CHEM 111 — Exam 1 Study Guide

Chapters 1–3 · Fall 2026

Contents

Chapter 1 — Matter & Measurement
Chapter 2 — Atomic Theory & Nomenclature
Chapter 3 — The Mole & Stoichiometry
Exam Strategy Tips

Chapter 1 — Matter & Measurement

Classifying Matter

Matter splits first into pure substances and mixtures. Pure substances have a fixed composition; mixtures can vary.

Pure substance
- Element — one type of atom (e.g., Fe, O₂). Cannot be broken down chemically.
- Compound — two+ elements chemically bonded in a fixed ratio (e.g., H₂O, NaCl). Can be broken down chemically.
Mixture — physically combined substances, separable by physical means
- Homogeneous — uniform composition throughout (you can't see separate parts). Ex: salt water, air.
- Heterogeneous — non-uniform; separate regions/phases visible. Ex: oil + water, sand in water.

Quick test: Ask "if I take a sample from two different spots, is it identical?" Yes → homogeneous. No → heterogeneous.

Uncertainty in Measurements

Estimating from instruments: Always record all certain digits plus one estimated digit (the smallest marked division, estimated to the nearest tenth of that division).

Significant Figures — Rules

Rule	Example
Nonzero digits always count	123 → 3 sig figs
Zeros between nonzero digits count	1002 → 4 sig figs
Leading zeros never count	0.0025 → 2 sig figs
Trailing zeros count only if there's a decimal point	100. → 3 sig figs; 100 → 1 sig fig
Trailing zeros after a decimal always count	1.500 → 4 sig figs

Math rules:

Multiply/divide → answer has the same number of sig figs as the input with the fewest.
Add/subtract → answer has the same number of decimal places as the input with the fewest decimal places.

Metric Prefixes (memorize these four + base)

Prefix	Symbol	Meaning (× base unit)
kilo	k	1,000 (10³)
— (base)	—	1
centi	c	0.01 (10⁻²)
milli	m	0.001 (10⁻³)
micro	µ	0.000001 (10⁻⁶)

Trick: Move the decimal point — converting to a smaller unit (e.g., m→cm) means a bigger number, and vice versa.

The 7 Base SI Units

Quantity	Unit	Symbol
Length	meter	m
Mass	kilogram	kg
Time	second	s
Temperature	kelvin	K
Amount of substance	mole	mol
Electric current	ampere	A
Luminous intensity	candela	cd

Unit Conversion — Dimensional Analysis

desired unit = given quantity × (conversion factor)

Chain factors so units cancel diagonally. Always double-check the final unit matches what's asked.

Density Conversions

Density (d) = mass (m) / volume (V)

Density acts as a conversion factor between mass and volume: e.g., g → mL using g/mL.

Example: A liquid has density 1.25 g/mL. Find the volume of 50.0 g.
50.0 g × (1 mL / 1.25 g) = 40.0 mL

Temperature Conversions

°F = (°C × 9/5) + 32 | °C = (°F − 32) × 5/9

K = °C + 273.15

Know both directions (F→C and C→F) — you'll likely be given this formula on the exam, but practice using it quickly.

Chapter 2 — Atomic Theory & Nomenclature

Historical Atomic Theory

Scientist	Experiment/Idea	Conclusion
Dalton	Dalton's Atomic Theory (Laws of Definite/Multiple Proportions, Conservation of Mass)	Matter is made of indivisible atoms; atoms of an element are identical; compounds form in fixed ratios
J.J. Thomson	Cathode Ray Tube experiment	Discovered the electron — a negatively charged particle; proposed the plum pudding model (negative electrons scattered in a positive "pudding")
Millikan	Oil Drop experiment	Determined the charge of a single electron (and thus its mass) by balancing gravity vs. an electric field on charged oil droplets
Rutherford	Gold Foil experiment	Most alpha particles passed through foil, but some deflected sharply → atom is mostly empty space with a small, dense, positively charged nucleus (disproved plum pudding model)

Subatomic Particles in Atoms & Ions

Particle	Determines
# protons	Atomic number (Z) — defines the element; never changes for a given element
# neutrons	Mass number − atomic number (A − Z); varies → isotopes
# electrons	= protons in a neutral atom. For ions: subtract charge from protons Cation (+): fewer electrons than protons Anion (−): more electrons than protons

Example: Fe³⁺ has 26 protons, so electrons = 26 − 3 = 23 electrons.

Isotopes & Nuclear Notation

^A_ZX (A = mass number, Z = atomic number, X = element symbol)

Isotopes = atoms of the same element (same Z) with different numbers of neutrons (different A).

Mass number (A) = protons + neutrons (always a whole number for a single isotope).

Natural abundance = the % of each isotope found in nature, used as weighting factors for average atomic mass.

Average Atomic Mass

Avg. mass = Σ (isotope mass × fractional abundance)

Example: Cl-35 (mass 34.97, abundance 75.76%) and Cl-37 (mass 36.97, abundance 24.24%):
(34.97 × 0.7576) + (36.97 × 0.2424) = 35.45 amu

Elements 1–86

You need to know symbol ↔ name for elements 1–86. Rather than reproduce all 86 here, drill yourself with these high-yield groups, then use flashcards/periodic table practice for the rest:

Group	Elements
Common "tricky" symbols	Na (sodium), K (potassium), Fe (iron), Cu (copper), Ag (silver), Au (gold), Hg (mercury), Sn (tin), Pb (lead), Sb (antimony), W (tungsten)
Alkali metals (Grp 1)	H, Li, Na, K, Rb, Cs
Alkaline earth (Grp 2)	Be, Mg, Ca, Sr, Ba
Halogens (Grp 17)	F, Cl, Br, I
Noble gases (Grp 18)	He, Ne, Ar, Kr, Xe
Common transition metals	Sc–Zn (row 4), plus Ag, Cd, Au, Hg, Pt

Study tip: print a blank periodic table (1–86) and quiz yourself filling in names/symbols until it's automatic.

Naming Ionic Compounds

Basic pattern: cation name + anion name (anion of a single element ends in -ide).

Example: NaCl → sodium chloride. CaO → calcium oxide.

Common Polyatomic Ions (memorize charges!)

Ion	Name	Ion	Name
NH₄⁺	ammonium	OH⁻	hydroxide
NO₃⁻	nitrate	NO₂⁻	nitrite
SO₄²⁻	sulfate	SO₃²⁻	sulfite
CO₃²⁻	carbonate	HCO₃⁻	bicarbonate/hydrogen carbonate
PO₄³⁻	phosphate	PO₃³⁻	phosphite
ClO⁻	hypochlorite	ClO₂⁻	chlorite
ClO₃⁻	chlorate	ClO₄⁻	perchlorate
C₂H₃O₂⁻ (or CH₃COO⁻)	acetate	CN⁻	cyanide
MnO₄⁻	permanganate	Cr₂O₇²⁻	dichromate
CrO₄²⁻	chromate	C₂O₄²⁻	oxalate

Pattern trick: "-ate" ions usually have one more O than the "-ite" version, same charge (e.g., NO₃⁻ vs NO₂⁻). "Per-...-ate" has one more O than "-ate"; "hypo-...-ite" has one fewer O than "-ite."

Naming with Transition Metals (variable charge)

Many transition metals form more than one charge, so you must indicate charge with a Roman numeral in parentheses (this is the "Stock system").

Example: Fe₂O₃ → each O is 2−, three O's = 6− total, so 2 Fe must total 6+ → each Fe is 3+ → iron(III) oxide.
FeO → iron(II) oxide.

Exceptions that don't need Roman numerals (fixed charge): Ag⁺, Zn²⁺, Cd²⁺.

Naming Molecular (Covalent) Compounds

Use Greek prefixes for both elements (mono- is dropped for the first element only).

#	Prefix	#	Prefix
1	mono-	6	hexa-
2	di-	7	hepta-
3	tri-	8	octa-
4	tetra-	9	nona-
5	penta-	10	deca-

Example: N₂O₄ → dinitrogen tetroxide. CO₂ → carbon dioxide (not "monocarbon").

Naming Acids

Acid type	Rule	Example
Binary acid (H + one nonmetal)	hydro- + root + -ic acid	HCl → hydrochloric acid
Oxyacid from "-ate" ion	root + -ic acid	H₂SO₄ (sulfate) → sulfuric acid
Oxyacid from "-ite" ion	root + -ous acid	H₂SO₃ (sulfite) → sulfurous acid

Chapter 3 — The Mole & Stoichiometry

The Mole & Avogadro's Number

1 mole = 6.022 × 10²³ particles (Avogadro's number)

Molar Mass

Molar mass (g/mol) = sum of atomic masses (from the periodic table) of all atoms in a formula.

Example: H₂O = 2(1.008) + 16.00 = 18.02 g/mol

Mole ↔ Mass ↔ Particle Conversions

grams ⇄ moles: use molar mass (g/mol)

moles ⇄ particles: use Avogadro's number (6.022×10²³/mol)

Set up as a chain: grams → moles → particles (or reverse), canceling units at each step.

Percent Composition

% element = (mass of element in formula / molar mass of compound) × 100%

Example: % H in H₂O = (2.016/18.02) × 100 = 11.19%

Empirical Formula

The simplest whole-number ratio of atoms in a compound.

Assume 100 g sample → convert each given % directly to grams.
Convert grams of each element to moles (divide by atomic mass).
Divide all mole values by the smallest mole value.
If not whole numbers, multiply all by a small integer to get whole numbers.

Molecular Formula

Molecular formula = empirical formula × n, where n = (molar mass given) / (empirical formula mass)

Example: Empirical formula CH₂O (mass 30 g/mol), actual molar mass = 180 g/mol → n = 6 → molecular formula = C₆H₁₂O₆

Combustion Analysis

Used to find empirical formulas of compounds containing C, H, (and O) by burning a sample completely in O₂ and measuring the CO₂ and H₂O produced.

All C in sample ends up in CO₂ → mol C = mol CO₂ produced.
All H in sample ends up in H₂O → mol H = 2 × mol H₂O produced.
If the compound contains O: mass O = (mass sample) − (mass C) − (mass H); convert to moles.
Find mole ratios → empirical formula (same method as above).

Balancing Chemical Equations

Law of Conservation of Mass: atoms of each element must be equal on both sides.

Write correct formulas for all reactants/products first (don't touch subscripts).
Balance one element at a time, usually starting with the most complex molecule.
Balance elements that appear alone (like O₂ or H₂) last.
Double-check every element's atom count matches on both sides.
Reduce coefficients to smallest whole-number ratio if needed.

Example: C₃H₈ + O₂ → CO₂ + H₂O
Balance C: C₃H₈ + O₂ → 3CO₂ + H₂O
Balance H: C₃H₈ + O₂ → 3CO₂ + 4H₂O
Balance O (right side has 6+4=10 O): C₃H₈ + 5O₂ → 3CO₂ + 4H₂O

Exam Strategy Tips

Sig figs & dimensional analysis show up everywhere — make sure unit cancellation is second nature before the exam.
Memorize polyatomic ion charges cold. Nomenclature questions are fast points if you know these; costly if you don't.
Practice going both directions: name → formula, and formula → name, for ionic, molecular, and acid compounds.
For the mole chapter, draw out the conversion map (mass ⇄ moles ⇄ particles) so you always know which conversion factor to grab.
Redo the historical atomic-theory experiments as a table (scientist → method → conclusion) — these are easy to mix up under time pressure.
Work several empirical/molecular formula and combustion analysis problems start-to-finish — these have multiple steps where small arithmetic errors compound.