Naming Compounds Writing Formulas and Equations Naming Compounds

Naming Compounds Writing Formulas and Equations

Naming Compounds The chemical formula represents the composition of each molecule. In writing the chemical formula, in almost all cases the element farthest to the left of the periodic table is written first. So for example the chemical formula of a compound that contains one sulfur atom and six fluorine atoms is SF 6. If the two elements are in the same period, the symbol of the element of that is lower in the group (i. e. heavier) is written first e. g. IF 3.

Naming Ionic Compounds • Ionic compounds are combinations of positive and negative ions. • In writing the chemical formula the positive ion is written first, It is then followed by the name of the negative ion. • Monatomic anions end in ide. Special endings apply for polyatomic ions • Examples Na. Cl Sodium chloride Ba. F 2 Barium Fluoride Zn. O Zinc Oxide

Names of Polyatomic Ions with Oxygen Polyatomic ions usually contain oxygen in addition to another element. Normally they have a negative charge. They end in either "ate" or "ite" depending on the number of oxygen atoms present. Cl. O 2 Cl. O 3 - hypochlorite chlorate Cl. O 4 - perchlorate NO 2 - Nitrite NO 3 PO 33 - Nitrate phosphite PO 43 - phosphate SO 32 SO 42 - sulfite sulfate

Polyatomic Ion -- Exceptions Most polyatomic ions contain oxygen Their names end in “ite” or “ate”. There are several exceptions OH- hydroxide CN- cyanide SCN- thiocyanate

Elements with Multiple Cations 1. 2. 3. 4. When an element can form more than one cation a Roman numeral is used to distinguish the oxidation state of the compound. Iron, Tin, Lead, Copper, and are common elements with more than one cation. Examples Pb. SO 4 = lead (II) sulfate This compound is formed from Pb 2+ and SO 42 Pb(SO 4)2 = lead (IV) sulfate This compound is formed from Pb 4+ and SO 42 Fe(OH)2 = iron (II) hydroxide This compound is formed from Fe 2+ and OH- Fe(OH)3 = iron (III) hydroxide This compound is formed from Fe 3+ and OH-

Examples of Ionic Compounds Na. Cl Zn. F 2 KOH Ca(NO 3)2 Ba. SO 3 Al 2(SO 4) 3 Ca 3(PO 3)2 NH 4 Cl chloride 9. (NH 4)2 CO 3 1. 2. 3. 4. 5. 6. 7. 8. = = = = Sodium chloride Zinc fluoride Potassium hydroxide Calcium nitrate Barium Sulfite Aluminum sulfate Calcium phosphite Ammonium = Ammonium carbonate

Naming Covalent Compounds • When naming covalent compounds, the name of the first element in the formula is unchanged. • The suffix “-ide” is added to the second element. Often a prefix to the name of the second element indicates the number of the element in the compound • Examples: • SF 6 – sulfur hexafluoride P 4 O 10 – tetraphosphorous decoxide CO – carbon monoxide CO 2 – carbon dioxide

Covalent molecules with multiple possibilities A Roman Numeral is used to indicate the state of the more positive element Examples 1. N 2 O = Nitrogen (I) oxide Since oxygen has a 2 - charge, the nitrogen must be 1+ to balance the charges. Also known as dinitrogen monoxide 2. N 2 O 3 = Nitrogen (III) oxide Since oxygen has a 2 - charge, the nitrogen must be 3+ to balance the charges Also known as dinitrogen trioxide

Binary compounds of Hydrogen The binary compounds of hydrogen are special cases. They were discovered before a convention was adopted and hence their original names have stayed. Water H 2 O is not called dihydrogen monoxide Hydrogen forms binary compounds with almost all nonmetals except the noble gases. Examples HF - hydrogen fluoride HCl - hydrogen chloride H 2 S - hydrogen sulfide

Acids When many hydrogen compounds are dissolve in water they take on the form of an acid. Special rules apply to acids. The “ite” suffix becomes “ous” and the “ate” suffix becomes “ic” HCl Hydrochloric Acid Cl- Chloride HNO 2 Nitrous Acid NO 2 - Nitrite HNO 3 Nitric Acid NO 3 - Nitrate H 2 SO 3 Sulfurous Acid SO 32 - Sulfite H 2 SO 4 Sulfuric Acid SO 42 - Sulfate H 3 PO 3 Phosphorous Acid PO 33 - Phosphite H 3 PO 4 Phosphoric Acid PO 43 - Phosphate H 2 CO 3 Carbonic Acid CO 32 - Carbonate

Writing Formulas for Ionic Compounds Write the positive ion (cation) first, then the negative ion. The positive charges must balance the negative charges. Use subscripts to show many times each ion must appear in order for the charges to balance. A subscript is not used if the ion appears only once Use parenthesis around polyatomic ions that appear more than once in the formula

Examples 1. 2. 3. 4. 5. 6. 7. 8. 9. Na+ and Cl- = Na. Cl Zn 2+ and Br- = Zn. Br 2 K+ and OH- = KOH Ca 2+ and OH- = Ca(OH)2 Fe 2+ and SO 42 - = Fe. SO 4 Fe 3+ and SO 42 - = Fe 2(SO 4) 3 Ca 2 + and PO 43 - = Ca 3(PO 4)2 NH 4+ and Cl- = NH 4 Cl NH 4+ and CO 32 - = (NH 4)2 CO 3

Chemical Reactions Elements and compounds frequently undergo chemical reactions to form new substances In a chemical reaction, chemical bonds are frequently broken and new chemical bonds are formed Atoms are neither created nor destroyed in an ordinary chemical change

Chemical Reactions A balanced chemical reaction is used to describe the process that occurs in a chemical change. For example: Zinc reacts with hydrochloric acid to produce zinc chloride and hydrogen gas. This chemical reaction could be written as Zn + 2 HCl Zn. Cl 2 + H 2

Reactants and Products In the chemical reaction Zn + 2 HCl Zn. Cl 2 + H 2 Reactants Products This shorthand way of describing a chemical reaction is known as a chemical equation The starting materials are shown on the left and are known as reactants The substances formed are shown on the right and are known as the products

Balancing a Chemical Reaction A proper chemical reaction must be balanced Zn + 2 HCl Zn. Cl 2 + H 2 Reactants Products Each element must appear on both sides of the arrow and equal number of times Chemical reactions can be balanced by inserting numbers in front of formulas. These numbers are called coefficients

Balancing Chemical Reactions Most simple equations can be balanced by inspection Example: Balance the following equation Ba. Cl 2 + K 3 PO 4 Ba 3 (PO 4)2 + KCl • • • There are 3 Ba on the right so we need coefficient of 3 in front of Ba. Cl 2 There are 2 PO 4 on the right so we need a coefficient of 2 in front of K 3 PO 4. This leaves 6 K on the left so we need a coefficient of 6 in front of the KCl on the right The balanced equation is 3 Ba. Cl 2 + 2 K 3 PO 4 Ba 3 (PO 4)2 + 6 KCl

Balancing Chemical Reactions An equation is balanced when there are the same number and kind of atoms on both sides of the arrow 3 Ba. Cl 2 + 2 K 3 PO 4 Ba 3(PO 4)2 + 6 KCl Reactants (Left) Products (Right) Ba 3 Ba 3 Cl 3 x 2 = 6 Cl 6 K 2 x 3 = 6 K 6 P 2 P 2 O 2 x 4 = 8 O 2 x 4 = 8

State Symbols State symbols are often added to chemical equations. Ca. CO 3 (s) + 2 HCl (aq) Ca. Cl 2 (aq) + CO 2 (g) + H 2 O (l) Symbols (s) (l) Solid Liquid (g) (aq) Gas Aqueous (Water Solution)

Types of Reactions There are many kinds of chemical reactions that occur. Some are very simple while others are very complex and may occur in multiple steps. A number of reactions conform to some relatively simple patterns Understanding and identifying these patterns can be helpful in predicting the products of similar reactions

Direct Combination In a direct combination, two elements or compounds combine to form a more complicated product Examples Ca. O + CO 2 Ca. CO 3 2 H 2 + O 2 2 H 2 O Fe. Cl 2 + Cl 2 Fe. Cl 3 N 2 + O 2 2 NO

Decomposition In a dcecomposition, a single compound is broken down into two or more simplier substances Examples 2 KCl. O 3 2 KCl + 3 O 2 Zn. CO 3 Zn. O + CO 2 Cu(OH)2 Cu. O + H 2 O

Single Replacement In a single replacement, one substance (usually an element) takes the place of another in a compound Examples Zn + H 2 SO 4 Zn. SO 4 + H 2 Cl 2 + 2 KBr 2 KCl + Br 2 Mg + Cu. Cl 2 Mg. Cl 2 + Cu

Double Replacement In a double replacement, two substances exchange places in their respective compounds Examples Ag. NO 3 + Na. Cl Ag. Cl + Na. NO 3 3 Ca. Cl 2 + 2 K 3 PO 4 Ca 3(PO 4)2 + 6 KCl Ba. Cl 2 + Na 2 SO 4 Ba. SO 4 + 2 Na. Cl

Diatomic Molecules Certain elements exist as diatomic molecules in nature H 2 Hydrogen N 2 Nitrogen F 2 Fluorine O 2 Oxygen I 2 Iodine Cl 2 Chlorine Br 2 Bromine

Diatomic Molecules Certain elements exist as diatomic molecules in nature H 2 Hydrogen Have N 2 Nitrogen No F 2 Fluorine Fear O 2 Oxygen Of I 2 Ice Iodine Cl 2 Chlorine Cold Br 2 Bromine Beer