o Change anion ending:
▪ –ate becomes –ic acid.
1. Ionic Compounds (Metal + Nonmetal): ▪ H₂CO₃ (contains carbonate) → Carbonic acid
• Name the metal first.
• Name the nonmetal next, but change its ending to –ide. List of Mentioned Compounds
o Example: K₂O → Potassium oxide Formula Compound Name Type
• If the metal can have more than one charge, use Roman numerals:
K₂O Potassium oxide Ionic
o FeCl₂ → Iron (II) chloride
FeCl₂ Iron (II) chloride Ionic (Variable)
o FeCl₃ → Iron (III) chloride
FeCl₃ Iron (III) chloride Ionic (Variable)
2. Polyatomic Ions: SF₆ Sulfur hexafluoride Molecular
• Some compounds have groups of atoms called polyatomic ions (like N₂O₄ Dinitrogen tetroxide Molecular
SO₄²⁻ or NO₃⁻).
HCl Hydrochloric acid Binary acid
• You must memorize their names, like carbonate (CO₃²⁻) or nitrate
(NO₃⁻). H₂CO₃ Carbonic acid Oxyacid
3. Molecular Compounds (Nonmetal + Nonmetal): Common Prefixes for Molecular Compounds
• Use prefixes to tell how many atoms are present. Number Prefix
o Example: SF₆ → Sulfur hexafluoride
1 mono-
o Example: N₂O₄ → Dinitrogen tetroxide
2 di-
3 tri-
4. Acids:
• Binary acids (with just hydrogen and one other element): 4 tetra-
o Use “hydro-” + element name with “–ic” + acid. 5 penta-
▪ HCl → Hydrochloric acid 6 hexa-
• Oxyacids (contain hydrogen, oxygen, and another element): 7 hepta-
Number Prefix Formula Name Type
8 octa- CO Carbon monoxide Molecular
9 nona- CO₂ Carbon dioxide Molecular
10 deca- NO Nitric oxide Molecular
NO₂ Nitrogen dioxide Molecular
Ionic Compounds (Metal + Nonmetal or Polyatomic Ion) PCl₃ Phosphorus trichloride Molecular
Formula Name Type PCl₅ Phosphorus pentachloride Molecular
NaCl Sodium chloride Ionic Cl₂O₇ Dichlorine heptoxide Molecular
CaBr₂ Calcium bromide Ionic BCl₃ Boron trichloride Molecular
MgO Magnesium oxide Ionic SiF₄ Silicon tetrafluoride Molecular
Al₂S₃ Aluminum sulfide Ionic
Li₃N Lithium nitride Ionic Acids
CuCl Copper(I) chloride Ionic (variable) Binary Acids (Hydrogen + Nonmetal)
CuCl₂ Copper (II) chloride Ionic (variable) Formula Name Type
PbO Lead (II) oxide Ionic (variable) HF Hydrofluoric acid Binary Acid
PbO₂ Lead (IV) oxide Ionic (variable) HCl Hydrochloric acid Binary Acid
AgNO₃ Silver nitrate Ionic (polyatomic) HBr Hydrobromic acid Binary Acid
CaCO₃ Calcium carbonate Ionic (polyatomic) HI Hydroiodic acid Binary Acid
Na₂SO₄ Sodium sulfate Ionic (polyatomic) H₂S Hydrosulfuric acid Binary Acid
KNO₂ Potassium nitrite Ionic (polyatomic) Oxyacids (Hydrogen + Polyatomic Ion)
Ba (OH)₂ Barium hydroxide Ionic (polyatomic) Formula Name Type
HNO₃ Nitric acid Oxyacid
Molecular Compounds (Nonmetal + Nonmetal)
Formula Name Type Formula Name Type
HNO₂ Nitrous acid Oxyacid hydroxide
H₂SO₄ Sulfuric acid Oxyacid
Name
H₂SO₃ Sulfurous acid Oxyacid
H₃PO₄ Phosphoric acid Oxyacid phosphite
H₃PO₃ Phosphorous acid Oxyacid
phosphate
H₂CO₃ Carbonic acid Oxyacid
hydrogen phosphate
HClO₃ Chloric acid Oxyacid
HClO₂ Chlorous acid Oxyacid dihydrogen phosphate
Name
perchlorate
ammonium
chlorate
nitrite
chlorite
nitrate
hypochlorite
sulfite
bromate
sulfate
iodate
hydrogen sulfate
(aka: bisulfate) Name Formula
thiosulfate acetate CH3COO−
acetate C2H3O2−
oxalate
carbonate CO32−
Name Formula IUPAC Name older name Formula
hydrogen carbonate HCO3− lead(IV) plumbic Pb4+
(aka: bicarbonate)
mercury(I) mercurous Hg22+
2−
chromate CrO4
mercury(II) mercuric Hg2+
dichromate Cr2O72−
tin(II) stannous Sn2+
permanganate MnO4−
tin(IV) stannic Sn4+
2−
peroxide O2
1. Monoatomic Elements
cyanide CN− • These are made of only one atom.
• They are stable on their own and don't need to bond with other atoms.
cyanate OCN−
• Examples: Helium (He), Neon (Ne), Argon (Ar) – these are all noble
gases.
thiocyanate SCN−
Think of it like: One single Lego piece that doesn’t need to connect to
Some Metallic Cations others.
IUPAC Name older name Formula
2. Diatomic Elements
copper(I) cuprous Cu+ • These are made of two atoms of the same element bonded together.
• They are found this way in nature because a single atom of these
copper(II) cupric Cu2+ elements isn’t stable on its own.
• Examples:
iron(II) ferrous Fe2+
o Hydrogen gas (H₂)
iron(III) ferric Fe3+ o Nitrogen gas (N₂)
o Fluorine (F₂), Chlorine (Cl₂) – both are halogens.
2+
lead(II) plumbous Pb
Think of it like: Two identical Lego pieces snapped together to be stable.
3. Diatomic Ions Pure Element vs. Compound
• These also have two atoms, but they carry an electric charge. • A pure element has only one type of atom.
• Examples: Examples:
o Peroxide ion (O₂²⁻) o Helium (He) – one helium atom
o Superoxide ion (O₂⁻) o Oxygen gas (O₂) – two oxygen atoms
o Mercury(I) ion (Hg₂²⁺) o Neon (Ne) – single neon atoms
• A compound has different types of atoms bonded together.
Key Difference: Diatomic elements are neutral, but diatomic ions have a
charge. Examples:
o Water (H₂O) – hydrogen and oxygen
4. Triatomic and Polyatomic Ions o Carbon dioxide (CO₂) – carbon and oxygen
• Triatomic ions have three atoms, and polyatomic ions have many
atoms.
What is a Molecule?
• These are charged groups of atoms that act like one unit in chemical
• A molecule forms when atoms bond.
reactions.
Example:
• Examples:
o O₂ is a molecule made of two oxygen atoms.
o Triiodide ion (I₃⁻) – 3 iodine atoms
o H₂O is a compound molecule because it has different types of
o Sulfate (SO₄²⁻) – 1 sulfur + 4 oxygen atoms atoms.
o Carbonate (CO₃²⁻) – 1 carbon + 3 oxygen atoms
o Disulfide ion (S₂²⁻) – 2 sulfur atoms Atoms vs. Ions
Think of it like: A team of Lego blocks joined together with a charge, • Atoms are neutral (same number of protons and electrons).
working as one in chemical reactions
• Ions have a charge because they gained or lost electrons.
Atom vs. Molecule
o Cation = positive ion (lost electrons)
• An atom is a single particle of an element.
o Anion = negative ion (gained electrons)
Example: Helium (He) is made of single atoms.
Example:
• A molecule is made of two or more atoms bonded together.
▪ Phosphorus-31 becomes P³⁻ when it gains 3 electrons.
Example: Hydrogen gas (H₂) has two hydrogen atoms joined
together.
How to Find the Number of Electrons To understand atoms, you need to know how to calculate the number of
protons, neutrons, and electrons. Let’s use sodium as an example.
• Electrons = Atomic Number − Charge
Example: Aluminum (atomic number 13)
o As an atom: 13 electrons Atomic Number = Protons
o As a 3+ ion: 13 − 3 = 10 electrons • The atomic number tells us how many protons an atom has.
• It also tells us what element it is.
Types of Compounds Sodium’s atomic number is 11, so it always has 11 protons.
• Ionic Compounds: Formed from a metal and a nonmetal
o Electrons are transferred Mass Number vs. Atomic Mass
Examples:
• The mass number is the total of protons + neutrons.
▪ Sodium chloride (NaCl)
• It’s different from the average atomic mass on the periodic table
▪ Calcium oxide (CaO) (which is a weighted average of all isotopes).
• Covalent (Molecular) Compounds: Made of nonmetals The most common isotope of sodium is Sodium-23 (11 protons + 12
neutrons).
o Electrons are shared
Examples:
▪ Water (H₂O) Isotopes
▪ Sulfur hexafluoride (SF₆) • Isotopes are atoms of the same element with different numbers of
neutrons.
Tip: Use the periodic table:
Sodium-22 and Sodium-24 are less common and less stable than Sodium-
• Metals = left side
23.
• Nonmetals = upper right side
Neutrons = Mass Number – Atomic Number
Exceptions
To find the number of neutrons, subtract the atomic number from the mass
• Ammonium chloride (NH₄Cl): number:
o Made of nonmetals but is still ionic because it has charged Sodium-23:
ions. 23 (mass number) – 11 (atomic number) = 12 neutrons
Understanding Atomic Structure: Protons, Neutrons, and Electrons Silicon-29:
29 – 14 = 15 neutrons
Average Atomic Mass=∑ (isotope mass×relative abundance) \text {Average
Electrons Atomic Mass} = \sum (\text {isotope mass} \times \text {relative abundance})
Average Atomic Mass=∑ (isotope mass×relative abundance)
• In neutral atoms, the number of electrons = protons.
Step-by-step Example:
Neutral sodium has 11 protons and 11 electrons.
Suppose an element has two isotopes:
• In ions, the number of electrons changes: • Isotope A: mass = 10.01 u, abundance = 20% (0.20)
o If the atom loses electrons, it becomes a positive ion (cation). • Isotope B: mass = 11.01 u, abundance = 80% (0.80)
o If the atom gains electrons, it becomes a negative ion (anion). Average Atomic Mass= (10.01×0.20) +(11.01×0.80)
=2.002+8.808=10.81 u\text {Average Atomic Mass} = (10.01 \times 0.20) +
Aluminum³⁺ has 13 protons but only 10 electrons (13 – 3).
(11.01 \times 0.80) = 2.002 + 8.808 = 10.81 \, \text{u}Average Atomic Mass=
(10.01×0.20) +(11.01×0.80) =2.002+8.808=10.81u
How to Identify an Element
So, the average atomic mass of the element is 10.81 u.
• You can find an element by calculating its atomic number, which
equals its number of protons.
What are isotopes? Isotopes are substances that are composed of the same
An atom with 22 protons is Titanium. element but consist of different mass numbers and number of neutrons. They
• To find the mass number, add protons + neutrons: share the same atomic number and therefore the same number of protons. This
Titanium: 22 protons + 26 neutrons = 48 (mass number) video contains plenty of examples and practice problems associated with
isotopes and writing the atomic symbol for it.
Key Formulas:
Ion Fundamentals
• Protons = Atomic Number 1. An ion is a particle with unequal numbers of electrons and protons,
• Neutrons = Mass Number – Atomic Number resulting in an electrical charge, while atoms are electrically
• Electrons (neutral atom) = Protons neutral due to equal numbers of protons and electrons.
• Electrons (ion) = Atomic Number – Charge (for cations) or + Charge 2. Positively charged ions are called cations, having more protons than
(for anions) electrons, while negatively charged ions are called anions, having
more electrons than protons.
To calculate the average atomic mass of an element, you need to know the Atomic Structure and Charge
mass and relative abundance of each of its isotopes. Here's the formula:
3. The atomic number equals the number of protons, the mass number is o Reduction: Gain of electrons.
the sum of protons and neutrons, and the charge is the difference o Dehydrogenation: Removal of hydrogen from a molecule.
between protons and electrons. 4. States of Matter & Phase Changes
4. To write an atomic symbol of an ion, identify the element using the o Sublimation: Solid to gas without becoming liquid.
atomic number, include the mass number, and add the charge in o Condensation: Gas to liquid.
parentheses. o Melting: Solid to liquid.
Calculations o Evaporation: Liquid to gas.
5. The net charge of an ion is calculated by subtracting the number 5. Atomic and Molecular Structure
of electrons from the number of protons (e.g., 29 protons - 27 electrons o Atoms: Basic units of elements.
= +2 charge). o Molecules: Two or more atoms bonded together.
6. The number of electrons in an ion can be determined by subtracting o Chemical symbols: Abbreviations for elements (e.g., Au for
the net charge from the atomic number (e.g., 32 protons - 4 charge = gold, Ag for silver).
28 electrons). o Chemical formulas: Representation of compounds (e.g., H₂O
Core Chemistry Concepts for water).Studocu+1Studocu+1Studocu
1. Chemical Bonding 6. Quantitative Chemistry
o Covalent bonds: Equal sharing of electrons between atoms. o The mole: Unit representing 6.022 × 10²³ particles.
o Ionic bonds: Transfer of electrons from one atom to another. o Molar mass: Mass of one mole of a substance.
o Hydrogen bonds: Weak attractions between polar molecules. 7. Periodic Table & Elemental Properties
o Coordinate covalent bonds: Both shared electrons come from o Abundance of elements (e.g., nitrogen as the most abundant
the same atom.Studocu+2Studocu+2Studocu+2 gas in Earth's atmosphere).
2. Chemical Reactions o Chemical symbols and their corresponding elements.Studocu
o Synthesis: Combining two or more substances to form a new 8. Separation Techniques
compound. o Methods to separate mixtures based on physical properties
o Decomposition: Breaking down a compound into simpler (e.g., filtration, distillation).
substances.
o Displacement: An element replaces another in a compound.
o Neutralization: Acid and base react to form water and a salt.
3. Redox Processes
o Oxidation: Loss of electrons.
