Chemistry Regents Study Guide

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Based on this guide's real question bank — 469 practice questions across 16 units. Slide to match your situation.

Days until your exam 14 days

Minutes you can study per day 30 min

Questions you'll complete

Units covered

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Term

Definition

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Core Formulas (Table T)

Density

D = m / V

% Error

|meas − accept| / accept × 100

Moles

n = m / GFM

Avogadro

1 mol = 6.02×10²³ particles

Kelvin

K = °C + 273

q = mCΔT

Sloped sections — T changing

q = mH_f

Melting/freezing (flat at MP)

q = mH_v

Boiling/condensing (flat at BP)

pH = −log[H⁺]

pH + pOH

= 14 (at 25°C)

[H⁺][OH⁻] = 1×10⁻¹⁴

Titration (1:1)

M_A × V_A = M_B × V_B

Molarity

M = mol solute / L solution

ppm

(m_solute / m_solution) × 10⁶

Dilution

M₁V₁ = M₂V₂

Combined Gas

P₁V₁/T₁ = P₂V₂/T₂

% Composition

(mass part / total mass) × 100

Half-Life

A = A₀ × (½)^(t/t₁/₂)

ΔH (bonds)

ΔH = bonds broken − bonds formed

Atomic Mass

Σ(mass × fractional abundance)

Molar Volume (STP)

1 mol gas = 22.4 L

% Yield

actual / theoretical × 100

ΔG (Gibbs)

ΔG = ΔH − TΔS (spont. if < 0)

Stoichiometry

g ÷ GFM_given × ratio × GFM_want

Nuclear Particle Changes

Decay Type	Symbol	Mass # Change	Atomic # Change	Penetration
Alpha (α)	⁴₂He	−4	−2	Stopped by paper
Beta⁻ (β⁻)	⁰₋₁e	0	+1	Stopped by aluminum
Positron (β⁺)	⁰₊₁e	0	−1	Stopped by aluminum
Gamma (γ)	⁰₀γ	0	0	Needs lead/concrete

Indicator Table (Table M)

Indicator	Transition pH	Acid Color	Base Color
Methyl orange	3.1–4.4	Red	Yellow
Bromcresol green	3.8–5.4	Yellow	Blue
Litmus	4.5–8.3	Red	Blue
Bromthymol blue	6.0–7.6	Yellow	Blue
Thymol blue	8.0–9.6	Yellow	Blue
Phenolphthalein	8.2–10.0	Colorless	Pink/Fuchsia

Oxidation Number Rules (Priority Order)

Free elements: any uncombined element = 0 (Na, O₂, Cl₂, Fe)
Monatomic ions: oxidation # = ion charge (Na⁺=+1, Cl⁻=−1, Fe³⁺=+3)
Group 1 in compounds: always +1 (Li, Na, K, Rb, Cs)
Group 2 in compounds: always +2 (Mg, Ca, Ba, Sr)
Fluorine: always −1 (most electronegative element)
Hydrogen: +1 usually; −1 in metal hydrides (NaH, CaH₂, LiH)
Oxygen: −2 usually; −1 in peroxides (H₂O₂, Na₂O₂); +2 with F only
Sum rule: sum of all oxidation numbers = 0 (neutral) or = ion charge (polyatomic)

Half-Life Reference Table

# Half-Lives	Fraction Left	% Remaining	% Decayed
0	1	100%	0%
1	1/2	50%	50%
2	1/4	25%	75%
3	1/8	12.5%	87.5%
4	1/16	6.25%	93.75%
5	1/32	3.125%	96.875%

VSEPR Molecular Geometry

Bonding Pairs	Lone Pairs	Geometry	Bond Angle	Example
2	0	Linear	180°	CO₂, BeCl₂
3	0	Trigonal planar	120°	BF₃, SO₃
4	0	Tetrahedral	109.5°	CH₄, CCl₄
3	1	Trigonal pyramidal	107°	NH₃
2	2	Bent/V-shaped	104.5°	H₂O

Electrochemistry Summary

Property	Voltaic (Galvanic) Cell	Electrolytic Cell
Energy conversion	Chemical → Electrical	Electrical → Chemical
Spontaneous?	Yes	No (needs external power)
Anode sign	Negative (−)	Positive (+)
Cathode sign	Positive (+)	Negative (−)
At anode	Oxidation (always)	Oxidation (always)
At cathode	Reduction (always)	Reduction (always)
Electron flow	Anode → Cathode (wire)	Cathode ← Anode (wire)

Strong Acids & Bases (Memorize)

Strong Acids (6)

HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄

Strong Bases

NaOH, KOH, Ca(OH)₂, Ba(OH)₂, LiOH

Weak Acid examples

HF, CH₃COOH, H₂CO₃, H₃PO₄

Weak Base examples

NH₃, Al(OH)₃

Reference Table Guide (All Tables)

Table A

STP: 273 K, 101.3 kPa, 1 atm, 760 mmHg

Table B

Water constants: C=4.18 J/g·K, H_f=334 J/g, H_v=2260 J/g

Table C

Selected polyatomic ions (NH₄⁺, OH⁻...)

Table D

Selected units and measurements

Table E

Polyatomic ions (NO₃⁻, SO₄²⁻, CO₃²⁻, PO₄³⁻...)

Table F

Solubility rules — predicts precipitates

Table G

Solubility curves (g/100g H₂O vs temp)

Table H

Vapor pressure curves (BP = where VP = 101.3 kPa)

Table I

Heats of reaction (ΔH values in kJ/mol)

Table J

Activity series — metal/halogen reactivity ranking

Table K

Common acids and their formulas

Table L

Common bases and their formulas

Table M

Indicators — pH ranges and color changes

Table N

Radioisotopes — half-lives, decay modes, medical uses

Table O

Nuclear symbols (α, β, γ, positron)

Table P

Organic functional groups

Table Q

Organic prefixes (meth=1, eth=2, prop=3...)

Table R

Organic reactions

Table S

Properties of selected elements (EN, radius, ionization E)

Table T

Important formulas (density, gas laws, pH, etc.)

Electronegativity & Bond Type

ΔEN Range	Bond Type	Example
0 – 0.4	Nonpolar covalent	Cl-Cl, C-H, H-H
0.4 – 1.7	Polar covalent	H-O, H-N, H-Cl
≥ 1.7	Ionic	Na-Cl, Mg-O, Ca-F

Gas Laws Quick Reference

Law	Formula	Constant	Relationship
Boyle's	P₁V₁ = P₂V₂	T, n	P and V inverse
Charles's	V₁/T₁ = V₂/T₂	P, n	V and T direct
Gay-Lussac's	P₁/T₁ = P₂/T₂	V, n	P and T direct
Combined	P₁V₁/T₁ = P₂V₂/T₂	n	All three
Avogadro's	Equal V at same T,P = equal n	T, P	V and n direct

Organic Chemistry Quick Reference

Series	General Formula	Bond Type	Example
Alkanes	CₙH₂ₙ₊₂	Single only (saturated)	CH₄, C₂H₆, C₃H₈
Alkenes	CₙH₂ₙ	One double bond	C₂H₄, C₃H₆
Alkynes	CₙH₂ₙ₋₂	One triple bond	C₂H₂ (acetylene)
Cycloalkanes	CₙH₂ₙ	Ring, single bonds	C₃H₆ (cyclopropane)

Functional Group	Name Class	Suffix	Example
−OH	Alcohol	-ol	ethanol (C₂H₅OH)
−COOH	Carboxylic acid	-oic acid	ethanoic acid (CH₃COOH)
−COO−	Ester	-anoate	ethyl ethanoate
−CHO	Aldehyde	-al	methanal (formaldehyde)
C=O (chain)	Ketone	-one	propanone (acetone)
−NH₂	Amine	-amine	methylamine

Reaction (Table R)	Reactants	Products	Key fact
Fermentation	C₆H₁₂O₆	2C₂H₅OH + 2CO₂	Yeast, anaerobic, makes alcohol
Esterification	Acid + Alcohol	Ester + H₂O	Reversible
Saponification	Ester + NaOH	Soap + Alcohol	Reverse of esterification
Addition	Alkene + H₂/Cl₂	Saturated product	Breaks double bond
Substitution	Alkane + halogen	Haloalkane + HX	Replaces H with halogen
Combustion	CₓHᵧ + O₂	CO₂ + H₂O	Complete combustion always
Polymerization	n(monomer)	Polymer	Ethene → polyethylene

Naming prefixes

meth=1, eth=2, prop=3, but=4, pent=5, hex=6, hept=7, oct=8

Isomers rule

Same molecular formula, different structural formula → isomers

Allotropes (Q7 exam)

O₂ vs O₃: same element, different structure → different properties

Organic = contains C

Exceptions: CO, CO₂, CO₃²⁻, CN⁻ are inorganic

Crystallization & Stoichiometry Quick Reference

Crystallization trigger

Cool saturated solution OR evaporate solvent → supersaturated → crystals form

Amount crystallized

Solubility(T_high) − Solubility(T_low) per 100 g H₂O

Seed crystal

Nucleation surface; triggers crystallization from supersaturated solution

Best candidate for recrystallization

Substance with steep solubility curve (e.g., KNO₃)

Stoichiometry roadmap

g ÷ GFM_given × (mol wanted/mol given) × GFM_wanted

Limiting reagent test

Convert both reactants to mol of product → smaller answer = limiting

% yield

actual ÷ theoretical × 100 (never > 100%)

Gas stoichiometry (STP)

Volumes react in same ratio as coefficients; 1 mol = 22.4 L

Periodic Trends Summary

Property	Across Period (L→R)	Down Group (↓)	Largest Value
Atomic radius	Decreases	Increases	Fr (bottom-left)
Ionization energy	Increases	Decreases	He (top-right)
Electronegativity	Increases	Decreases	F (4.0)
Metallic character	Decreases	Increases	Fr (bottom-left)
Ionic radius (cations)	Decreases	Increases	Cs⁺

Table I — Heats of Dissolving (Tested on Regents!)

Compound	ΔH (dissolving)	Type	Effect on solution
NaOH(s)	Negative (−)	Exothermic	Solution gets WARM
NH₄NO₃(s)	Positive (+)	Endothermic	Solution gets COLD (cold packs)
KNO₃(s)	Positive (+)	Endothermic	Solution gets COLD
NH₄Cl(s)	Positive (+)	Endothermic	Solution gets COLD

Physical vs Chemical Equilibrium

Type	Example	What's equal	Closed system?
Physical equilibrium	Sealed flask with liquid H₂O	Evaporation rate = condensation rate	Yes
Chemical equilibrium	N₂ + 3H₂ ⇌ 2NH₃	Forward rate = reverse rate	Yes

Reading conc-vs-time graph

Equilibrium = where ALL lines become flat and stay flat

PE during evaporation

PE INCREASES as molecules escape IMFs; KE (temp) of remaining liquid DECREASES → cooling

Heat flow direction

Always flows from HIGH temperature to LOW temperature

Control group (lab)

The unchanged baseline — e.g., dry towel in evaporation experiment

Top 15 Exam Traps

Always convert °C to Kelvin before using gas law formulas (K = °C + 273).
Dissolving is a PHYSICAL change — NaCl can be recovered by evaporation.
Activation energy (Ea) is measured from REACTANTS up to the peak — not from zero.
Catalyst lowers Ea but does NOT change ΔH or shift equilibrium position.
CO₂ has polar bonds but is a nonpolar MOLECULE — linear geometry cancels dipoles.
Ionic solids do NOT conduct electricity — only when molten or dissolved (ions mobile).
Saturated ≠ concentrated. Saturated = at maximum capacity (on the solubility curve line).
pH scale is LOGARITHMIC: each 1-unit difference = 10× change in [H⁺].
Single replacement: check Table J first. If replacing element is lower → NR.
Only COEFFICIENTS can be changed when balancing — never subscripts.
Voltaic vs. electrolytic: anode terminal signs flip, but oxidation is ALWAYS at the anode.
Nuclear equations: BOTH mass numbers AND atomic numbers must balance.
Half-life is constant — unaffected by temperature, pressure, or chemical form.
Excited state still has the correct number of electrons — they're just not in lowest levels.
Strength ≠ concentration — a dilute HCl solution is still a STRONG acid (100% ionized).

Exam Day Checklist

Use Table T first — it lists all the formulas you'll need.
Know which Reference Table to use: J (activity series), F (precipitates), M (indicators), N (radioisotopes), G (solubility curves).
Show all work with units. Dimensional analysis prevents errors.
Start with questions you know best; circle and skip difficult ones.
Gas laws: write down given values, identify what's constant, pick the right formula.
Nuclear equations: always verify mass numbers AND atomic numbers balance.
For PE diagrams: mark reactants, peak, and products before reading Ea or ΔH.
Treat heat as a substance (reactant or product) for Le Chatelier problems.
Oxidation number assignments: use the 8 rules in priority order.
Check units in your final answer before moving on.

Quick ways to lock in the facts you keep forgetting. Read the big trick, then the small note tells you what it unlocks. Say them out loud — silly is memorable.

Matter, Measurement & Atoms

Can you get it back?

Physical change: the original substance can be recovered (melt ice, freeze it back). Chemical change: you can't un-burn wood — a genuinely new substance formed.

Fixed = compound, flexible = mixture

Compounds have a fixed ratio of elements (H₂O is always 2:1). Mixtures can be any ratio — that's what makes them physically separable.

Distill by degree, filter what's firm

Distillation separates liquids by different boiling points. Filtration catches an insoluble solid in a filter.

Zeros: leading lies, trailing tells, captive counts

Leading zeros (0.0025) are never significant. Trailing zeros only count WITH a decimal point (1500. = 4 sig figs). Captive zeros (2005) always count.

Add keeps the least decimal, multiply keeps the least digit

Adding/subtracting: round to the fewest decimal places among your numbers. Multiplying/dividing: round to the fewest total sig figs.

Same P, different N = isotope

Isotopes have the same number of protons but a different number of neutrons — mass number varies, atomic number doesn't.

Periodic Table & Bonding

Down gets big, across gets small

Atomic radius increases going down a group (more energy levels) and decreases going across a period (more protons pull electrons in).

Ionization energy is the opposite of radius

Small atom = electrons held tight = high ionization energy. IE increases across a period, decreases down a group — the reverse of atomic radius.

Fluorine is the greediest

Electronegativity peaks at fluorine (top-right of the table, excluding noble gases) — it pulls shared electrons hardest of any element.

Metal + nonmetal = ionic, nonmetal + nonmetal = covalent

Ionic bonds transfer electrons (big electronegativity gap). Covalent bonds share electrons (both nonmetals, similar electronegativity).

Symmetric cancels, lopsided doesn't

A molecule with polar bonds can still be nonpolar overall if its shape is symmetric (like CO₂) — the bond dipoles cancel out.

VSEPR: count the groups around the center

4 bonding groups, 0 lone pairs = tetrahedral. 4 groups, 1 lone pair = trigonal pyramidal. 4 groups, 2 lone pairs = bent.

Quantities & Reactions

Moles are always the bridge

Every stoichiometry problem goes grams → moles → moles → grams. You can never skip straight from one substance's grams to another's.

Limiting runs out first, excess is left over

Whichever reactant produces the smaller amount of product is the limiting reagent — it runs out first and stops the reaction.

Yield can't beat 100%

% yield = actual ÷ theoretical × 100. An answer over 100% means a math error somewhere.

One becomes many, many become one

Synthesis: two reactants → one product. Decomposition: one reactant → many products — think of them as opposites.

Higher on the list wins

Table J's activity series: a metal can only replace a less active metal below it. Asking it to replace something above it means "no reaction."

Balance coefficients, never subscripts

Only the big numbers in front of formulas can change when balancing. Changing a subscript changes the substance itself.

Phases & Gases

Flat means it's changing phase

On a heating curve, a flat plateau means energy is breaking intermolecular forces, not raising temperature — melting or boiling in progress.

Gas particles don't care about each other

KMT: gas particles have negligible volume and no attraction to each other — that's why gases compress easily and fill any container.

Tyndall = you can see the beam

A light beam becomes visible passing through a colloid because its particles are big enough to scatter light — true solutions don't do this.

Squeeze it, it heats up pressure-wise

Boyle's Law: P and V are inversely related at constant T — squeeze the volume down, pressure goes up.

Charles says heat it, it grows

Charles's Law: V and T are directly related at constant P — heat a balloon, it expands.

STP is the universal starting line

At STP (0°C, 1 atm), 1 mole of any gas = 22.4 L. Always convert Celsius to Kelvin before plugging into a gas law.

Solutions & Thermochemistry

Above the line, it won't all dissolve

On a solubility curve, a point above the curve means undissolved solid remains — the solution at that point is saturated at most.

Concentrated ≠ saturated

Saturated means at the max the solvent can hold. Concentrated just means "a lot of solute" — a dilute solution can still be saturated if the solute barely dissolves.

Dilution: moles don't change

M₁V₁ = M₂V₂ works because adding water changes concentration, not the actual moles of solute present.

Exo EXits heat, endo takes it IN

Exothermic reactions release heat to the surroundings. Endothermic reactions absorb heat from the surroundings.

Flat parts on a PE diagram are phase changes

Sloped line = temperature changing (q = mCΔT). Flat line = phase changing at constant temperature (q = mH_f or mH_v).

Activation energy is measured from where you start

Ea is measured from the reactants up to the peak of the curve, not from zero — a common exam trap.

Equilibrium & Acids/Bases

Stress it, it shifts away

Le Chatelier's Principle: add stress (more reactant, less product, heat, pressure) and equilibrium shifts to relieve it, away from the stress.

Catalysts speed up, don't show up

A catalyst lowers activation energy and speeds a reaction, but isn't consumed and does NOT change ΔH or shift the equilibrium position.

Down the pH scale, up the acid

Each drop of 1 on the pH scale means [H⁺] is 10× bigger — pH is logarithmic, not linear.

Arrhenius: acids give H⁺, bases give OH⁻

That's the whole definition — an Arrhenius acid donates H⁺ ions in water, an Arrhenius base donates OH⁻ ions.

Strong means it splits all the way

Strong acids/bases ionize 100% in water. A dilute strong acid is still STRONG — don't confuse concentration with strength.

More collisions, more successful hits

Reaction rate increases with more concentration, more surface area, higher temperature, or a catalyst — all mean more effective collisions.

Redox & Nuclear

LEO the lion says GER

Lose Electrons = Oxidation. Gain Electrons = Reduction. Oxidation number goes up when you lose electrons, down when you gain them.

An OX and a RED CAT

Anode = OXidation (always). Cathode = REDuction (always) — true in both voltaic and electrolytic cells, even though the charge signs flip between them.

Free elements start at zero

Any uncombined element (Na, O₂, Fe) has an oxidation number of 0 — rule #1 before anything else.

Alpha is heavy, beta is light, gamma is nothing

Alpha decay: mass −4, atomic number −2. Beta decay: mass unchanged, atomic number +1. Gamma: neither changes — pure energy.

Half-life doesn't care about anything else

Half-life is constant no matter the temperature, pressure, or chemical form the isotope is in — a purely nuclear property.

Fission splits big, fusion joins small

Fission breaks a heavy nucleus (like U-235) into smaller pieces. Fusion joins light nuclei (like hydrogen) into a bigger one — the sun runs on fusion.

Organic Chemistry

Meth-Eth-Prop-But: count the carbons

Methane(1)-Ethane(2)-Propane(3)-Butane(4) — the prefix always tells you the carbon count, no matter the functional group.

-ane is full, -ene has one, -yne has two

Saturated hydrocarbons (all single bonds) end in -ane. One double bond = -ene. One triple bond = -yne.

Same family, one CH₂ apart

A homologous series is a family with the same general formula and functional group, where each member differs from the next by exactly one CH₂ unit.

Chemistry Study Guide Expanded Edition

See what your study plan could look like

Core Formulas (Table T)

Nuclear Particle Changes

Indicator Table (Table M)

Oxidation Number Rules (Priority Order)

Half-Life Reference Table

VSEPR Molecular Geometry

Electrochemistry Summary

Strong Acids & Bases (Memorize)

Reference Table Guide (All Tables)

Electronegativity & Bond Type

Gas Laws Quick Reference

Organic Chemistry Quick Reference

Crystallization & Stoichiometry Quick Reference

Periodic Trends Summary

Table I — Heats of Dissolving (Tested on Regents!)

Physical vs Chemical Equilibrium

Top 15 Exam Traps

Exam Day Checklist

Matter, Measurement & Atoms

Periodic Table & Bonding

Quantities & Reactions

Phases & Gases

Solutions & Thermochemistry

Equilibrium & Acids/Bases

Redox & Nuclear

Organic Chemistry