Chemistry Regents Review Sheet

Q: How should I study for the Chemistry Regents?

Work real past-exam questions by topic, review this one-page summary of the must-know concepts, and read the worked solution every time you miss one. The exam reuses the same ideas year after year.

Chemistry is bookkeeping: atoms in, atoms out. Get comfy with the periodic table and half the exam solves itself.

Practice Chemistry free Reference & formulas

Big ideas

Everything is made of atoms. Atoms combine via bonds. Bonds break and reform in reactions. Nothing else.
The periodic table isn't decoration — it TELLS you charges, sizes, reactivity, and electron configurations.
The mole is the chemist's dozen: 6.022 × 10²³ of anything. It's how we count invisible particles by weighing them.

Atoms, ions, isotopes

Protons make it an element. Neutrons make it an isotope. Electrons make it an ion.

Atomic number vs mass

Atomic number = protons. Mass number = protons + neutrons.

Example: Carbon-14. Protons? Neutrons? → 6 protons (C is #6), 14 − 6 = 8 neutrons

Ion charges

Metals lose electrons → positive ions (cations). Nonmetals gain → negative (anions). Charge = group number logic.

Example: Mg (group 2) → ? → Loses 2e⁻ → Mg²⁺

Electron configuration

Fill orbitals lowest energy first: 1s → 2s → 2p → 3s → 3p → 4s → 3d ... Regents just needs valence electrons.

Example: Sodium, 11 electrons, config? → 2-8-1 (valence = 1)

Isotope math

\bar{m} = \sum (\text{fraction}) \times \text{mass}

Atomic mass on the periodic table is a WEIGHTED average of isotopes.

Example: Cl-35 (75%) and Cl-37 (25%). Average? → 0.75(35) + 0.25(37) = 35.5

Trap: Mass number is a WHOLE number for one isotope. Atomic mass on the table (like 35.45) is the weighted average.

Tip: Draw a Bohr diagram if you're stuck on electron count: 2 in first shell, 8 in second, 8 in third for basic subjects.

Periodic trends

Learn the direction, derive the rest. Everything follows from electron pull vs shielding.

Atomic radius

Bigger DOWN (more shells). Smaller ACROSS a period (more protons pull electrons closer).

Example: Larger: Li or Cs? → Cs (further down)

Ionization energy

Energy to REMOVE an electron. Opposite of radius: smaller atom, harder to strip.

Example: Higher IE: Na or Cl? → Cl (smaller, tighter grip)

Electronegativity

Pull on bonding electrons. F is #1. Increases across, decreases down.

Example: Rank EN: Br, Cl, F, I → F > Cl > Br > I

Metallic character

Opposite of EN. Most metallic in bottom-left (Cs, Fr). Most nonmetallic in top-right (F).

Example: More metallic: K or Ca? → K (further left)

Trap: IE and radius are OPPOSITE trends. Draw arrows on the table if you get confused.

Tip: Noble gases (group 18) don't play well — they're chemically inert. Don't try to give them an ionization pattern.

Bonding

Big EN gap → ionic (transfer electrons). Small gap → covalent (share). Metal-metal → metallic (sea of electrons).

Ionic vs covalent

Metal + nonmetal usually ionic. Two nonmetals covalent. EN difference > ~1.7 = ionic.

Example: NaCl vs H₂O? → NaCl ionic (metal + nonmetal), H₂O covalent (2 nonmetals)

Polar covalent

Covalent but unequal sharing (EN difference 0.5–1.7). Creates partial + and − ends.

Example: H–Cl bond? → Polar covalent (EN diff ~0.9)

Molecular shape

VSEPR: lone pairs and bonds repel to spread out. Linear, bent, trigonal, tetrahedral, etc.

Example: H₂O shape? → Bent (2 bonds + 2 lone pairs on O)

Intermolecular forces

Between molecules (not within). Weakest: London dispersion. Then dipole-dipole. Strongest: hydrogen bonding (H bonded to N, O, or F).

Example: Why does water have high boiling point? → Hydrogen bonding

Trap: Polar BOND ≠ polar MOLECULE. CO₂ has polar bonds but the molecule is linear so dipoles cancel out.

Tip: For solubility: 'like dissolves like.' Polar solvents dissolve polar solutes; nonpolar dissolves nonpolar.

Reactions & stoichiometry

Balance atoms on both sides, then use moles to relate amounts.

Reaction types

Synthesis A+B→AB. Decomposition AB→A+B. Single replacement A + BC → AC + B. Double replacement AB + CD → AD + CB. Combustion CxHy + O₂ → CO₂ + H₂O.

Example: 2H₂ + O₂ → 2H₂O → Synthesis

Balancing

Same atoms on both sides. Change COEFFICIENTS only, never subscripts. Balance one element at a time.

Example: H₂ + O₂ → H₂O → 2H₂ + O₂ → 2H₂O

Moles ↔ mass

n = m/M

n = moles, m = grams, M = molar mass. Molar mass = sum of atomic masses.

Example: 18 g of H₂O → M = 18 g/mol → 1 mole

Mole ratios

Coefficients in a balanced equation are MOLE ratios. Convert given amount to moles, use ratio, convert back.

Example: How many mol O₂ needed for 4 mol CH₄? (CH₄ + 2O₂ → ...) → 4 × 2/1 = 8 mol O₂

Trap: Coefficients ratio MOLES, not grams. Convert to moles FIRST.

Tip: Limiting reactant: the one you run out of first. Convert both reactants to moles of product, take the SMALLER.

Gas laws & solutions

Gases respond to pressure, volume, temperature predictably. Solutions are all about concentration.

Combined gas law

\dfrac{P_1V_1}{T_1} = \dfrac{P_2V_2}{T_2}

T in Kelvin always. K = °C + 273.

Example: V₁ = 2 L at 300 K, T₂ = 600 K, P constant. V₂? → V₂ = 4 L

Ideal gas law

PV = nRT

One condition. R = 0.0821 L·atm/(mol·K). Use when you need moles.

Example: 1 mole of gas at STP → V = 22.4 L

Molarity

M = \dfrac{\text{mol}}{\text{L}}

Standard concentration. Moles of solute per liter of solution.

Example: 58.5 g NaCl in 1 L → 1 M (58.5 g/mol NaCl)

Dilution

M_1V_1 = M_2V_2

Concentration × volume conserved.

Example: Dilute 500 mL of 2 M HCl to 4 M... wait, dilute means add water, so lower M. → Cannot dilute to a HIGHER concentration. Trick question.

Trap: Always convert Celsius to Kelvin for gas law problems. Forgetting this is Regents kryptonite.

Tip: STP = 273 K, 1 atm. One mole of ANY gas at STP = 22.4 L. Great for quick gas volume estimates.

Acids, bases, & redox

H⁺ vs OH⁻. Electrons lost vs gained.

Arrhenius definitions

Acid donates H⁺ in water. Base donates OH⁻. Strong = fully ionizes; weak = partially.

Example: HCl in water? → → H⁺ + Cl⁻ (strong acid)

pH scale

\text{pH} = -\log[\text{H}^+]

0–6 acidic, 7 neutral, 8–14 basic. Each unit = 10× change in [H⁺].

Example: [H⁺] = 10⁻⁴ M → pH = 4 (acidic)

Titration

M_aV_a = M_bV_b

For 1:1 acid-base. At equivalence point, moles acid = moles base.

Example: 25 mL 0.1 M HCl neutralized by 0.2 M NaOH → V_b = 12.5 mL

Redox

OIL RIG: Oxidation Is Loss of electrons, Reduction Is Gain. Track oxidation numbers.

Example: Zn + Cu²⁺ → Zn²⁺ + Cu → Zn oxidized (0 → +2), Cu reduced (+2 → 0)

Trap: pH goes DOWN as acidity goes UP. pH 3 is 100× MORE acidic than pH 5 (two units = 10²).

Tip: Half-cell notation: 'Zn|Zn²⁺||Cu²⁺|Cu.' Left is oxidation (anode). Right is reduction (cathode). 'An Ox' and 'Red Cat.'

Kinetics, equilibrium, thermodynamics

How fast, how far, how hot.

Rate factors

Speeds up with higher concentration, higher temperature, more surface area, added catalyst.

Example: Break a chunk of chalk into powder → Reaction speeds up (more surface area)

Le Châtelier

System at equilibrium shifts to OPPOSE any change. Add reactant → shifts right. Increase temp of endothermic → shifts right.

Example: N₂ + 3H₂ ⇌ 2NH₃ + heat. Add heat? → Shifts LEFT (exothermic — heat is a product)

Endo vs exo

Endothermic ΔH > 0 (absorbs heat, feels cold). Exothermic ΔH < 0 (releases heat, feels hot).

Example: Cold pack dissolving → Endothermic

Catalyst

Lowers activation energy. Speeds BOTH forward and reverse. Doesn't shift equilibrium.

Example: Add catalyst to equilibrium → Reaches equilibrium faster, position unchanged

Trap: A catalyst does NOT change ΔH or the equilibrium position — only the SPEED.

Tip: Draw a potential-energy diagram: reactants → activation energy hump → products. ΔH = products − reactants.

Frequently asked questions

How should I study for the Chemistry Regents?