1. The building blocks of matter
Matter is anything with mass and volume. Atoms are the smallest unit of an element that still retains the element's chemical properties. An element is a pure substance made of only one type of atom; a compound is two or more elements chemically bonded in fixed proportions; a mixture is two or more substances physically combined and separable by physical means.
Quick distinction
Pure substances (elements and compounds) have a fixed composition and cannot be separated by physical means. Mixtures (homogeneous like salt water, or heterogeneous like sand and gravel) can be separated by physical means like filtration, distillation, or evaporation.
2. Subatomic particles
Atoms are made of three particles. Two live in the nucleus (which holds nearly all the mass), one orbits in the electron cloud.
| Particle | Symbol | Charge | Relative Mass | Location |
|---|---|---|---|---|
| Proton | p+ or 1₁p | +1 | 1 amu | Nucleus |
| Neutron | n0 or 1₀n | 0 (neutral) | 1 amu | Nucleus |
| Electron | e− or 0₋₁e | −1 | 1/1836 amu (≈ 0) | Electron cloud |
Three facts the Regents drills repeatedly
- The nucleus is positively charged and contains almost all the mass.
- A neutral atom has equal numbers of protons and electrons.
- An electron has approximately zero mass compared to a proton or neutron.
3. Atomic number and mass number
Every element is uniquely identified by its atomic number (Z) = the number of protons in the nucleus. The atomic number defines the element. Change the number of protons and you have a different element entirely.
Mass Number (A) = number of protons + number of neutrons
Number of Neutrons = A − Z
On the Periodic Table, the atomic number is the whole number (always an integer). The atomic mass shown is a decimal — it's the weighted average of all naturally occurring isotopes of that element.
Reading an isotope symbol
The notation 14₆C means carbon with mass number 14 and atomic number 6. That's 6 protons, 6 electrons (if neutral), and 14 − 6 = 8 neutrons. The mass number is the superscript on the left, the atomic number is the subscript on the left.
4. Isotopes
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons, and therefore different mass numbers. They are chemically identical (same electron configuration) but physically slightly different (different masses, some are radioactive).
| Isotope | Protons | Neutrons | Mass Number |
|---|---|---|---|
| Carbon-12 | 6 | 6 | 12 |
| Carbon-13 | 6 | 7 | 13 |
| Carbon-14 | 6 | 8 | 14 (radioactive) |
5. Calculating average atomic mass
The number on the Periodic Table is a weighted average based on the natural abundance of each isotope. To calculate it:
Worked example
Chlorine has two natural isotopes:
- Cl-35 with mass 34.97 amu, abundance 75.78%
- Cl-37 with mass 36.97 amu, abundance 24.22%
Avg = (34.97 × 0.7578) + (36.97 × 0.2422) = 26.50 + 8.95 = 35.45 amu. Notice this matches the Periodic Table value.
6. The evolution of atomic models
The Regents will ask you to put these in order, identify which scientist proposed what, and describe the key experiment that changed each model.
| Year | Scientist | Model | Key idea / experiment |
|---|---|---|---|
| ~400 BC | Democritus | "Atomos" | Matter is made of indivisible particles. Philosophical, not experimental. |
| 1803 | Dalton | Solid sphere ("billiard ball") | Atoms are indivisible; all atoms of an element are identical; compounds form in fixed ratios. |
| 1897 | Thomson | "Plum pudding" | Discovered the electron with the cathode-ray tube. Proposed atoms are positive spheres with electrons embedded like raisins in pudding. |
| 1911 | Rutherford | Nuclear model | Gold foil experiment: most alpha particles passed through, but a few deflected sharply. Conclusion: atoms are mostly empty space with a small, dense, positively charged nucleus. |
| 1913 | Bohr | Planetary / fixed orbits | Electrons travel in fixed circular orbits at specific energy levels. Electrons absorb energy to jump to higher levels and release it as light when falling back. |
| 1920s+ | Schrödinger, Heisenberg | Wave-mechanical (current model) | Electrons occupy regions of probability called orbitals, not fixed paths. We can describe where an electron is likely to be, but not exactly where. |
The two questions the Regents always asks about models
- What did Rutherford's gold foil experiment prove? That the atom is mostly empty space with a small, dense, positively charged nucleus.
- What's the difference between Bohr and the wave-mechanical model? Bohr said electrons travel in fixed circular orbits. The current model says electrons exist in regions of probability (orbitals).
7. Electron arrangement
Electrons are arranged in energy levels (also called shells) around the nucleus. Each level can hold a maximum number of electrons. For Regents, you write the electron configuration as numbers separated by dashes for each shell from inside out.
The configuration is given directly on the Periodic Table. For example, sodium (Na, Z = 11) is listed as 2-8-1: two electrons in the first shell, eight in the second, one in the third (its single valence electron).
Valence electrons
Valence electrons are the electrons in the outermost (highest) energy level. They determine how an atom bonds. Group 1 elements have 1 valence electron, Group 2 have 2, Group 13 have 3, Group 14 have 4, and so on up to Group 18 noble gases with a full 8 (except helium with 2). Valence electrons are the entire reason for the octet rule and chemical bonding.
Ground state vs. excited state
An atom is in the ground state when its electrons are in the lowest available energy levels — the configuration listed on the Periodic Table.
If an electron absorbs energy (heat, light, electricity), it can jump from a lower energy level to a higher one. The atom is now in an excited state. Its electron configuration will not match the one on the Periodic Table.
Spotting excited state on the exam
Sodium ground state: 2-8-1. If you see Na written as 2-7-2 or 2-8-0-1, that's an excited state — an electron has been bumped up to a higher level. The total number of electrons must still equal the atomic number (here, 11).
8. Bright-line spectra
| Formula | Variables | Units | Use it for |
|---|---|---|---|
| Ephoton = Ehigh − Elow | E = energy of emitted photon | J (or eV) | Why each line has a fixed wavelength |
| E = hf | h = Planck's constant; f = frequency | h = 6.626 × 10⁻³⁴ J·s; f in Hz | Convert between photon energy and color |
| c = λf | c = speed of light; λ = wavelength | c = 3.00 × 10⁸ m/s; λ in m | Convert between wavelength and frequency |
| Bigger ΔE → bluer / higher f / shorter λ | — | — | Ranking lines by energy at a glance |
When an excited electron drops back down to a lower energy level, it releases the absorbed energy as a photon of light. The wavelength (color) depends on the size of the energy drop.
Because each element has a unique set of energy levels, each element emits a unique pattern of bright lines — its bright-line spectrum. This is why we can identify elements in distant stars: we read their spectra.
Common Regents question
You'll be shown bright-line spectra for several elements plus a "mixture." You identify which elements are in the mixture by matching its lines to the individual element spectra. The answer is: every element whose lines appear in the mixture is present.
The bright-line spectrum is direct experimental evidence that electrons exist in discrete energy levels, not a continuous range. If energy levels were continuous, you'd see a continuous rainbow of light from any excited atom — not a few sharp lines.
9. Ions: cations and anions
Atoms become ions when they gain or lose electrons. Protons never change in normal chemistry — they only change in nuclear reactions.
| Ion type | What happens | Charge | Typical of |
|---|---|---|---|
| Cation | Loses electrons | Positive (+) | Metals (left side) |
| Anion | Gains electrons | Negative (−) | Nonmetals (right side, except noble gases) |
Calculating ion charge
A neutral atom has equal protons and electrons. The charge on an ion equals
protons − electrons.
Example: Mg²⁺ has 12 protons (the atomic number doesn't change) and 10 electrons.
Charge = 12 − 10 = +2. ✓
Example: O²⁻ has 8 protons and 10 electrons. Charge = 8 − 10 = −2. ✓
Atoms ionize to achieve a stable noble gas configuration — a full octet of valence electrons (or 2 for the smallest atoms). Sodium (2-8-1) loses 1 electron to become Na⁺ (2-8), matching neon's configuration. Chlorine (2-8-7) gains 1 electron to become Cl⁻ (2-8-8), matching argon's. This drive for stability is the whole point of bonding (Topic 4).
Key terms to know cold
| Atom | Smallest unit of an element retaining its properties |
| Element | Pure substance made of one type of atom |
| Nucleus | Dense, positive center; contains protons + neutrons |
| Proton | +1 charge, 1 amu, in nucleus; defines the element |
| Neutron | 0 charge, 1 amu, in nucleus; varies in isotopes |
| Electron | −1 charge, ~0 mass, in cloud/orbitals |
| Atomic number (Z) | Number of protons; identifies the element |
| Mass number (A) | Protons + neutrons |
| Isotope | Same protons, different neutrons |
| Atomic mass | Weighted average of all natural isotopes (decimal on PT) |
| Orbital | Region of probability where an electron is likely found (wave-mech model) |
| Valence electrons | Electrons in the outermost energy level |
| Ground state | Electrons in their lowest possible energy levels |
| Excited state | An electron has absorbed energy and jumped to a higher level |
| Bright-line spectrum | Pattern of light emitted when excited electrons drop back down; unique to each element |
| Ion | Atom that has gained or lost electrons; has a net charge |
| Cation | Positive ion (lost electrons); typical of metals |
| Anion | Negative ion (gained electrons); typical of nonmetals |
Practice questions
Q1 · Multiple Choice (Part A style)
Which subatomic particles are located in the nucleus of an atom?
- protons and electrons
- protons and neutrons
- neutrons and electrons
- only electrons
Show answer
(2) protons and neutrons. Electrons orbit the nucleus in the electron cloud; they are not inside it.
Q2 · Multiple Choice
An atom of nitrogen-15 contains how many neutrons?
- 7
- 8
- 15
- 22
Show answer
(2) 8. Nitrogen has atomic number 7 (7 protons). Mass number is 15, so neutrons = 15 − 7 = 8.
Q3 · Multiple Choice
Which electron configuration represents a sulfur atom in an excited state?
- 2-6
- 2-8-6
- 2-7-7
- 2-8-8
Show answer
(3) 2-7-7. Sulfur has atomic number 16, so a sulfur atom has 16 total electrons. Configuration (2) shows the ground state on the Periodic Table. Configuration (3) also totals 16 electrons but the second shell is missing an electron (only 7 instead of 8) and the third shell has an extra one — that electron has been excited to a higher level. (1) only totals 8 electrons (wrong atom) and (4) totals 18 (wrong atom or an ion, but the question says "atom").
Q4 · Part B-2 / Short response
The element copper has two naturally occurring isotopes: Cu-63 (mass 62.93 amu, abundance 69.17%) and Cu-65 (mass 64.93 amu, abundance 30.83%). Calculate the average atomic mass of copper. Show your work.
Show answer
Avg = (62.93 × 0.6917) + (64.93 × 0.3083)
= 43.53 + 20.02
= 63.55 amu
(This matches the value 63.55 shown on the Periodic Table — confirming the calculation.)
Q5 · Part C / Extended response
In Rutherford's gold foil experiment, alpha particles were directed at a thin sheet of gold foil. Most particles passed straight through, but a small number were deflected at large angles, and a few even bounced backward.
(a) State the conclusion Rutherford drew about the structure of the atom based on the deflected particles. (b) Why did most of the alpha particles pass straight through?
Show answer
(a) The atom contains a small, dense, positively charged center (the
nucleus). The deflected alpha particles were repelled by this concentrated positive charge.
(b) Most particles passed through because the atom is mostly empty space.
The nucleus occupies only a tiny fraction of the atom's total volume.