular isotope.

fractional abundance

A percent abundance written as a decimal (percent abundance ÷ 100); it is what you actually multiply by in the average.

average atomic mass

The weighted average mass of all naturally occurring isotopes of an element, weighted by each isotope's abundance.

weighted average

An average in which each value is multiplied by how common it is, so more abundant values count more.

mass spectrometer

An instrument that sorts atoms by mass and measures the relative abundance of each isotope.

Inside a mass spectrometer

A mass spectrometer is essentially a racetrack for atoms. It strips electrons off a sample to make positive ions, accelerates them, then bends their path with a magnet — and here is the trick: heavier ions swing wide while lighter ones curve tightly, so each isotope lands in a different spot. Counting how many ions strike each spot gives the percent abundances that feed the weighted average, which is exactly how we know chlorine is about 76% chlorine-35 and 24% chlorine-37.

The carbon-14 clock

Living things constantly take in carbon, including a trace of radioactive carbon-14. When an organism dies it stops refreshing that supply, and the carbon-14 already inside it decays at a steady rate — half of it every 5,730 years. By measuring how much carbon-14 is left in a bone, a scrap of linen, or the charcoal of an ancient fire, archaeologists read off how long ago it was alive, back to roughly 50,000 years. Isotopes, it turns out, keep time.

The only isotopes with their own names

Hydrogen is so light that its isotopes behave noticeably differently, so chemists gave them their own names: ordinary protium (no neutrons), deuterium (one neutron), and tritium (two neutrons). Water built from deuterium — heavy water, D₂O — is dense enough that ice cubes of it sink instead of float, and it plays a genuine role in some nuclear reactors. Deuterium and tritium are also the fuel that fusion researchers hope will one day power the electric grid.

The mistake almost everyone makes first

The single most common error is taking a plain average instead of a weighted one. For boron (isotope masses 10 and 11) students often write (10 + 11) ÷ 2 = 10.5 — but boron-11 is four times more common, so the true average is 10.8, pulled toward the abundant isotope. Always multiply each mass by its fractional abundance first, and lean on the built-in check: the answer must land between the lightest and heaviest isotope, nearest the most common one.

Isotopes as forensic fingerprints

Because different places and processes leave slightly different isotope ratios, those ratios act like a hidden fingerprint. The oxygen isotopes in your hair record the water you drank and therefore roughly where you have lived; the carbon and nitrogen isotopes in food can expose fake 'local' honey or mislabeled organic beef; investigators even use isotope analysis to catch athletes doping with synthetic testosterone. Same-element atoms that behave identically in chemistry still quietly carry a story locked in their mass.

In the real world

Mass spectrometers that measure isotope abundances let scientists carbon-date fossils, trace pollution back to its source, and even tell where a food or an athlete's sample truly came from.

Exit check

Isotopes 10u at 20% and 11u at 80% give what average?

Electron Configuration

Learning goals

• I can describe how electrons are organized into energy levels, sublevels (s, p, d, f), and orbitals.
• I can write ground-state electron configurations using the Aufbau order, including noble-gas shorthand.
• I can identify an atom's valence electrons from its configuration.

Readings for this unit

• From a spectrum to an atomic model

The quantum model describes electrons using orbitals and allowed energy states rather than fixed circular paths. The principal quantum number n labels shells and relates to orbital size and energy. For a one-electron atom, shell energy depends on n; in many-electron atoms, sublevel energies also depend on interactions and shielding. An electron in a larger-n shell is not automatically higher in energy than every smaller-n orbital. Build configurations using the appropriate ground-state filling pattern, not a literal picture of stadium tiers.

Each energy level is divided into sublevels named s, p, d, and f, and each sublevel is built from orbitals — regions of space that, by the Pauli exclusion principle, hold at most two electrons apiece. An s sublevel has 1 orbital (2 electrons), a p sublevel has 3 orbitals (6 electrons), a d sublevel has 5 orbitals (10 electrons), and an f sublevel has 7 orbitals (14 electrons). Add these up level by level and a neat rule appears: the maximum number of electrons an energy level can hold is 2n². Level 1 holds 2, level 2 holds 8, level 3 holds 18.

The Aufbau approach builds approximate ground-state configurations by filling lower-energy orbitals before higher-energy ones. A useful introductory sequence is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p. Orbital energies depend on occupation, so the statement that 4s is always below 3d is not universal. Some neutral atoms have exceptions, and transition-metal cations normally lose the highest-n s electrons before d electrons. Hund’s rule places electrons singly with parallel spins in equivalent orbitals before pairing in the ground-state model.

An electron configuration records the result. You write each occupied sublevel in order and use a superscript for how many electrons it holds, and those superscripts must add up to the total number of electrons. Oxygen (8 electrons) is 1 s² 2 s² 2p⁴ — check: 2 + 2 + 4 = 8. Writing out long configurations gets tedious, so chemists use noble-gas shorthand: replace the inner electrons with the previous noble gas in brackets. Calcium's full form, 1 s² 2 s² 2p⁶ 3 s² 3p⁶ 4 s², collapses to the tidy [Ar] 4 s².

For main-group elements, valence electrons are usually the electrons in the outermost occupied shell. Sulfur has 3s²3p⁴ in its highest-n shell, giving six valence electrons. Inner electrons contribute shielding and are not literally inactive; transition-metal bonding can also involve d electrons from a lower principal shell. A simple outer-shell count helps identify many main-group patterns, but it is not a complete model of every element’s chemical behavior.

This is the hinge of the whole unit. The periodic table is, quite literally, a picture of electron configurations: its rows correspond to energy levels being filled and its columns to matching valence arrangements. Once you can build a configuration, you can read an element's place on the table — and predict its behavior — straight from its electrons. That connection is exactly what the essential question is driving at: the arrangement of electrons is what the entire periodic table is organizing.

Model check: The 2n² shell capacity is not the actual filling order of sublevels. 1 s²2 s²2p⁶3s¹; eleven electrons total and one outer-shell electron.

Maximum Electrons in an Energy Level

maximum electrons = 2n²

n is the principal energy level (n = 1, 2, 3, …); for example n = 3 holds 2 × 3² = 18 electrons. Sublevel capacities: s = 2, p = 6, d = 10, f = 14.

Worked examples

Avoid these mistakes

• The 2n² shell capacity is not the actual filling order of sublevels.

Words to know

energy level (n)

A principal shell around the nucleus; higher n means higher energy and, on average, greater distance from the nucleus.

sublevel

A subdivision of an energy level, labeled s, p, d, or f, each with a characteristic shape and capacity.

orbital

A region within a sublevel that can hold a maximum of two electrons.

electron configuration

The notation showing how many electrons occupy each sublevel of an atom, such as 1 s² 2 s² 2p⁶.

Aufbau principle

Electrons fill the lowest-energy available sublevel first before moving to higher ones.

Pauli exclusion principle

No orbital can hold more than two electrons, and those two must have opposite spins.

Hund's rule

Within a sublevel, electrons occupy separate orbitals singly before any orbital is doubled up.

valence electrons

Electrons relevant to the chemical bonding model. Main-group outer-shell patterns are a useful case; transition elements can involve nearby d states.

Where fireworks and flame tests get their colors

Heat an atom and its electrons jump up to higher energy levels; when they fall back down, they release the exact energy difference as a particle of light of one specific color. Because every element has a unique ladder of energy levels, each emits its own signature colors — sodium's deep yellow, copper's blue-green, lithium's crimson. This is the fingerprint behind flame tests and firework recipes, and by reading these same emission lines in starlight, astronomers can name the elements in a star without ever leaving Earth.

The rule-breakers: chromium and copper

Electron configurations follow the Aufbau order beautifully — until they don't. Chromium 'should' be [Ar] 4 s² 3d⁴ but is really [Ar] 4s¹ 3d⁵, and copper is [Ar] 4s¹ 3d¹⁰ instead of 4 s² 3d⁹. The reason is that a half-filled or completely filled d sublevel is unusually stable, so the atom borrows one electron from 4 s to reach it. These exceptions are a healthy reminder that the filling rules are an excellent model, not an unbreakable law.

Orbitals are clouds, not racetracks

It is tempting to picture electrons looping around the nucleus like tiny planets, but the quantum model says otherwise: an orbital is a three-dimensional region of probability, not a path. An s orbital is a fuzzy sphere; a p orbital is shaped like a dumbbell. We can map where an electron is likely to be, but never its precise route, because the Heisenberg uncertainty principle says you cannot pin down an electron's exact position and motion at the same time. Bohr's tidy orbits were a useful ladder we then had to climb beyond.

Why the periodic table has the shape it has

The blocks of the periodic table are electron sublevels in disguise. An s sublevel holds 2 electrons, so the s-block is 2 columns wide; a p sublevel holds 6, so the p-block is 6 wide; the d-block is 10 columns and the f-block 14. Line those widths up and you get the exact silhouette of the table — the two tall columns on the left, the broad ten-wide transition block, the detached f-block below. The table's strange shape is nothing more than electron configuration drawn out on paper.

Spin: why exactly two electrons share an orbital

Each electron carries a property called spin, which comes in only two varieties, usually drawn as up and down arrows. The Pauli exclusion principle says two electrons may share an orbital only if their spins are opposite — which is precisely why an orbital tops out at two. Spin is far more than bookkeeping: the unpaired electron spins in iron are what make a magnet magnetic, and the aligned spins of atomic nuclei are exactly what an MRI scanner reads to build an image of the body.

In the real world

The specific colors of neon signs, fireworks, and streetlights come from electrons dropping between exact energy levels — each element releases its own signature colors, which is how astronomers identify what distant stars are made of.

Exit check

Write sodium ground-state configuration and valence count.

The Periodic Table

Learning goals

• I can explain how the periodic table is organized by atomic number into periods, groups, and s/p/d/f blocks.
• I can connect an element's position to its electron configuration and valence electrons.
• I can classify elements as metals, nonmetals, or metalloids and name the major families.

Readings for this unit

• From a spectrum to an atomic model

When Dmitri Mendeleev arranged the known elements by their properties in 1869, he built the first periodic table and even left gaps for elements not yet discovered, predicting their properties in advance. The modern table refines his insight by ordering elements strictly by atomic number — increasing proton count from left to right, top to bottom. The word 'periodic' captures the magic: chemical properties repeat at regular intervals, so similar elements line up in the same columns like a pattern that keeps coming back.

Within a main-group column, elements generally have similar outer-electron configurations and often similar chemical behavior. Helium is an important exception to a simple group-number count, and transition-metal groups require a more detailed treatment of s and d electrons. Periods organize recurring configuration patterns as atomic number increases; their number connects to the largest principal shell occupied in the usual introductory ground-state configurations. Similar patterns support predictions but do not make members of a group identical.

The table's shape mirrors the sublevels from the previous lesson, dividing into four blocks. The tall left columns (Groups 1–2) are the s-block, where an s sublevel is filling. The right columns (Groups 13–18) are the p-block. The broad center (Groups 3–12) is the d-block, home of the transition metals, and the two detached rows at the bottom are the f-block. Read this way, an element's address on the table is essentially a shorthand for its electron configuration: position and electrons are two views of the same thing.

The table also sorts elements by character. Metals — the great majority, filling the left and center — are shiny, malleable, and conduct heat and electricity well because their outer electrons move freely. Nonmetals, clustered in the upper right, are usually dull, brittle when solid, and poor conductors. Straddling the zigzag staircase between them are the metalloids like silicon and germanium, which behave partway between the two — exactly the property that makes them ideal semiconductors for computer chips.

Certain columns are important enough to have family names. The alkali metals (Group 1) are so reactive they are stored under oil. The alkaline earth metals (Group 2) are reactive but less so. The halogens (Group 17) are one electron short of a full shell, making them aggressive electron-grabbers that form salts. And the noble gases (Group 18) have complete valence shells, which leaves them famously unreactive — useful precisely because they refuse to react. Knowing which family an element belongs to lets you predict its behavior before any experiment.

This is the payoff of the essential question. Because the table is arranged by electron configuration, its geometry encodes chemistry: an element's row tells you its highest energy level, its column tells you its valence count, and its block tells you which sublevel is filling. From that single position you can anticipate how an element bonds, what charge its ion will take, and — as the next lesson shows — how properties like size and reactivity slide smoothly across the table.

Model check: The simple group-number valence rule applies to main-group patterns, not all transition elements. Outer n=3 gives period 3; five valence electrons give main-group 15; this is phosphorus.

Valence Electrons of a Main-Group Element

valence electrons = group number − 10 (for groups 13–18)

applies to main-group (representative) elements; for groups 1–2 the valence-electron count simply equals the group number. This shortcut does not apply to the transition (d-block) metals.

Worked examples

Avoid these mistakes

• The simple group-number valence rule applies to main-group patterns, not all transition elements.

Words to know

period

A horizontal row of the periodic table; its elements have their outermost electrons in the same highest energy level (same n).

group (family)

A vertical column of the periodic table; its elements share the same number of valence electrons and similar chemical behavior.

block (s, p, d, f)

A region of the table named for the sublevel that is being filled across it, such as the d-block for transition metals.

metal

An element that is typically shiny, malleable, and a good conductor; metals fill the left and center of the table.

nonmetal

An element that is typically dull, brittle, and a poor conductor; nonmetals cluster in the upper right.

metalloid

An element along the staircase line with a mix of metal and nonmetal properties, such as silicon.

transition metals

The d-block metals in the center of the table (groups 3–12), including iron, copper, and gold.

noble gases

The Group 18 elements with full valence shells, which makes them extremely unreactive.

The gaps Mendeleev dared to leave

When Mendeleev laid out his table in 1869, he did something genuinely risky — he left blank squares where the pattern demanded an element no one had yet found, and predicted their properties in detail. He described a missing 'eka-silicon' with a specific mass and density; when germanium was discovered in 1886, its measured properties matched his forecast almost exactly, and gallium and scandium confirmed two more predictions. That is the moment the table stopped being a clever chart and became a prediction machine.

Why the table is ordered by protons, not mass

Mendeleev ordered elements by atomic mass, but a few stubborn pairs — like tellurium and iodine — only fell into the right families if he swapped them out of mass order. The puzzle was solved in 1913 when Henry Moseley used X-rays to measure each element's nuclear charge and showed the true ordering is by atomic number (proton count), not mass. Moseley's insight fixed every anomaly at once; tragically he was killed in the First World War at just 27, one of science's great lost talents.

Seven elements that refuse to be alone

Seven elements never exist as single atoms under normal conditions — hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, and iodine all travel as diatomic molecules like O₂ and Cl₂. The table hides other quirks of personality too: of all the elements, only two are liquid at room temperature — bromine and mercury — while every other element is a solid or a gas. Knowing these small exceptions keeps you from writing a lone 'O' when the real substance is O₂.

Finishing the table: the newest elements

The bottom row of the periodic table was completed only recently, with synthetic superheavy elements assembled a few atoms at a time inside particle accelerators. Elements 113, 115, 117, and 118 were officially named in 2016 — including oganesson, number 118, named after a living physicist. These atoms are so unstable they survive for only fractions of a second, yet the table predicted where they would sit and roughly how they should behave long before anyone made them.

The element that won't sit still: hydrogen

Hydrogen is usually parked at the top of Group 1, but it is not truly an alkali metal — it is a colorless gas that happens to share the family's single valence electron. Some chemists float it above the whole table, or place it over the halogens instead, since it is 'one electron short' of a full shell just as they are. Hydrogen's awkward placement is a good reminder that the periodic table is a model — powerful, but not a perfect box for every element.

In the real world

Because position predicts properties, chemists could describe how newly synthesized elements at the bottom of the table should behave before ever making more than a few atoms of them.

Exit check

Locate [Ne]3 s²3p³.

Periodic Trends

Learning goals

• I can describe how atomic radius, ionization energy, and electronegativity change across a period and down a group.
• I can explain these trends using effective nuclear charge, shielding, and distance from the nucleus.
• I can use the trends to compare two elements and predict relative reactivity.

Readings for this unit

• From a spectrum to an atomic model

Because the periodic table is arranged by electron configuration, its properties do not scatter randomly — they slide in smooth, predictable periodic trends. Almost every trend comes down to a tug-of-war between two forces: the nucleus pulling valence electrons inward, and the distance and shielding pushing back. Learn to weigh those two factors and you can predict an element's behavior from its address on the table alone.

The key concept is effective nuclear charge (Zeff) — the net pull a valence electron actually feels once the inner electrons have canceled out part of the nuclear charge by shielding it. Moving left to right across a period, protons are added to the nucleus while the new electrons all crowd into the same energy level, adding almost no extra shielding. So Zeff climbs steadily across a period, gripping the valence electrons ever more tightly. Moving down a group, each step adds a whole new shell of electrons, so the valence electrons sit farther out and are better shielded — the nucleus's grip on them loosens.

Atomic radius follows directly. Across a period, rising Zeff pulls the electron cloud inward, so atoms get smaller left to right — even though you are adding electrons. Down a group, each new energy level puts the outermost electrons farther from the nucleus, so atoms get larger top to bottom. Put together: the biggest atoms sit at the bottom-left of the table (like cesium) and the smallest nonmetal atoms at the top-right (like fluorine).

Ionization energy — the energy needed to yank off the outermost electron — runs as the mirror image of radius, and for the same reason. In a small atom with high Zeff, the outer electron is held close and tight, so it takes a lot of energy to remove: ionization energy increases across a period. In a large atom, the outer electron is far away and well shielded, so it comes off easily: ionization energy decreases down a group. Whenever an electron is held tightly, it is both hard to remove and hard to pull away — so small atoms tend to have high ionization energies.

Electronegativity measures how strongly an atom pulls on the electrons it *shares* in a bond, and it follows the same pattern as ionization energy: it increases across a period and decreases down a group. That makes fluorine, at the top right (noble gases aside), the most electronegative element of all. On the widely used Pauling scale fluorine is set near 4.0, while the reactive metals at the bottom left score below 1.0. Noble gases are usually left out because they rarely bond at all.

These three trends together explain reactivity, which is the bridge to chemical bonding. Metals on the lower left have large radii, low ionization energies, and low electronegativities, so they give up electrons easily — that is why cesium reacts explosively while its group-mate lithium reacts more gently. Nonmetals on the upper right have small radii, high ionization energies, and high electronegativities, so they grab electrons hungrily — which is why fluorine is the most aggressive nonmetal there is. Read an element's position, weigh Zeff against distance, and the periodic table becomes a working prediction machine — exactly the promise of this unit's essential question.

Model check: General trends have exceptions; do not use a trend arrow as an exact numerical table. K has an occupied outer shell farther from the nucleus with more inner shielding, outweighing its greater proton count in this comparison.

Effective Nuclear Charge (simple approximation)

Zeff ≈ (number of protons) − (number of core electrons)

core electrons are the inner-shell (non-valence) electrons that shield the valence electrons; a higher Zeff means the valence electrons are held more tightly, giving a smaller radius and a higher ionization energy.

Worked examples

Avoid these mistakes

• General trends have exceptions; do not use a trend arrow as an exact numerical table.

Words to know

atomic radius

A measure of the size of an atom, roughly the distance from the nucleus to the outermost electrons.

ionization energy

The energy required to remove the most loosely held electron from a neutral atom.

electronegativity

A measure of how strongly an atom attracts the shared electrons in a chemical bond.

effective nuclear charge (Zeff)

The net positive pull actually felt by a valence electron, after inner electrons cancel out some of the nuclear charge.

shielding (screening)

The way inner-shell electrons block, or screen, valence electrons from feeling the full charge of the nucleus.

periodic trend

A predictable pattern in a property that repeats as you move across periods or down groups.

The trend that flips: ionic radius

Atomic radius has a twin that behaves in reverse. When an atom loses electrons to become a positive cation, it shrinks dramatically — often to less than half its former size — because the electrons that remain feel the nucleus's full pull. When an atom gains electrons to become a negative anion, it swells instead. Line up ions that have the same electron count (like O²⁻, F⁻, Na⁺, and Mg²⁺) and the one with the most protons is the smallest, its nucleus reeling the same cloud in tighter.

The trends aren't perfectly smooth

Ionization energy generally climbs across a period, but it hiccups in a couple of revealing spots. Pulling an electron from oxygen is actually a touch easier than from nitrogen, because nitrogen's half-filled 2p sublevel is extra stable while oxygen has one paired electron that repels its partner. These little dips are not flaws in the trend — they are evidence that the stability of half-filled and full sublevels (the same effect that bends chromium's configuration) is a genuine physical force.

Effective nuclear charge, made concrete

Why doesn't sodium's outer electron feel the full +11 of its nucleus? Because the ten inner core electrons stand between them and cancel out most of that pull through shielding. The lone valence electron effectively feels only about +1, which is exactly why it is so loosely held and so easily lost. Effective nuclear charge is just this idea — the net tug that survives after the core electrons take their cut — and it is the engine running underneath nearly every periodic trend.

Why sodium explodes and gold never rusts

These trends are really a map of reactivity. Metals at the lower left have low ionization energies, so they surrender electrons eagerly — drop cesium or sodium into water and it reacts violently. Metals with high ionization energies cling to their electrons instead, which is why gold and platinum can sit in jewelry for centuries without tarnishing. Reactivity is not a grab-bag of facts to memorize but a direct reading of an element's address on the table.

From trends to bonds: predicting how atoms join

The payoff of electronegativity is that the difference between two atoms predicts the kind of bond they will form. A large difference — a hungry nonmetal meeting a generous metal, like sodium and chlorine — means electrons are handed over outright in an ionic bond. A small difference means the electrons are shared in a covalent bond, and between two metals the electrons roam freely in a metallic bond. That single comparison, read straight off the periodic table, is the doorway into the entire next unit on chemical bonding.

In the real world

Battery designers chase metals with low ionization energy (they give up electrons easily), while electronegativity differences let chemists predict whether two elements will form an ionic or a covalent bond.

Exit check

Explain why K is usually larger than Na.

Electron Transitions & Emission Spectra

Learning goals

• I can connect spectral lines to allowed electron energy changes.
• I can calculate photon energy and evaluate a pattern match.

Readings for this unit

• From a spectrum to an atomic model

An electron in an atom has allowed energy states. The lowest available arrangement is the ground state. Absorbing sufficient energy can move an electron to a higher allowed state. An excited state is less stable; a transition to a lower state can release a photon with energy equal to the difference between states.

The energy difference is ΔE = Ehigh − Elow and photon energy is E = hf = hc/λ. Here h = 6.626×10⁻³⁴ J·s and c = 3.00×10⁸ m/s. Higher frequency means higher photon energy; longer wavelength means lower energy. Use wavelength in meters when using these SI constants.

Atomic line spectra contain selected wavelengths rather than every wavelength because allowed energy differences are discrete. Different elements have different state structures and patterns. An unknown can be compared with a reference pattern; one matching line is weaker evidence than several matches within measurement uncertainty.

A line spectrum is evidence for quantized energy, not for electrons literally following painted circular tracks. Orbitals describe probability distributions. Shell diagrams organize electron counts, while energy-level diagrams organize allowed transitions; their distances on a drawing do not directly represent physical orbital distances.

Absorption spectra show missing wavelengths from a continuous source when atoms absorb matching photons. Emission spectra show light released by excited atoms. The same transition energy links the two descriptions, though actual observed intensities also depend on population, temperature, and instrument response.

Use provided spectra or enclosed educational equipment for independent study. Do not burn metal salts, open electrical lamps, or stare at the Sun. For an unidentified star or gas, compare multiple features and state instrument resolution and possible mixtures before claiming an identity.

Worked examples

Avoid these mistakes

• A photon is emitted on a downward transition; excitation requires energy input.

Words to know

ground state

The lowest-energy available electron arrangement.

excited state

An arrangement with energy above the ground state.

photon

A quantum of electromagnetic energy.

In the real world

Use the chapter investigation to connect this unit to measured evidence and a practical decision.

Exit check

State the relation between wavelength and photon energy.