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Chemistry tables

All 118 elements, ions, compounds, mole calculations, bonding, equilibrium and organic functional groups.

Common ions

Ion	Formula	Charge
Sodium	Na⁺	+1
Calcium	Ca²⁺	+2
Ammonium	NH₄⁺	+1
Chloride	Cl⁻	−1
Hydroxide	OH⁻	−1
Nitrate	NO₃⁻	−1
Sulfate	SO₄²⁻	−2
Carbonate	CO₃²⁻	−2

A positive ion has lost electrons and a negative ion has gained electrons. Combine ions in ratios that balance total charge.

Worked example

One Ca²⁺ and two Cl⁻ form CaCl₂. Two Na⁺ and one SO₄²⁻ form Na₂SO₄.

Amount of substance

n = mass / molar mass. Use grams and g/mol to obtain moles.

Molar mass is the mass per mole. It differs from atomic number and needs to be calculated for the full chemical formula.

Worked example

With NaCl molar mass approximated as 58.44 g/mol, 5.844 g contains 0.1000 mol.

Concentration

Molar concentration = moles of solute / litres of solution.

Use the final solution volume in litres, not the initial volume of solvent alone. Molarity is measured in mol/L.

Worked example

Dissolving 0.2 mol to make 0.5 L of solution gives 0.4 mol/L.

Balancing equations

Change coefficients, never chemical subscripts. Example: 2H₂ + O₂ → 2H₂O.

Count atoms of each element on both sides. Use whole-number coefficients in the smallest suitable ratio.

Worked example

2H₂ + O₂ → 2H₂O has four hydrogen atoms and two oxygen atoms on each side.

pH

At school-level dilute-solution approximation, pH = −log₁₀[H⁺]. Neutral water is approximately pH 7 at 25 °C.

The concentration formula is an introductory approximation; the thermodynamic definition uses hydrogen ion activity. Temperature affects neutral pH.

Worked example

For an ideal dilute solution with [H⁺]=10⁻³ mol/L, pH≈3. A one-unit pH change corresponds to a tenfold activity ratio.

Common laboratory relationships

Relationship	Meaning and conditions	Example
c₁V₁=c₂V₂	Dilution with the same solute and conserved amount; use matching volume units.	100 mL of 1 M stock diluted to 500 mL gives 0.2 M.
PV=nRT	Ideal gas approximation. Use absolute temperature in kelvin and a consistent gas constant.	For 1 mol at 300 K and 100,000 Pa with R≈8.314, V≈0.02494 m³.
N=nNₐ	Particle count equals moles times Avogadro’s constant, 6.02214076×10²³ mol⁻¹.	0.5 mol contains 3.01107038×10²³ specified particles.

Atomic structure and isotopes

Atomic number Z is proton count. Mass number A is protons plus neutrons in one isotope. A neutral atom has Z electrons; cations have fewer and anions more. Isotopes share Z but differ in neutron count. Average atomic mass reflects isotope abundances.

Sodium-23 has 11 protons and 12 neutrons. Neutral sodium has 11 electrons; Na⁺ has 10. Orbital subshell capacities are s:2, p:6, d:10 and f:14. Each orbital holds at most two opposite-spin electrons.

All 118 elements: atomic number, symbol and name

Atomic numbers below count protons, not atomic mass. Chemical symbols are case-sensitive: Co is cobalt, while CO is carbon monoxide. Use the search box to locate an element.

Atomic number	Symbol	Element
1	H	Hydrogen
2	He	Helium
3	Li	Lithium
4	Be	Beryllium
5	B	Boron
6	C	Carbon
7	N	Nitrogen
8	O	Oxygen
9	F	Fluorine
10	Ne	Neon
11	Na	Sodium
12	Mg	Magnesium
13	Al	Aluminium
14	Si	Silicon
15	P	Phosphorus
16	S	Sulfur
17	Cl	Chlorine
18	Ar	Argon
19	K	Potassium
20	Ca	Calcium
21	Sc	Scandium
22	Ti	Titanium
23	V	Vanadium
24	Cr	Chromium
25	Mn	Manganese
26	Fe	Iron
27	Co	Cobalt
28	Ni	Nickel
29	Cu	Copper
30	Zn	Zinc
31	Ga	Gallium
32	Ge	Germanium
33	As	Arsenic
34	Se	Selenium
35	Br	Bromine
36	Kr	Krypton
37	Rb	Rubidium
38	Sr	Strontium
39	Y	Yttrium
40	Zr	Zirconium
41	Nb	Niobium
42	Mo	Molybdenum
43	Tc	Technetium
44	Ru	Ruthenium
45	Rh	Rhodium
46	Pd	Palladium
47	Ag	Silver
48	Cd	Cadmium
49	In	Indium
50	Sn	Tin
51	Sb	Antimony
52	Te	Tellurium
53	I	Iodine
54	Xe	Xenon
55	Cs	Caesium
56	Ba	Barium
57	La	Lanthanum
58	Ce	Cerium
59	Pr	Praseodymium
60	Nd	Neodymium
61	Pm	Promethium
62	Sm	Samarium
63	Eu	Europium
64	Gd	Gadolinium
65	Tb	Terbium
66	Dy	Dysprosium
67	Ho	Holmium
68	Er	Erbium
69	Tm	Thulium
70	Yb	Ytterbium
71	Lu	Lutetium
72	Hf	Hafnium
73	Ta	Tantalum
74	W	Tungsten
75	Re	Rhenium
76	Os	Osmium
77	Ir	Iridium
78	Pt	Platinum
79	Au	Gold
80	Hg	Mercury
81	Tl	Thallium
82	Pb	Lead
83	Bi	Bismuth
84	Po	Polonium
85	At	Astatine
86	Rn	Radon
87	Fr	Francium
88	Ra	Radium
89	Ac	Actinium
90	Th	Thorium
91	Pa	Protactinium
92	U	Uranium
93	Np	Neptunium
94	Pu	Plutonium
95	Am	Americium
96	Cm	Curium
97	Bk	Berkelium
98	Cf	Californium
99	Es	Einsteinium
100	Fm	Fermium
101	Md	Mendelevium
102	No	Nobelium
103	Lr	Lawrencium
104	Rf	Rutherfordium
105	Db	Dubnium
106	Sg	Seaborgium
107	Bh	Bohrium
108	Hs	Hassium
109	Mt	Meitnerium
110	Ds	Darmstadtium
111	Rg	Roentgenium
112	Cn	Copernicium
113	Nh	Nihonium
114	Fl	Flerovium
115	Mc	Moscovium
116	Lv	Livermorium
117	Ts	Tennessine
118	Og	Oganesson

Periodic trends and common valencies

Across a period atomic radius generally decreases while ionisation energy and electronegativity generally rise, with exceptions. Down a group radius generally increases. Group 1 metals commonly form +1 ions; group 2 form +2; aluminium commonly +3; halides commonly −1. Transition elements often have multiple oxidation states.

Species	Charge or common oxidation state	Example
Magnesium	+2	MgCl₂
Aluminium	+3	Al₂O₃
Iron	+2 or +3	FeCl₂ or FeCl₃
Copper	+1 or +2	Cu₂O or CuO
Zinc	+2	ZnSO₄
Phosphate	−3	PO₄³⁻
Hydrogen carbonate	−1	HCO₃⁻
Permanganate	−1	MnO₄⁻
Dichromate	−2	Cr₂O₇²⁻

Combine ions to make overall charge zero: Al³⁺ and SO₄²⁻ form Al₂(SO₄)₃. Parentheses preserve the polyatomic ion. Oxidation state is formal electron bookkeeping and need not be an actual ionic charge.

Common compounds and school molar masses

Compound	Formula	Approximate molar mass (g/mol)
Water	H₂O	18.02
Carbon dioxide	CO₂	44.01
Sodium chloride	NaCl	58.44
Calcium carbonate	CaCO₃	100.09
Sulfuric acid	H₂SO₄	98.08
Hydrochloric acid solute	HCl	36.46
Sodium hydroxide	NaOH	40.00
Ammonia	NH₃	17.03
Glucose	C₆H₁₂O₆	180.16
Baking soda	NaHCO₃	84.01
Washing soda	Na₂CO₃·10H₂O	286.14
Gypsum	CaSO₄·2H₂O	172.17
Plaster of Paris	CaSO₄·½H₂O	145.15

Use atomic weights supplied in your question when they differ slightly. Hydrated salts include water in the formula mass. 9.01 g water is approximately 0.500 mol. A solution does not have one fixed molar mass like a pure compound.

Stoichiometry and limiting reagent

moles = mass / molar mass; mole ratios come from balanced coefficients

Balance the equation, convert reactant amounts to moles, divide each amount by its coefficient and identify the smallest ratio. This reactant limits the theoretical product. Percentage yield=100×actual/theoretical yield on the same basis.

Worked example

For 2H₂+O₂→2H₂O, 3 mol H₂ and 2 mol O₂ give ratios 1.5 and 2. H₂ limits; 3 mol water can form and 0.5 mol O₂ remains.

Common mistake: Comparing masses directly instead of coefficient-adjusted mole amounts.

Concentration units and dilution

M=nsolute/Vsolution(L); molality=nsolute/msolvent(kg)

Mass%=100×solute mass/solution mass. Mole fraction xᵢ=nᵢ/Σn. Mass-based ppm=10⁶×mass fraction; in very dilute aqueous solutions of density near 1 kg/L, mg/L is approximately ppm. Dilution conserves solute amount: M₁V₁=M₂V₂.

Worked example

5.844 g NaCl is 0.1 mol; make up to 0.5 L solution for 0.2 M. Diluting 50 mL of 2 M solution to 250 mL gives 0.4 M.

Common mistake: Using solvent volume instead of final solution volume for molarity.

Acids, bases, indicators and salts

A Brønsted acid donates a proton and a base accepts one. Strong versus weak describes extent of ionisation; concentrated versus dilute describes amount per volume. For dilute ideal aqueous solutions at 25°C, pH+pOH≈14. Neutral pH changes with temperature.

Indicator	Acidic condition	Basic condition
Litmus	Red	Blue
Phenolphthalein	Colourless	Pink in its alkaline transition range
Methyl orange	Red	Yellow
Universal indicator	Red/orange/yellow depending on pH	Blue/purple depending on pH

HCl+NaOH→NaCl+H₂O is neutralisation. For a dilute ideal 0.001 M strong monoprotic acid, pH≈3. Salt solutions are not always neutral: Na₂CO₃ solutions are alkaline, while NH₄Cl solutions are acidic.

Chemical bonding and molecular shape

Ionic bonding: electrostatic attraction; covalent bonding: shared electron pairs

Metallic bonding involves delocalised electrons. Molecular geometry depends on bonding and lone electron pairs: methane is tetrahedral, ammonia trigonal pyramidal and water bent. A polar bond does not guarantee a polar molecule.

Worked example

CO₂ is linear, so its two C=O bond dipoles cancel. H₂O is bent, so its bond dipoles do not cancel.

Common mistake: Describing every covalent molecule as non-polar.

Oxidation, reduction and electrochemistry

Oxidation loses electrons; reduction gains electrons

An oxidising agent is reduced and a reducing agent is oxidised. Anode is the site of oxidation and cathode of reduction in both galvanic and electrolytic cells. Electrode signs differ between these cell types.

Worked example

Zn+Cu²⁺→Zn²⁺+Cu: zinc loses two electrons and is the reducing agent. Copper ions gain electrons. For a galvanic cell, E°cell=E°cathode−E°anode using reduction potentials.

Common mistake: Assuming the anode is always negative; an electrolytic anode is positive.

Equilibrium and reaction quotient

For aA+bB⇌cC+dD: Kc=[C]ᶜ[D]ᵈ/([A]ᵃ[B]ᵇ)

Use equilibrium concentrations; omit pure solids and liquids. Reaction quotient Q uses current values. Q<K favours forward net reaction and Q>K favours reverse. K changes with temperature; a catalyst accelerates approach to equilibrium without changing K.

Worked example

For N₂+3H₂⇌2NH₃, Kc=[NH₃]²/([N₂][H₂]³). Higher pressure favours fewer gas moles, here the product side.

Common mistake: Changing K merely because a reactant was added at fixed temperature.

Thermochemistry and reaction rates

ΔHreaction=ΣνΔHf(products)−ΣνΔHf(reactants)

Exothermic reactions have negative ΔH; endothermic positive. Hess’s law adds enthalpy changes along a path. Rate law rate=k[A]ᵐ[B]ⁿ has orders determined experimentally, not generally by coefficients of the overall equation.

Worked example

If product formation enthalpies sum to −500 kJ and reactants to −300 kJ for the stated reaction, ΔH=−200 kJ. A catalyst lowers the effective activation barrier.

Common mistake: Assuming a favourable enthalpy alone guarantees a rapid reaction.

Organic functional groups

Class	Group	Example
Alkane	C–C single bonds	Ethane C₂H₆
Alkene	C=C	Ethene C₂H₄
Alkyne	C≡C	Ethyne C₂H₂
Alcohol	–OH	Ethanol C₂H₅OH
Aldehyde	–CHO	Ethanal CH₃CHO
Ketone	>C=O	Propanone CH₃COCH₃
Carboxylic acid	–COOH	Ethanoic acid CH₃COOH
Ester	–COO–	Ethyl ethanoate CH₃COOC₂H₅
Amine	–NH₂ (primary)	Methylamine CH₃NH₂

Homologous members share a functional group and successive members commonly differ by CH₂. Isomers have the same molecular formula but different arrangement. Saturated open-chain alkanes follow CₙH₂ₙ₊₂; do not apply that formula to rings.

Solubility and reaction recognition

At introductory level, alkali-metal and ammonium salts and most nitrates are soluble in water. Many chlorides are soluble, with important exceptions such as AgCl and PbCl₂. Many carbonates and hydroxides are poorly soluble except relevant alkali-metal and ammonium salts. Temperature and the particular substance matter.

Reaction type	Example
Combination	2Mg+O₂→2MgO
Decomposition	CaCO₃→CaO+CO₂ (heating)
Displacement	Zn+CuSO₄→ZnSO₄+Cu
Precipitation	AgNO₃+NaCl→AgCl↓+NaNO₃
Acid–carbonate	CaCO₃+2HCl→CaCl₂+H₂O+CO₂
Combustion	CH₄+2O₂→CO₂+2H₂O

Equations explain classroom concepts. Carry out practical chemistry only with appropriate instruction and laboratory supervision.

Practice with answers

How many neutrons in carbon-14?

Z=6, so 14−6=8 neutrons.

Formula from Mg²⁺ and NO₃⁻?

Mg(NO₃)₂ balances the charges.

Moles in 22.005 g CO₂?

Using 44.01 g/mol gives 0.500 mol.

pH for ideal dilute [H⁺]=10⁻⁴ mol/L?

pH≈4.

Which is oxidised in Mg+2H⁺→Mg²⁺+H₂?

Mg loses two electrons and is oxidised.

Does a catalyst change the equilibrium constant?

No. It changes reaction rates, not K at a fixed temperature.

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