AP Chemistry Unit 1: Atomic Structure & Properties
Quick-reference study sheet
This unit delves into the structure and properties of atoms. Matter, which makes up everything around us, is anything with mass that takes up space. Atoms are the basic building blocks of matter and the smallest units that still have the properties of an element. When atoms join together, they create molecules, which form the foundation of chemical compounds.
1. Moles & Molar Mass
The mole is the unit of measurement in the International System of Units (SI) for amount of substance. It is defined as the amount of a chemical substance that contains as many elementary entities (e.g., atoms, molecules, ions, electrons, or photons). This number is expressed by the Avogadro constant, which has a value of 6.022 × 1023. The mole is one of the base units of the SI, and has the unit symbol mol. Molar mass (M) is the mass of one mole (n) of a substance in grams; it is numerically equal to the atomic mass (in amu) listed on the periodic table.
Avogadro’s Number
NA = 6.022 × 1023 particles/mol
Moles from Mass
n = mM
Mass
m = n × M
Particles
N = n × NA
Where:
- n = moles (mol)
- m = mass (g)
- M = molar mass (g/mol)
- N = number of particles, atoms, molecules, formula units
- NA = Avogadro’s number (6.022 × 1023/mol)
The Mole Map
Example: How many moles are in 36.0 g of H2O? (M = 18.0 g/mol)
- Use n = m / M
- n = 36.0 g18.0 g/mol
- Answer: 2.00 mol H2O
2. Mass Spectra of Elements
A mass spectrometer is an analytical laboratory instrument that measures the mass-to-charge ratio of gas-phase ions. Scientists use it to identify unknown chemical compounds, determine isotopic abundances, quantify known substances, and elucidate the chemical structures of complex biological molecules.
Reading a Spectrum
- Peak position (x-axis) → mass of that isotope
- Peak height (y-axis) → relative abundance (%) of that isotope
- Number of peaks → number of naturally occurring isotopes
Average Atomic Mass
avg. mass = Σ (isotope mass × fractional abundance)
Example: Chlorine has two isotopes — 35Cl (34.97 amu, 75.77%) and 37Cl (36.97 amu, 24.23%). Find the average atomic mass.
- (34.97 × 0.7577) = 26.50
- (36.97 × 0.2423) = 8.96
- Add: 26.50 + 8.96 = 35.45 amu (matches the periodic table value for Cl)
3. Elemental Composition of Pure Substances
Percent Composition
% element = mass of element (in 1 mol compound)molar mass of compound × 100%
Empirical vs. Molecular Formula
| Empirical Formula | Molecular Formula |
|---|---|
| Simplest whole-number ratio of atoms in a compound. e.g. CH2O | Actual number of atoms in one molecule; a whole-number multiple of the empirical formula. e.g. C6H12O6 |
Example: A compound is 40.0% C, 6.7% H, and 53.3% O by mass. Find the empirical formula.
- Assume a 100 g sample: 40.0 g C, 6.7 g H, 53.3 g O
- Convert to moles: C = 3.33 mol, H = 6.7 mol, O = 3.33 mol
- Divide by the smallest mole (3.33): C = 1, H = 2, O = 1
- Reduce if necessary
- Answer: CH2O
To get the molecular formula: divide the compound’s given molar mass by the empirical formula mass to find the whole-number multiplier, then multiply every subscript in the empirical formula by it.
4. Composition of Mixtures
Unlike pure substances, mixtures have a variable composition — the ratio of components is described using percent by mass, molarity, or parts per million rather than a fixed chemical formula.
Percent by Mass
% = mass of parttotal mass ×100%
Molarity
M = mol soluteL solution
Dilution
M1V1 = M2V2
Example: What is the molarity of a solution made with 2.0 mol NaCl dissolved in 4.0 L of water?
- M = mol / L
- M = 2.0 mol4.0 L
- Answer: 0.50 M
For very dilute mixtures, concentration is often given in parts per million (ppm): 1 ppm ≈ 1 mg of solute per 1 L of solution (for dilute aqueous solutions, where density ≈ 1 g/mL).
5. Atomic Structure & Electron Configuration
| Particle | Charge | Mass (amu) | Location |
|---|---|---|---|
| Proton | +1 | ≈ 1 | Nucleus |
| Neutron | 0 | ≈ 1 | Nucleus |
| Electron | -1 | ≈ 0 (1/1840) | Electron cloud / orbitals |
Electron Configuration Rules
| Aufbau Principle | Electrons fill the lowest-energy orbitals available first. |
| Pauli Exclusion | Each orbital holds at most 2 electrons, with opposite spins. |
| Hund’s Rule | Electrons fill degenerate (equal-energy) orbitals singly before pairing up. |
Sublevel Capacities
Aufbau Filling Order
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s
Example: Write the electron configuration for iron (Fe, Z = 26).
- Full: 1s2 2s2 2p6 3s2 3p6 4s2 3d6
- Check electron count: 2+2+6+2+6+2+6 = 26 ✓
- Noble gas shorthand: [Ar] 4s2 3d6
6. Photoelectron Spectroscopy (PES)
Photoelectron spectroscopy (PES) is an analytical technique used to determine the binding energies of electrons in atoms or molecules. It shines high-energy radiation (like UV or X-rays) onto a sample to eject electrons, measuring their kinetic energy to reveal electronic structure, subshell populations, and elemental composition.
Reading a PES Spectrum
- x-axis = binding energy (plotted decreasing left to right)
- Peaks farthest left (highest binding energy) → electrons closest to the nucleus (1s, core electrons)
- Peaks farthest right (lowest binding energy) → valence electrons, least tightly held
- Peak height/area is proportional to the number of electrons in that subshell
- Number of peaks = number of occupied subshells; relative heights confirm the electron configuration
Example: Nitrogen (Z = 7), configuration 1s2 2s2 2p3
| Peak (left to right) | Subshell | Relative Binding Energy | Relative Height (electrons) |
|---|---|---|---|
| 1st | 1s | Highest | 2 |
| 2nd | 2s | Medium | 2 |
| 3rd | 2p | Lowest | 3 |
7. Periodic Trends
Trends are driven by two competing effects: effective nuclear charge (Zeff = Z − shielding electrons), which increases across a period and pulls electrons in tighter, and the addition of new occupied energy levels down a group, which pushes electrons farther from the nucleus.
Definitions
- Atomic Radius: half the distance between the nuclei of two identical bonded atoms; a measure of the size of an atom.
- Ionization Energy: the minimum energy required to remove the most loosely held electron from a gaseous atom or ion.
- Electronegativity: a measure of an atom’s ability to attract shared (bonding) electrons toward itself in a chemical bond.
- Electron Affinity: is the amount of energy released when a neutral atom in the gaseous phase gains an extra electron to form a negative ion. A higher or more negative value means an atom is more stable and eager to accept an electron.
| Property | Across a Period (left → right) | Down a Group |
|---|---|---|
| Atomic Radius | Decreases — higher Zeff pulls electrons closer | Increases — new energy levels added |
| Ionization Energy | Increases — electrons held more tightly | Decreases — outer electrons farther from nucleus |
| Electronegativity | Increases — stronger pull on shared/bonding electrons | Decreases — weaker pull, more shielding |
| Electron Affinity | Generally increases (more negative/exothermic) | Generally decreases |
Noble gases are exceptions to ionization energy and electron affinity trends due to their already-stable, filled valence shells. Francium (Fr) has the lowest ionization energy; fluorine (F) has the highest electronegativity of any element.
8. Valence Electrons & Ionic Compounds
Valence electrons are the electrons in an atom’s outermost occupied energy level; they determine an element’s chemical bonding behavior. For main-group elements, the number of valence electrons equals the group number (using the 1–18 numbering, groups 13–18 have 3–8 valence electrons respectively).
Common Ion Charges by Group
| Group | 1A | 2A | 13 | 15 | 16 | 17 | 18 |
|---|---|---|---|---|---|---|---|
| Ion charge | +1 | +2 | +3 | -3 | -2 | -1 | 0 |
Ionic Bond Formation
Metals lose valence electrons to form positive cations; nonmetals gain electrons to form negative anions. Both achieve a stable, noble-gas electron configuration (octet rule). The oppositely charged ions attract electrostatically to form an ionic compound, which must be electrically neutral overall.
Example: Find the formula for the ionic compound formed from aluminum and oxygen.
- Aluminum forms Al3+; oxygen forms O2-
- Criss-cross the charges (drop signs) to use as subscripts: Al2O3
- Check neutrality: 2(+3) + 3(−2) = 0 ✓
- Answer: Al2O3
9. Nomenclature
Naming Ionic Compounds
Type I metals have one fixed charge (main-group metals, plus Ag+, Zn2+, Cd2+). Type II metals (transition metals, Pb, Sn, etc.) have variable charge, so a Roman numeral showing that charge is included in the name.
Formula → Name
- Identify the cation (metal) and anion (nonmetal)
- Name the cation: element name (Type I), or element name + (Roman numeral) charge (Type II)
- Name the anion: monatomic → root + “-ide”; polyatomic → use its name as-is
- Combine: cation name + anion name
Name → Formula
- Write the ion symbols with their charges
- Criss-cross the charges (drop signs) into subscripts
- Reduce subscripts to the lowest whole-number ratio
Examples
- Na2O → sodium oxide (Type I)
- FeCl3 → iron(III) chloride (Type II — 3 Cl- means Fe must be +3)
- Ca(NO3)2 → calcium nitrate (polyatomic anion)
- “Copper(II) sulfate” → Cu2+ + SO42- → CuSO₄
- “Aluminum oxide” → Al3+ + O2- → Al₂O₃ (criss-crossed)
Naming Covalent (Molecular) Compounds
Used for compounds of two nonmetals. Greek prefixes show how many atoms of each element are present. The first element drops the prefix “mono-”; the second element always gets a prefix and ends in “-ide.”
Formula → Name
- Name element 1 with its prefix (omit “mono-” for element 1 only)
- Name element 2 with its prefix + root + “-ide”
- Drop a vowel for easier pronunciation where needed (e.g. “monoxide,” “tetroxide”)
Name → Formula
- Each prefix gives the subscript for that element
- No prefix (element 1 only) means a subscript of 1
- Write the two element symbols with those subscripts
Examples
- CO2 → carbon dioxide
- CO → carbon monoxide
- N2O5 → dinitrogen pentoxide
- “Sulfur hexafluoride” → SF₆
- “Diphosphorus pentoxide” → P₂O₅
Common Polyatomic Ions
Ions with -1 Charge
Ions with -2 Charge
Ions with -3 Charge
Ions with +1 Charge
Halogen Oxyanion Series (all -1 charge)
Bromine, chlorine, and iodine each form a parallel family of oxyanions — more oxygens shift the prefix/suffix pattern:
| Pattern | Bromine | Chlorine | Iodine |
|---|---|---|---|
| per-...-ate (most O) | Perbromate BrO₄- | Perchlorate ClO₄- | Periodate IO₄- |
| -ate | Bromate BrO₃- | Chlorate ClO₃- | Iodate IO₃- |
| -ite | Bromite BrO₂- | Chlorite ClO₂- | Iodite IO₂- |
| hypo-...-ite (least O) | Hypobromite BrO- | Hypochlorite ClO- | Hypoiodite IO- |
Naming Acids
Binary acids (H + one nonmetal, no oxygen) and oxyacids (H + a polyatomic ion containing oxygen) follow different naming patterns:
| Anion Pattern | Acid Name Pattern | Example |
|---|---|---|
| “-ide” (binary, no oxygen) | “hydro-” + root + “-ic acid” | HCl (chloride) → hydrochloric acid |
| “-ate” (polyatomic + O) | root + “-ic acid” | HNO3 (nitrate) → nitric acid |
| “-ite” (polyatomic + O, less O) | root + “-ous acid” | HNO2 (nitrite) → nitrous acid |
Formula → Name
- Identify the anion attached to H
- Match its ending (-ide, -ate, -ite) to the acid pattern above
- Apply the matching prefix/suffix
Name → Formula
- Find the matching anion (hydro-...-ic → -ide; -ic → -ate; -ous → -ite)
- Add enough H+ to balance the anion’s charge
- Write the resulting formula
Examples
- H2S (sulfide, no oxygen) → hydrosulfuric acid
- H2SO4 (sulfate) → sulfuric acid
- H2SO3 (sulfite) → sulfurous acid
- “Carbonic acid” (carbonate CO32-) → H₂CO₃