Chemical Symbols and Formulas

1. Chemical Symbols and Formulas

2. Chemical Symbols • Each element is assigned a unique one, or two letter symbol. • Some of the more common elements have a single symbol, such as O for oxygen and H for hydrogen. • The elements are given systematic names that represent their atomic number.

3. Chemical Symbols • Sometimes the letters of the symbol do not correspond to the English element but relate to the Latin or Greek name: • For example: Some periodic tables have K (postassium) labeled as (Latin – Kalium), and Na (sodium) labeled as (Latin – Natrium)

4. Diatomic Molecules • When writing the symbols of uncombined elements, always written as monatomic, that is without a script. • A subscript is a number to the right and slightly below a symbol that tells the number of atoms present • A subscriptis not written if only one atom is present.

5. Diatomic Molecules

6. Chemical Formulas • Compounds are composed of combinations of elements chemically combined in definite proportions by weight (mass). • Formulas use chemical symbols and numbers to show both qualitative and quantitative information about the substance.

7. Chemical Formulas • Qualitative information relates to things that cannot be counted or measured; such as what elements are in the compound. • Quantitative information deals with things that can either be counted or measured; such as the number of atoms of each is present in a unit of a compound.

8. Chemical Formulas • In the formula of a compound, the symbols for the elements supply the qualitative information. • The formula CO tells the reader that the compound consist of carbon and oxygen.

9. Chemical Formulas • Notice the difference between CO and Co. • The first combination of two elements in a compound, carbon dioxide, while the second is the symbol for cobalt.

10. Chemical Formulas

11. Types of Formulas • Two basic types of formulas provide different types of information about a compound. • Empirical formulas include all types of compounds. • Molecular formulas are important when considering compounds formed atoms sharing electrons.

12. Empirical Formulas • Empirical formulas represents the simplest integer ratio in which atoms combine to form a compound. • Ionic substances do not form discrete of molecules, but rather an array of ions (charged particles).

13. Empirical Formulas • Ionic formulas indicate the ration of ions in a compound in the simplest ratio. • The formula MgCl tells us that for every magnesium ion in the compound there are two chloride ions. • Formulas for ionic substances are empirical formulas. 2

14. Molecular Formulas • Covalent Bonds: chemical bonds that involve the sharing of a electron pair. • Covalently bonded substances form discrete units called molecules.

15. Molecular Formulas • In some cases, such as H O, the empirical formula not only represents the simplest ratio, but it also represents that actual ratio of the atoms in a molecule of water. • In other cases, the molecular formula may be a multiple of the empirical formula. 2

16. Molecular Formulas • For example: • The molecular formula of glucose is C H O which is six times the empirical formula CH O 6 12 6 6

17. Molecular vs. Empirical

18. Atoms, Compounds, and Ions • How do you know what elements form a compound and in what proportion? • To understand how elements form compounds, an understanding of atoms and ions is essential.

19. Atoms, Compounds, and Ions • Atoms and compounds are electrically neutral; that is, they do not have a net charge. • Both atoms and compounds contain positively charged protons and negatively charged electrons, but there are equal numbers of positive and negative charges, producing a neutral atom or compound.

20. Atoms, Compounds, and Ions • Ions, however, are not neutral and may be either positively or negatively charged. • An ion that contain more protons than electrons will be positively charged, while and ion with more electrons will have a negative charge.