This unit builds the modern picture of the atom, from the experiments that revealed its parts to the electron arrangements that explain the shape of the periodic table. By the end you should be able to count subatomic particles, calculate average atomic mass, convert between grams, moles and particles, write electron configurations, and predict periodic trends.
Read the overview first, then work through each topic's notes and try the checkpoint question before you look at the answer.
Unit 3 answers two big questions. What is an atom made of, and why do elements in the same column of the periodic table act alike? The math here, especially mole conversions, shows up again in every later unit.
Honors Chemistry Unit 3 covers the history of atomic theory (Dalton, Thomson, Rutherford, Bohr and the quantum mechanical model), the structure of the atom (protons, neutrons, electrons, atomic number, mass number, isotopes and ions), average atomic mass and the mole (Avogadro's number, molar mass and gram-mole-particle conversions), electron configurations (aufbau principle, Pauli exclusion principle, Hund's rule, orbital diagrams and noble gas shorthand), and periodic trends (atomic radius, ionization energy, electronegativity and ion size). Schools order these topics differently, and some save empirical formulas or electron configurations for a later unit, but this set is typical for a first-year honors course.
Dalton's solid spheres could not explain Thomson's electrons, and Thomson's plum pudding could not explain Rutherford's bouncing alpha particles. Bohr's orbits explained the spectra of hydrogen and other one-electron species but did not accurately predict the spectra of multi-electron atoms, so the quantum mechanical model took over and is still the current model. Tests often ask which observation led to which model.
One mole is particles. One mole of atoms of an element has a mass in grams numerically equal to its average atomic mass in amu. That means 12.01 g of carbon contains about carbon atoms. Molar mass is the conversion factor between grams and moles.
The s, p, d and f blocks of the table match the sublevel being filled. Main-group elements in the same group usually have similar valence-electron configurations, so they react in similar ways. Helium is an exception to the simple count because its only occupied shell is full with two electrons. Sodium and potassium each have one valence electron, and both react strongly with water.
Change the proton count and you have a different element. Change the neutron count and you get an isotope with a different mass but nearly identical chemical behavior, although isotope mass can affect reaction rates and some physical properties. The way electrons fill energy levels, and how strongly the nucleus pulls on the outermost ones, explains the main periodic trends in this unit.
Atomic theory is the story of evidence forcing scientists to redraw the atom.
An atom has a nucleus containing protons and usually neutrons, with electrons occupying orbitals around it.
The atomic mass on the periodic table is a weighted average of all of an element's natural isotopes.
An electron configuration lists which sublevels an atom's electrons occupy.
Most periodic trends come from nuclear charge, shielding and how far valence electrons sit from the nucleus.
Atomic theory is the story of evidence forcing scientists to redraw the atom. Democritus proposed indivisible atomos around 400 BCE, but with no experiments behind it, the idea sat mostly unused for about 2,000 years.
Lavoisier's law of conservation of mass and Proust's law of definite proportions gave John Dalton the evidence for his 1803 theory. Dalton said elements are made of tiny indivisible atoms, all atoms of one element are identical, atoms combine in whole-number ratios, and reactions rearrange atoms without creating or destroying them.
In 1897 J.J. Thomson used cathode ray tubes to find the electron, a particle about 1/1836 the mass of a hydrogen atom. Atoms could be divided after all. His plum pudding model scattered electrons through a ball of positive charge.
In 1909 Hans Geiger and Ernest Marsden, working under Ernest Rutherford, fired alpha particles at thin gold foil. Most passed through, but a few deflected at large angles. In 1911 Rutherford concluded that positive charge and nearly all the mass sit in a tiny nucleus, and most of the atom is empty space.
Bohr (1913) placed electrons in fixed energy levels to explain hydrogen's line spectrum. The quantum mechanical model developed in the 1920s through work by de Broglie, Heisenberg, Schrödinger, Born and others. It replaced fixed paths with orbitals, which describe where an electron is likely to be found.
Check yourself. Why did the gold foil results rule out the plum pudding model? Answer. Plum pudding spreads positive charge thinly through the whole atom, so nothing should be strong enough to bounce an alpha particle backward. The few large deflections meant the positive charge and mass were packed into a tiny nucleus.
| Model | Year | Picture of the atom | Evidence or reason |
|---|---|---|---|
| Dalton | 1803 | Solid, indivisible sphere | Conservation of mass and definite proportions |
| Thomson | 1904 | Electrons in a ball of positive charge | Cathode rays bent toward positive charge |
| Rutherford | 1911 | Dense positive nucleus, mostly empty space | Gold foil alpha particle scattering |
| Bohr | 1913 | Electrons in fixed energy levels | Hydrogen's line spectrum |
| Quantum mechanical | 1920s | Electrons in probability orbitals | Wave behavior of electrons |
An atom has a nucleus containing protons and usually neutrons, with electrons occupying orbitals around it. Hydrogen-1 is the exception because its nucleus has no neutron. The nucleus holds almost all of the mass in a tiny fraction of the volume.
The atomic number (Z) is the number of protons, and it defines the element. Every atom with 6 protons is carbon. The mass number (A) is protons plus neutrons, so neutrons = A − Z.
Atoms of the same element with different neutron counts are isotopes. Carbon-12 and carbon-14 both have 6 protons, but carbon-12 has 6 neutrons and carbon-14 has 8. Isotopes usually have nearly identical chemical behavior because they have the same electron configuration, although their different masses can affect reaction rates and some physical properties.
A neutral atom has equal numbers of protons and electrons. An ion has gained or lost electrons, never protons. Losing electrons makes a positive cation, and gaining them makes a negative anion.
Worked example. Find the protons, neutrons and electrons in . Iron's atomic number is 26, so it has 26 protons. Neutrons are . The 3+ charge means three electrons were lost, so electrons are .
For anions, add electrons instead. has 17 protons, neutrons, and electrons.
Check yourself. How many protons, neutrons and electrons are in ? Answer. It has 16 protons, 16 neutrons, and 18 electrons, because the 2− charge means two extra electrons.
| Particle | Charge | Approximate mass (amu) | Location |
|---|---|---|---|
| Proton | +1 | 1 | Nucleus |
| Neutron | 0 | 1 | Nucleus |
| Electron | −1 | 0.00055 (about 1/1836 of a proton) | Outside the nucleus, in orbitals |
The atomic mass on the periodic table is a weighted average of all of an element's natural isotopes. Isotopes that are more common count more toward the average. That's why chlorine is listed as 35.45 amu even though no single chlorine atom has that mass.
Worked example. Chlorine is 75.78% chlorine-35 (34.969 amu) and 24.22% chlorine-37 (36.966 amu). Convert the percents to decimals, multiply each by its mass, and add.
The answer lands closer to 35 because chlorine-35 is more common. Always run that sanity check.
A mole is a counting unit like a dozen. One mole equals particles (Avogadro's number). Molar mass is the mass of one mole in g/mol, found by adding atomic masses from the formula. For H₂O it is g/mol.
Worked example. How many molecules are in 10.0 g of water? First convert grams to moles: . Then convert moles to molecules: molecules.
You can write a grams-to-particles conversion in one line, but moles are still the bridge. Use molar mass to convert grams to moles, then use Avogadro's number to convert moles to particles.
Check yourself. What is the mass of 2.50 mol of CO₂? Answer. The molar mass is 12.01 + 2(16.00) = 44.01 g/mol, so 2.50 mol × 44.01 g/mol = 110. g.
An electron configuration lists which sublevels an atom's electrons occupy. Three rules decide the arrangement.
The aufbau principle says electrons fill the lowest-energy sublevels first. The usual filling order is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. For neutral potassium and calcium, 4s is occupied before 3d. Orbital energies shift as electrons are added, so transition-metal ions generally lose 4s electrons before 3d electrons. Sublevels hold s = 2, p = 6, d = 10 and f = 14 electrons.
The Pauli exclusion principle says one orbital holds at most two electrons, and they must have opposite spins. Hund's rule says electrons spread out one per orbital, with the same spin, before any of them pair up. So nitrogen's three 2p electrons sit in three separate orbitals.
The periodic table maps this filling order. Groups 1-2 and helium form the s-block, groups 13-18 except helium form the p-block, the transition metals form the d-block, and the two detached bottom rows form the f-block.
Worked example. Iron has 26 electrons. Fill in order to get (2 + 2 + 6 + 2 + 6 + 2 + 6 = 26). In noble gas shorthand, argon covers the first 18, giving [Ar] . For Fe²⁺, remove the two 4s electrons first, giving [Ar] .
Chromium and copper are common exceptions, written as [Ar] and [Ar] .
Check yourself. Write the full electron configuration of sulfur (16 electrons) and say how many unpaired electrons it has. Answer. Sulfur is . The four 3p electrons fill three orbitals as one pair and two singles, so sulfur has 2 unpaired electrons.
Most periodic trends come from nuclear charge, shielding and how far valence electrons sit from the nucleus. Nuclear charge is how many protons pull on the electrons. Shielding is how much inner electrons block that pull from reaching the outer electrons.
Moving left to right across a period, nuclear charge increases. Added electrons do not fully shield one another from the added protons, so effective nuclear attraction generally increases. In the main-group sections, electrons enter the same outer energy level; in a transition series, they mainly enter an inner d sublevel. So atomic radius generally decreases across a period, while first ionization energy (the energy to remove an outer electron) and electronegativity (pull on shared electrons in a bond) generally increase. Sodium is bigger than chlorine.
Moving down a group, each row adds a new energy level, so valence electrons sit farther out behind more shielding. Radius increases while ionization energy and electronegativity decrease. Cesium is far larger than lithium and loses its outer electron much more easily. Fluorine, at 3.98, is the most electronegative element.
Cations are smaller than their atoms, often because they lose a whole outer shell, so Na⁺ is much smaller than Na. Anions are larger because extra electrons repel each other, so Cl⁻ is larger than Cl.
Many metals on the left lose electrons and form cations. Many nonmetals on the right can gain electrons and form anions, although nonmetals also commonly share electrons in covalent bonds, and noble gases are generally unreactive.
Check yourself. Which has the higher first ionization energy, magnesium or chlorine, and why? Answer. Chlorine does, because it is farther right in the same period, so its greater nuclear charge holds the outer electrons more tightly with about the same shielding.
| Trend | Across a period (left to right) | Down a group |
|---|---|---|
| Atomic radius | Decreases | Increases |
| Ionization energy | Increases | Decreases |
| Electronegativity | Increases | Decreases |
| Metallic character | Decreases | Increases |
| Term | Definition |
|---|---|
| Atomic number | The number of protons in an atom's nucleus, which determines which element it is. |
| Mass number | The total number of protons and neutrons in the nucleus of a specific isotope. |
| Isotope | One of two or more forms of an element that have the same number of protons but different numbers of neutrons. |
| Average atomic mass | The abundance-weighted average mass of an element's natural isotopes, shown on the periodic table in amu. |
| Mole | A counting unit equal to $6.022 \times 10^{23}$ particles of a substance. |
| Molar mass | The mass of one mole of a substance in grams per mole, numerically equal to its atomic or formula mass. |
| Orbital | A region of space around the nucleus where up to two electrons are likely to be found. |
| Valence electrons | For main-group elements, valence electrons are the electrons in the highest occupied energy level, and more generally they are the electrons available for bonding, including some d electrons in transition metals. |
| Ionization energy | The energy required to remove the outermost electron from a gaseous atom. |
| Electronegativity | The tendency of an atom to attract shared electrons toward itself in a chemical bond. |
Mass number is a whole-number count of protons and neutrons in one specific isotope, like 35 for chlorine-35. Atomic mass is the decimal weighted average on the periodic table, like 35.45. Use the mass number to count neutrons, and use the atomic mass for mole and gram calculations.
If you just average 35 and 37 for chlorine you get 36, which is wrong. Isotopes count in proportion to how common they are, so multiply each mass by its decimal abundance before adding. Then check that your answer sits closest to the most abundant isotope.
Adding protons and electrons across a period generally shrinks the atom. The added electrons do not fully shield one another from the increasing nuclear charge, so the nucleus pulls the electron cloud inward more strongly. In main-group sections, electrons enter the same outer energy level; in transition series, they mainly enter an inner d sublevel. Atomic size increases sharply when a new principal energy level begins in the next period.
4s fills before 3d, so it feels natural to remove 3d electrons first. Once electrons are in place, though, the 4s electrons are the outermost and leave first. Fe is [Ar] , so Fe²⁺ is [Ar] , not [Ar] .