High School Chemistry Essentials Cheatsheet

Foundations: Elements & Bonding

PERIODIC TABLE BASICS

Periods (Rows)	Horizontal rows (1-7). Indicate the number of electron shells an atom has.
Groups (Columns)	Vertical columns (1-18). Indicate the number of valence electrons (for main group elements) and similar chemical properties.
Metals	Left and center (Groups 1-12, parts of 13-16). Tend to lose electrons to form positive ions (cations). Shiny, malleable, ductile, good conductors.
Nonmetals	Upper right. Tend to gain electrons to form negative ions (anions) or share electrons. Dull, brittle, poor conductors.
Metalloids	Along the ‘staircase’ line (B, Si, Ge, As, Sb, Te). Exhibit properties of both metals and nonmetals.
Common Group Charges	Group 1 (Alkali Metals): +1 (e.g., Na⁺) Group 2 (Alkaline Earth Metals): +2 (e.g., Mg²⁺) Group 17 (Halogens): -1 (e.g., Cl⁻) Group 18 (Noble Gases): 0 (stable)
Memory Tip	Periods are like ROWS in a play, they go across. Groups are like COLUMNS supporting a roof, they go up and down.

CHEMICAL BONDING

Ionic Bonds	Between a metal and a nonmetal. Involves the transfer of electrons from metal to nonmetal, forming ions (cations and anions) which are attracted electrostatically. Example: `NaCl` (Na loses 1e⁻ to Cl)
Covalent Bonds	Between two nonmetals. Involves the sharing of electrons to achieve a stable electron configuration (usually an octet). Example: `H₂O` (O shares electrons with 2 H atoms)
Metallic Bonds	Between metal atoms. Valence electrons are delocalized and form a ‘sea of electrons’ that can move freely, explaining metal properties like conductivity.
Octet Rule	Atoms tend to gain, lose, or share electrons to achieve a stable configuration of eight valence electrons (like noble gases). Hydrogen follows the ‘duet rule’ (2 electrons).
Lewis Dot Structures	Diagrams that show the valence electrons of atoms as dots around the element symbol. Used to illustrate bonding and non-bonding electron pairs.
Common Mistake	Confusing electron transfer (ionic) with electron sharing (covalent). Remember, ‘ionic’ sounds like ‘ions’ which are formed by transfer!
Polar vs. Nonpolar Covalent	Nonpolar: Equal sharing of electrons (e.g., `O₂`, `Cl₂`). Polar: Unequal sharing of electrons, creating partial charges (e.g., `H₂O` - Oxygen pulls electrons stronger).

Reactions, Matter & Gases

CHEMICAL REACTIONS

1. Synthesis (Combination)
Two or more reactants combine to form a single, more complex product.
A + B → AB
Example: 2Na(s) + Cl₂(g) → 2NaCl(s)

2. Decomposition
A single compound breaks down into two or more simpler substances.
AB → A + B
Example: 2H₂O(l) → 2H₂(g) + O₂(g)

3. Single Replacement (Displacement)
One element replaces another element in a compound.
A + BC → AC + B
Example: Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)

4. Double Replacement (Displacement)
The positive ions (cations) of two ionic compounds swap places, forming two new compounds (often one is a precipitate, gas, or water).
AB + CD → AD + CB
Example: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

5. Combustion
A rapid reaction with oxygen, usually producing heat and light. For hydrocarbons, products are always carbon dioxide and water.
Hydrocarbon + O₂ → CO₂ + H₂O
Example: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)

Balancing Chemical Equations

Law of Conservation of Mass: Atoms are neither created nor destroyed.
Steps:
1. Count atoms of each element on both sides.
2. Balance metals first, then nonmetals (excluding H & O).
3. Balance Oxygen (O).
4. Balance Hydrogen (H).
5. Check all elements. Coefficients must be lowest whole numbers.
  Example: _C₂H₆ + _O₂ → _CO₂ + _H₂O becomes 2C₂H₆ + 7O₂ → 4CO₂ + 6H₂O

Memory Tip:
‘Subscripts STAY, Coefficients CHANGE!’ You change the number of molecules (coefficients), not the formula of the molecules (subscripts).

STATES OF MATTER & GAS LAWS

Solids
Liquids
Gases
Boyle’s Law (P vs. V)
At constant temperature and moles, pressure and volume are inversely proportional. If pressure increases, volume decreases.
Charles’s Law (V vs. T)
At constant pressure and moles, volume and temperature (in Kelvin) are directly proportional. If temperature increases, volume increases.
Ideal Gas Law
Relates Pressure (P), Volume (V), moles (n), and Temperature (T). R = Ideal Gas Constant (0.0821 \frac{\text{L} \cdot \text{atm}}{\text{mol} \cdot \text{K}}). Units are crucial! P in atm, V in L, T in K.
Common Mistake:

Acids, Moles & Calculations

ACIDS, BASES & pH

pH Scale	Measures the acidity or basicity of a solution, from 0 to 14. 0-6.9: Acidic 7.0: Neutral 7.1-14: Basic (Alkaline)
pH Calculation	`pH = -log[H⁺]` `pOH = -log[OH⁻]` `pH + pOH = 14`
Acid Indicators	Substances that change color depending on the pH of the solution. Litmus Paper: Red in acid, blue in base. Phenolphthalein: Colorless in acid, pink in base.
Arrhenius Definition	Acid: Produces `H⁺` ions when dissolved in water (e.g., `HCl → H⁺ + Cl⁻`) Base: Produces `OH⁻` ions when dissolved in water (e.g., `NaOH → Na⁺ + OH⁻`)
Brønsted-Lowry Definition	Acid: A proton (`H⁺`) donor (e.g., `HCl` in `HCl + H₂O → H₃O⁺ + Cl⁻`) Base: A proton (`H⁺`) acceptor (e.g., `NH₃` in `NH₃ + H₂O ⇌ NH₄⁺ + OH⁻`)
Conjugate Acid-Base Pairs	When a Brønsted-Lowry acid donates a proton, it forms its conjugate base. When a Brønsted-Lowry base accepts a proton, it forms its conjugate acid. Example: `HCl (acid) + H₂O (base) ⇌ Cl⁻ (conjugate base) + H₃O⁺ (conjugate acid)`
Memory Tip:	Brønsted-Lowry is broader. Think of it like a `give and take` (donor/acceptor) of protons, rather than just what they produce in water.

MOLES & STOICHIOMETRY

The Mole (mol)	The SI unit for amount of substance. A mole of any substance contains Avogadro’s number of particles.
Avogadro’s Number	`6.022 x 10²³` This is the number of particles (atoms, molecules, ions, formula units) in exactly one mole of a substance.
Molar Mass (g/mol)	The mass of one mole of a substance. Numerically equal to the atomic mass (for elements) or formula/molecular mass (for compounds) in grams. Example: Molar mass of `H₂O` = `(2 x 1.008) + 16.00 = 18.016 g/mol`
Mole-to-Mass Conversion	`Moles x Molar Mass = Grams` `Grams / Molar Mass = Moles` Example: How many grams are in 0.5 mol of `H₂O`? `0.5 mol x 18.016 g/mol = 9.008 g`
Mole-to-Particle Conversion	`Moles x Avogadro's Number = Particles` `Particles / Avogadro's Number = Moles` Example: How many molecules in 2 mol of `CO₂`? `2 mol x 6.022 x 10²³ molecules/mol = 1.2044 x 10²⁴ molecules`
Mole-to-Mole Conversion (Stoichiometry)	Uses the mole ratio from a balanced chemical equation to convert between moles of different substances. Example: `2H₂ + O₂ → 2H₂O` If you have 4 moles of `H₂`, how many moles of `H₂O` are produced? `4 mol H₂ x (2 mol H₂O / 2 mol H₂) = 4 mol H₂O`
Common Mistake:	Forgetting to balance the chemical equation before doing any mole-to-mole (stoichiometry) calculations. The coefficients are the key!