Unit 4 Formulas and Equations

1. Unit 4Formulas and Equations Textbook Chapter 2, 6, & 8 Review Book Topic 2

2. Formulas use chemical symbols and numbers to show what elements and how many atoms of each are involved in each compound

3. 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

4. Some symbols show a relationship • Ex. Carbon ~ C Sodium ~ Na (Latin – natrium) • Symbols are assigned by IUPAC • International Union of Pure and Applied Chemists

5. 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

6. Ex. Element #109 • Un-nil-enn-ium (1)-(0)-(9) • Ex. Element #114

7. 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)

8. 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

9. No two compounds have identical properties

10. 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

11. 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

12. 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

13. 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

14. Molecular formulas can be seen as a multiple of an empirical formula • Ex. Glucose: C6H12O6 (molecular) CH2O (empirical) 6(CH20) = C6H12O6

15. Structural formula – shows the physical organization of the atoms in a molecule

16. 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

17. 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

18. 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

19. Atoms, Compounds and Ions • Atoms and compounds are electrically neutral • (# p+ = # e-) • Ions have a net charge, either (+) or (-) • (# p+≠ # e-) • (+) ions attract (-) ions in a ratio that produces a neutral compound