Chemistry Guides

How to calculate molar mass (g/mol): add up each element's relative atomic mass times its number of atoms. Worked examples for H₂O, H₂SO₄, Ca(OH)₂, CuSO₄·5H₂O plus common mistakes.

Use the dilution formula C₁V₁=C₂V₂ to find the concentration or volume before/after dilution. Includes a worked example and safety notes.

How to balance chemical equations using conservation of mass by adjusting coefficients only. Inspection, lowest-common-multiple and odd/even tips, with worked examples.

How to calculate pH: pH = −log[H⁺]. Strong acids ionise fully; weak acids use [H⁺]≈√(Ka·c); for bases find pOH then subtract from 14. With examples.

Electronegativity is an atom's ability to attract bonding electrons (Pauling scale). It increases left-to-right across a period and decreases down a group; fluorine is highest. With trends and uses.

Solubility rules for common ionic compounds: group-1/ammonium and nitrates all soluble, chlorides except silver, sulfates except barium/lead, most carbonates and phosphates insoluble. With table and uses.

The metal activity series (K Ca Na Mg Al Zn Fe Sn Pb (H) Cu Hg Ag Pt Au) and its three uses: reactions with acids, with salt solutions, and with water/oxygen.

Common chemical equations organised by type — combination, decomposition, single and double displacement — covering combustion, metal-acid and double-displacement reactions for quick review.

Calculate molarity: first n = m/M for moles, then c = n/V for concentration (mol/L). With a worked example for preparing a solution.

How to use the ideal gas law PV=nRT: R=8.314 J/(mol·K) with P in kPa, V in L, T in K. Worked examples for volume, pressure and moles.

Periodic trends in atomic radius: it decreases left-to-right across a period and increases down a group. With reasons, ionic radii and links to electronegativity and metallic character.

The difference between acids, bases and salts: acids ionise to give only H⁺, bases give only OH⁻, salts are a metal (or ammonium) cation plus an acid anion. With properties and examples.

Differences between metals and nonmetals in periodic-table position, physical properties (lustre, conductivity, malleability) and chemistry (losing vs gaining electrons), plus metalloids.

A list of common acids, bases and salts with formulas: HCl, H₂SO₄, NaOH, NaCl, CaCO₃ and more, with strong/weak labels.

Mass percent has two common meanings: solute mass percent = solute mass ÷ solution mass ×100%; element mass percent in a compound = element mass ÷ molar mass ×100%. Worked examples for NaCl solution, %H in H₂O and more.

The limiting reagent is the reactant used up first; it caps how much product forms. Method: divide each reactant's moles by its stoichiometric coefficient — the smallest ratio is the limiting reagent. Worked examples for H₂ + O₂ and N₂ + H₂.

The mole (mol) is the hub of chemical calculations. Mass↔mol uses n=m/M; gas volume↔mol uses 22.4 L/mol at STP; particles↔mol uses Avogadro's number 6.022×10²³. Includes a conversion workflow and examples.

A reference table of common element valences and ion charges: K, Na, H, Ag are +1; Ca, Mg, Ba, Zn are +2; Al is +3; O is −2; plus common polyatomic ion charges. Used for writing formulas and balancing.

Rules for writing chemical formulas: the positive element/ion goes first, the negative second, total charge is zero, use the crossover method for subscripts, then simplify. Examples for Al₂(SO₄)₃, Fe₂O₃ and common mistakes.

The four basic reaction types: combination A+B→AB, decomposition AB→A+B, single displacement A+BC→AC+B, and double displacement (metathesis) AB+CD→AD+CB, each with a general form and typical equations.

Lab preparation of oxygen (heating KMnO₄ or catalytic decomposition of H₂O₂), carbon dioxide (limestone + dilute HCl) and hydrogen (zinc + dilute H₂SO₄): equations, collection methods and safety notes.

Ionic bonds form by transferring electrons, between a metal and a nonmetal (e.g. NaCl); covalent bonds form by sharing electrons, between nonmetals (e.g. H₂O). Compared by formation, electronegativity difference, melting point and conductivity.

The periodic law: element properties repeat periodically with atomic number. A summary of how atomic radius, electronegativity, metallic/nonmetallic character and ionization energy change across a period and down a group, and why.

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons — almost identical chemistry, different mass. Covers notation (¹²C, ¹⁴C), examples (protium/deuterium/tritium, carbon-12/13/14, uranium-235/238) and uses.

The noble (inert) gases are group 18: helium He, neon Ne, argon Ar, krypton Kr, xenon Xe, radon Rn. Their outer shells are full, so they are extremely unreactive. Their individual properties and real-world uses.

Radioactive elements have unstable nuclei that spontaneously decay and emit radiation. Every element with atomic number ≥ 84 (polonium) is radioactive, and technetium (43) and promethium (61) also have no stable isotopes. Common examples and rules.

The human body needs about two dozen essential elements. Major: oxygen, carbon, hydrogen, nitrogen, calcium, phosphorus, potassium, sulfur, sodium, chlorine, magnesium. Trace: iron, zinc, copper, manganese, iodine, selenium and more — with their roles.

Sodium (Na) and potassium (K) are both group-1 alkali metals: soft, reactive and vigorous with water. Comparing atomic mass, melting point, density, electronegativity and reactivity — potassium is more reactive than sodium.

Iron (Fe) and copper (Cu) are both period-4 transition metals. Comparing atomic mass, melting point, density, electronegativity, conductivity, magnetism and corrosion: iron is more reactive, magnetic and rusts; copper conducts well, resists corrosion and is antimicrobial.

A list of common organic compounds: methane CH₄, ethanol C₂H₅OH, acetic acid CH₃COOH, glucose C₆H₁₂O₆, methanol, ethylene, benzene and more — with formulas, classes and everyday uses.