A guided, interactive walkthrough of how atoms are built and how that structure shows up in light, mass, and identity.
0The Ten Units
Chemistry is organized into ten major units. Each one builds on the last. This page covers Unit 1, Atomic Structure, the foundation that every other chemistry topic stands on.
I. Atomic Structure
II. The Periodic Table
III. Stoichiometry
IV. Chemical Bonding
V. Physical Behavior of Matter
VI. Equilibrium
VII. Organic Chemistry
VIII. Oxidation-Reduction
IX. Acids, Bases, and Salts
X. Nuclear Chemistry
By the end of this page you should be able to identify and place every subatomic particle, read isotope notation, calculate the average atomic mass of an element, recognize ground vs excited electron configurations, and explain how bright-line spectra prove that energy levels exist.
1The Three Subatomic Particles
An atom has three building blocks. Two of them sit in a tiny dense ball at the center, the nucleus. The third buzzes around the outside in a region called the electron cloud.
Almost all of an atom's mass is in the nucleus, where protons and neutrons live. Electrons contribute almost nothing to mass. But electrons take up almost all of an atom's volume, because the orbitals are huge compared to the nucleus.
If the nucleus were a marble in the middle of a football field, the electrons would be moving around the outer edge of the field. The atom is mostly empty space.
An amu (atomic mass unit) is a tiny mass unit chosen so that protons and neutrons each weigh almost exactly 1. The whole point of amu is to make atomic mass numbers small and easy to compare.
2Rutherford's Gold Foil Experiment
Before 1911, scientists thought atoms were like a "plum pudding," a soft positive blob with electrons stuck inside. Ernest Rutherford disagreed and ran a famous experiment to test it.
The Setup
Rutherford fired tiny positive particles called alpha particles at a very thin sheet of gold foil. He surrounded the foil with a detector to see where the particles went after passing through.
What He Expected vs What He Saw
If the plum-pudding model were correct, all the alpha particles should pass through with only tiny deflections. They didn't.
What actually happened:
· Most alpha particles passed straight through. The atom must be mostly empty space.
· Some were deflected at small angles.
· A few bounced almost straight back. Rutherford said this was "as if you fired a bullet at tissue paper and it bounced back at you."
Conclusions
1. Atoms have a tiny, dense, positively charged nucleus at the center.
2. Most of the atom is empty space.
3. Electrons orbit the nucleus at a relatively large distance, in this empty space.
Rutherford Simulation: Fire Alpha Particles at the Foil
0
Passed straight
0
Slight deflection
0
Bounced back
Each shot has a small chance of hitting near a nucleus. Most miss entirely.
3Atomic Number, Mass Number, and Isotopes
Atomic number (Z) = number of protons.
The atomic number is the identity of an element. If you change the number of protons, you change the element. Carbon has 6 protons. Anything with 6 protons is carbon, by definition.
Mass number (A) = protons + neutrons.
The mass number counts the heavy stuff in the nucleus. Electrons weigh almost nothing, so they do not count.
In a neutral atom, number of electrons = number of protons. Charge is only nonzero when those don't match.
Isotope Notation
Isotopes are written with the mass number on the upper left and the atomic number on the lower left, like $^{14}_{6}\text{C}$. The chemical symbol on the right tells you the element, and the numbers tell you the specific version.
Isotope Notation Builder
126C
What is an Isotope?
Isotopes are atoms of the same element with different masses, because they have different numbers of neutrons. Same protons, different neutrons.
Carbon has three notable isotopes:
· Carbon-12 ($^{12}_{6}$C): 6 protons, 6 neutrons. Most common.
· Carbon-13 ($^{13}_{6}$C): 6 protons, 7 neutrons. Stable, rare.
· Carbon-14 ($^{14}_{6}$C): 6 protons, 8 neutrons. Radioactive. Used in carbon dating.
All three are still carbon. They differ only in the nucleus.
4Average Atomic Mass
When you look up an element on a periodic table, you see a decimal mass like Cl = 35.45 amu. That is not a typo, and chlorine is not made of half-atoms. The decimal is a weighted average of all the naturally occurring isotopes of that element.
Atomic mass is the weighted average mass of the naturally occurring isotopes of an element.
Why "weighted"? Because rarer isotopes count less. The formula:
Atomic mass = (mass$_1$ × abundance$_1$) + (mass$_2$ × abundance$_2$) + ...
where each abundance is a decimal between 0 and 1, and they all add up to 1.
Worked Example: Chlorine
Isotopes. Cl-35 (mass 35) is 75.77% of natural chlorine. Cl-37 (mass 37) is 24.23%.
Total. $26.52 + 8.97 \approx 35.49$ amu. The periodic table lists 35.45, basically the same with sharper data.
Sanity check. The average is closer to 35 than to 37, because Cl-35 is more abundant. The weighted average always lands closer to the more common isotope.
Weighted Average Calculator
5Ions: When the Numbers Don't Match
A neutral atom has the same number of protons and electrons. When those counts differ, the atom has a net charge and is called an ion.
charge = (number of protons) − (number of electrons)
More electrons than protons means a negative charge (the extra electrons each carry -1). Fewer electrons than protons means a positive charge.
Example. An atom has 11 protons and 10 electrons. What is its charge and identity?
Identity. 11 protons means it is sodium (Na).
Charge. $11 - 10 = +1$. So it is Na$^+$, a sodium ion (cation).
Atom Builder: Build Any Atom or Ion
Protons
Neutrons
Electrons
Element
Mass #
Atomic #
Charge
6Energy Levels: Ground State and Excited State
Electrons are arranged into principal energy levels, sometimes called shells. They are numbered 1, 2, 3, and so on, with 1 being closest to the nucleus and lowest in energy. Each shell holds a limited number of electrons.
Shell 1 (n=1)Holds up to 2 electrons. Closest to the nucleus, lowest in energy.Shell 2 (n=2)Holds up to 8 electrons.Shell 3 (n=3)Holds up to 18 electrons (8 in the introductory model).Shell 4 (n=4)Holds up to 32 electrons.Valence shellThe outermost shell that contains electrons. These electrons have the most energy and do all the bonding.
Ground state: every electron is in the lowest available energy level. This is the resting, stable arrangement.
Excited state: at least one electron has absorbed energy and jumped to a higher level, leaving a gap below it. The atom is unstable and will fall back.
Carbon ground state: 2-4 (2 in shell 1, 4 in shell 2).
Carbon excited state example: 2-3-1 (one electron jumped from shell 2 to shell 3).
When you see a configuration written as a series of numbers separated by dashes, like 2-8-2 for magnesium, that is the shell configuration. Add the numbers up to get the total electron count, and check it against the atomic number to make sure it is the right element. To check ground vs excited, compare against what you would get by filling shells in order with the proper caps.
Ground vs Excited: Spot the Difference
Ground
Boron (Z = 5)
2-3
Shell 1 is full (2). Shell 2 has 3, the rest of B's 5 electrons. No higher shells used. Ground state.
Excited
Boron (Z = 5)
2-2-1
Total is still 5, but an electron sits in shell 3 while shell 2 is not full. That promoted electron is the giveaway.
Ground
Sodium (Z = 11)
2-8-1
Shells 1 and 2 are full. Shell 3 has the leftover 1 electron. All electrons are as low as they can be.
Excited
Sodium (Z = 11)
2-7-2
Total is still 11, but shell 2 is missing one electron and shell 3 has an extra. An electron is excited.
7Light Emission and Bright-Line Spectra
When an excited electron falls back down to a lower energy level, the atom releases the extra energy as a photon, a packet of light. The energy of the photon equals the difference between the two levels.
Light is produced when an atom transitions from the excited state back to the ground state.
The exact wavelength (color) of the light depends on the size of the energy drop. Bigger drop → more energetic photon → bluer/violet color. Smaller drop → less energetic photon → redder color.
Bohr Model: Watch an Electron Jump Down
Click a button to see an electron drop and emit a photon. The bigger the drop, the more energetic the light.
Bright-Line Spectra
If you heat up a gas of one element and pass the emitted light through a prism, you do not get a rainbow. You get a few sharp bright lines on a black background. That is called a bright-line spectrum, and every element has its own unique pattern.
Why? Because the energy levels of an atom are quantized, meaning only certain drops are possible. Each allowed drop produces a photon of a specific wavelength. Different elements have different shells with different gaps, so they emit different colors. The pattern is a fingerprint.
Sample Bright-Line Spectra (visible range)
Hydrogen
Helium
Sodium
Mercury
400 nm (violet)500 nm600 nm700 nm (red)
Each white tick marks an emission line for that element. Notice how each pattern is unique. That is how astronomers identify the elements in distant stars.
8The Wave-Mechanical Model and Orbitals
Bohr's model of fixed circular orbits was a great first step, but it was not the whole story. Modern chemistry uses the wave-mechanical model (also called the quantum-mechanical model). In this picture, electrons do not travel on neat circles. Instead, they exist as fuzzy regions of probability called orbitals.
An orbital is a region in space where an electron is most likely to be found. Not where it is for sure. Where it is most probably.
The wave-mechanical model still has principal energy levels (shells, n = 1, 2, 3, ...), and the electrons inside still have similar total energies as in Bohr's picture. The difference is the shape and the language of probability instead of certainty.
Each orbital has a shape. The s orbital is a sphere. The p orbital is a dumbbell. The d and f orbitals are more complex. The shape tells you the regions of high probability for finding the electron.
Orbital Shapes: Where the Electron Probably Is
s orbital
Spherical, surrounding the nucleus.
p orbital
Dumbbell, two lobes through the nucleus.
d orbital
Four-lobed cloverleaf around the nucleus.
f orbital
Complex, multi-lobed shape.
Bohr vs Wave-Mechanical: Side by Side
Bohr Model
Electrons travel on fixed circular orbits at definite distances. A useful first picture, but not the full reality.
Wave-Mechanical Model
Electrons exist in orbitals, fuzzy clouds of probability. The dots show where the electron is more likely to be found over many measurements.
9Flashcards: Self-Quiz Yourself
Click the card to flip. Use the arrows to move through the deck. Shuffle whenever you want a fresh order.
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10Practice Problems
Try each one before opening the answer.
1. An atom has 17 protons, 18 neutrons, and 17 electrons. What is its identity, mass number, and charge?
Identity: 17 protons = chlorine (Cl). Mass number = 17 + 18 = 35. Charge: protons - electrons = 17 - 17 = 0. So this is a neutral Cl-35 atom.
2. Naturally occurring copper is 69.17% Cu-63 (mass 62.93 amu) and 30.83% Cu-65 (mass 64.93 amu). Calculate the average atomic mass.
$(62.93 \times 0.6917) + (64.93 \times 0.3083) = 43.53 + 20.02 \approx $ 63.55 amu. The periodic table value is 63.55, matching exactly.
3. Which of these electron configurations represents an excited state of a sodium atom (Z = 11)?
(A) 2-8-1 (B) 2-7-2 (C) 2-8-2 (D) 2-8
(B) 2-7-2. Sodium has 11 electrons, so configurations summing to 11 are valid. (A) 2-8-1 is the ground state. (B) 2-7-2 sums to 11 but shell 2 is not full while shell 3 has electrons, so it is excited. (C) sums to 12, not sodium. (D) sums to 10, not sodium.
4. An ion has 12 protons and 10 electrons. What is its symbol and charge?
5. What did Rutherford's gold foil experiment demonstrate about the structure of the atom? Pick the best answer.
(A) Electrons exist in shells.
(B) The atom has a small, dense, positively charged nucleus.
(C) Atoms cannot be divided.
(D) Electrons emit photons when they fall energy levels.
(B). The deflection of alpha particles, especially the few that bounced back, told Rutherford the atom must have a tiny, dense, positively charged center.
6. Why does each element produce a unique bright-line spectrum?
Each element has its own unique set of energy levels. When an excited electron drops between levels, the photon released has an energy equal to the gap. Different gaps mean different photon energies, which means different wavelengths (colors). The full pattern of allowed transitions is unique to each element, so it acts as a fingerprint.
7. $^{40}_{20}\text{Ca}$ — how many protons, neutrons, and electrons are in this neutral atom?
Atomic number 20: 20 protons. Mass number 40: protons + neutrons = 40, so 20 neutrons. Neutral atom: electrons = protons = 20 electrons.
8. Which subatomic particle has the smallest mass?
The electron. Mass roughly 1/1836 amu, almost negligible compared to a proton or neutron at ~1 amu.
9. What is the difference between Bohr's model and the wave-mechanical (quantum) model of the atom?
Bohr's model places electrons on fixed circular orbits at definite distances. The wave-mechanical model treats electrons as probability clouds called orbitals. The orbital represents where the electron is most probably found, not exactly where it is. Both keep the idea of energy levels.
10. Bromine occurs as Br-79 (50.69%, mass 78.92) and Br-81 (49.31%, mass 80.92). Predict whether the average atomic mass is closer to 79 or 81, then calculate.
Both isotopes are nearly equal in abundance, so the average should sit roughly in the middle, slightly closer to 79 (since Br-79 is a bit more abundant). $(78.92 \times 0.5069) + (80.92 \times 0.4931) = 40.01 + 39.91 \approx $ 79.92 amu. Periodic table lists 79.90.
11Quick Review Cheat Sheet
ProtonCharge +1, mass ~1 amu, in nucleus.NeutronCharge 0, mass ~1 amu, in nucleus.ElectronCharge -1, mass ~1/1836 amu, in orbitals.Atomic # (Z)= number of protons. Identifies the element.Mass # (A)= protons + neutrons.Charge= protons − electrons.Atomic massWeighted average of all natural isotopes.IsotopeSame element, different number of neutrons.RutherfordGold foil → tiny dense positive nucleus, mostly empty atom.Ground stateAll electrons in lowest available levels.Excited stateAn electron has jumped to a higher level. Total still matches Z, but a lower shell is unfilled.Light emissionExcited electron falls back → releases a photon.Bright-line spectrumUnique pattern of emission lines for each element.Valence shellOutermost shell with electrons. Highest energy. Does the bonding.Bohr modelElectrons in fixed circular orbits.Wave-mechanical modelElectrons in orbitals = regions of probability.OrbitalMost probable region for finding an electron.