Naming Binary Compounds

A binary compound is a chemical made of two different elements.
Naming binary compounds is the first step in learning to name any chemical.
This video covers naming:
- Simple ionic compounds
- Multivalent ionic compounds
- Simple covalent compounds
The video does not cover polyatomic ions or organic compounds.

The more metallic or less electronegative element is written first in both the chemical name and formula (e.g., $CaCl_2$ ).
The metallic nature of an element can be determined by its position on the periodic table.
- Metals are typically to the left and towards the bottom of the periodic table.
The first element's name remains unchanged, while the ending of the second element changes to "ide".
- Example: Calcium chloride ( $CaCl2$ ) and Carbon dioxide ( $CO2$ ).

Ionic compounds:
- Typically made of a metal and a nonmetal (e.g., calcium chloride).
- Do not use prefixes.
Covalent compounds:
- Composed of only nonmetals or metalloids (e.g., carbon dioxide).
- Use prefixes to indicate the numbers in the formula.

Common prefixes for covalent compounds:
- 1: mono
- 2: di
- 3: tri
- 4: tetra
- 5: penta
- 6: hexa
- 7: hepta
- 8: octa
- 9: nona
- 10: deca
The "mono" prefix is not used for the first element (e.g., $CO_2$ is carbon dioxide, not monocarbon dioxide).
Mono is used for the second element (e.g. $N_2O$ is dinitrogen monoxide).
An 'a' or 'o' at the end of a prefix is dropped if followed by 'oxide' (e.g., $P2O4$ is diphosphorus tetroxide, not diphosphorus tetraoxide).
- Example: $B2S3$ is diboron trisulfide.
- Example: Arsenic pentachloride is $AsCl_5$ .

Ionic compounds generally do not use prefixes (e.g., $Li3N$ is lithium nitride, $Al2S_3$ is aluminum sulfide).
The octet rule helps determine the formula from the name.
- Main group elements tend to gain or lose electrons to have the same number of electrons as the nearest noble gas (typically 8 valence electrons).
- Elements with 1, 2, or 3 valence electrons lose them, resulting in +1, +2, and +3 charges, respectively.
- Nonmetals like nitrogen, oxygen, and fluorine gain electrons to achieve 8 valence electrons, resulting in -3, -2, and -1 charges, respectively.
All compounds need to be neutral, so the charges of the ions must balance.
- Example: Magnesium fluoride. Magnesium ( $Mg$ ) is in column 2, so it forms $Mg^{2+}$ ions. Fluorine ( $F$ ) is in the next to last column, so it forms $F^-$ ions. Therefore, the formula is $MgF_2$ .
The magnitude of one ion's charge becomes the subscript of the other element.
- Example: Aluminum oxide. Aluminum ions have a +3 charge ( $Al^{3+}$ ), and oxide ions have a -2 charge ( $O^{2-}$ ). So, the formula is $Al2O3$ .
Reduce the ratio to the smallest whole numbers for ionic compounds.
- Example: Calcium sulfide. Calcium ions have a +2 charge ( $Ca^{2+}$ ), and sulfide ions have a -2 charge ( $S^{2-}$ ). The initial formula would be $Ca2S2$ , but it reduces to $CaS$ .

Transition metals (and some others) can form more than one stable ion (multivalent ions).
Roman numerals indicate the charge of the metal ion.
- Example: Iron can form $Fe^{2+}$ and $Fe^{3+}$ ions. $FeCl2$ is iron(II) chloride, and $FeCl3$ is iron(III) chloride.
To name these compounds, calculate the total negative charge of the anions and then determine the charge of the metal cation.
*Example: Copper (I) phosphide
Copper (I) has a positive one charge, so that's $CU^+$ .
Phosphorus has five valence electrons, so it will gain three electrons to make phosphide which is $P^{3-}$ .
The formula is $Cu_3P$
- Example: $CrBr_3$ . Bromide has a -1 charge, and there are three bromides, so the total negative charge of the anions is -3. Therefore, chromium must have a +3 charge. The name is chromium(III) bromide.