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• Formulas use chemical symbols and numbers to show what elements and how many atoms of each are involved in each compound
Chemical Symbols
• Each element has been assigned a one-, two- or three- letter symbol for its identification
• First letter is ALWAYS capitalized, additional letters are lowercase
• Only recently discovered, unnamed elements are given three- letter symbols
• Some symbols show a relationship
– Ex. Carbon ~ C
Sodium ~ Na (Latin – natrium)
• Symbols are assigned by IUPAC
– International Union of Pure and Applied Chemists
• Roots used for naming elements:
0 : nil 1 : un 2 : bi 3: tri 4 : quad
5 : pent 6 : hex 7 : sept 8 : oct 9 : enn
Chemical Molecules
• Monatomic molecules – uncombined elements, written without a subscript
– Ex. Neon gas – Ne
Argon gas – Ar
• Diatomic molecules – elements can exist in nature as two identical atoms bonded together
– Ex. Hydrogen – H2
(F, O, N, Cl, Br, I)
Chemical Formulas
• Chemists have identified over 10 million compounds
• Compound – two or more elements that are chemically combined (bonded together) in definite proportions by mass
– Ex. H2O, C6H12O6, H2O2
• Chemical formula – shows the kinds and numbers of atoms in the smallest representative unit of the substance
– If monatomic: use chemical symbol (ex. Kr)
– If diatomic or a compound: use chemical symbols of elements involved, and subscripts to represent # of atoms present (ex. F2 or O3 or NaCl)
– Types of formulas: molecular, empirical, structural
• Subscript – smaller number after an element symbol that indicates how many atoms of that element are in the molecule
– Ex. H2O means there are 2 H and 1 O atom
• Coefficient – number in front of a molecule’s formula indicating how many molecules are present
– Ex. 2H2O means there are 2 water molecules
• Molecular formulas – shows the kinds and numbers of atoms present in a molecule of a compound
– Subscript written after the symbol indicates the # of atoms of each element
• If only 1 atom, subscript of 1 is omitted
– Show composition but NOT molecular structure
• Empirical formula (“formula unit”) – shows the lowest whole number ratio of ions in a compound
– Ex. MgCl2
• For every 1 Mg+, there are 2 Cl-
– Ex. H2O and H4O2
• Both have a ratio of 2 H : 1 O
• Molecular formulas can be seen as a multiple of an empirical formula
– Ex. Glucose: C6H12O6 (molecular)
CH2O (empirical)
6(CH20) = C6H12O6
• Law of definite proportions – in any compound, the masses of the elements involved are always in the same proportions
– Ex. NaCl always has 1 Na (23 amu) and 1 Cl (35 amu) = 58 amu total for one NaCl
– Ex. H2O always has 2 H (total 2 amu) and 1 O (16 amu) = 18 amu total for one H2O
– Proportions of mass equals the ratio proportions of the number of atoms of each element in the molecule
• Law of multiple proportions – whenever two elements form more than one compound (ex. H2O and H2O2), the different masses of one element (ex. O versus O2) that combine with the same mass of the other element (H2) are in the ratio of small whole numbers
– Ex. We have two compounds, each with 2 g of element B. Compound 1 has 5g element A, compound 2 has 10 g element A