Warm Up Define oxidation state A star with a volume of 3.6 x 106 km3 drifts into the neighborhood of Sol. It has a mass of 6.0 x 1018 Kg. What is it’s density?

Naming Compounds Writing Formulas and Equations Larry Scheffler Lincoln High School

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 SF6. If the two elements are in the same group, the symbol of the element of that is lower in the group (i.e. heavier) is written first e.g. IF3.

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 NaCl Sodium chloride BaF2 Barium Fluoride ZnO Zinc Oxide

5.7 Nomenclature Naming of Compounds Binary Compounds have two types of atoms (not diatomic which has only two atoms). Metals (Groups I, II, and III) and Non-Metals Metal _________ + Non-Metal _________ide Sodium Chlorine Sodium Chloride NaCl Metals (Transition Metals) and Non-Metals Metal ______ +Roman Numeral (__) + Non-Metal ________ide Iron III Bromine Iron (III) Bromide FeBr3 Compare with Iron (II) Bromide FeBr2

Let’s Practice! Name the following. CaF2 Calcium Flouride K2S Potassium Sulfide CoI2 Cobalt (II) Iodide or Cobaltous Iodide SnF2 Tin (II) Flouride or Stannous Flouride SnF4 Tin (IV) Flouride or Stannic Flouride OF2 Oxygen diflouride CuI2 Copper (II) Iodide or Cupric Iodide CuI Copper (I) Iodide or Cuprous Iodide SO2 Sulfur dioxide SrS Strontium Sulfide Lithium Bromide LiBr

Names of Polyatomic Ions with Oxygen ClO- hypochlorite ClO2- chlorite ClO3- chlorate ClO4- perchlorate NO2- Nitrite  NO3- Nitrate  PO33- phosphite  PO43- phosphate  SO32- SO42- sulfite  sulfate  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.

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 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 PbSO4  =  lead (II) sulfate   This compound is formed from Pb2+ and  SO42- Pb(SO4)2 =  lead (IV) sulfate   This compound is formed from Pb4+ and  SO42- Fe(OH)2  =  iron (II) hydroxide   This compound is formed from Fe2+ and  OH-  Fe(OH)3  =  iron (III)  hydroxide    This compound is formed from Fe3+ and  OH-

Examples of Ionic Compounds NaCl = Sodium chloride ZnF2 = Zinc fluoride KOH = Potassium hydroxide Ca(NO3)2 = Calcium nitrate BaSO3 = Barium Sulfite Al2(SO4) 3 = Aluminum sulfate Ca3(PO3)2 = Calcium phosphite NH4Cl = Ammonium chloride (NH4)2CO3 = 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: SF6 – sulfur hexafluoride P4O10 – tetraphosphorous decoxide CO – carbon monoxide CO2 – carbon dioxide

Nomenclature Naming of Compounds Binary Compounds have two types of atoms (not diatomic which has only two atoms). Metals (Transition Metals) and Non-Metals Older System Metal (Latin) _______ + ous or ic + Non-Metal ________ide Ferrous Bromine Ferrous Bromide FeBr2 Compare with Ferric Bromide FeBr3 Non-Metals and Non-Metals Use Prefixes such as mono, di, tri, tetra, penta, hexa, hepta, etc. CO2 Carbon dioxide CO Carbon monoxide PCl3 Phosphorus trichloride CCl4 Carbon tetrachloride N2O5 Dinitrogen pentoxide CS2 Carbon disulfide

Covalent molecules with multiple possibilities A Roman Numeral is used to indicate the state of the more positive element Examples N2O   =  Nitrogen (I) oxide   Since oxygen has a 2- charge, the nitrogen must be 1+ to  balance the charges.    Also known as dinitrogen monoxide N2O3 =  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 H2O is not called dihydrogen monoxide Hydrogen forms binary compounds with almost all non-metals except the noble gases. Examples HF - hydrogen fluoride HCl - hydrogen chloride H2S - 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 HNO2 Nitrous Acid NO2- Nitrite HNO3 Nitric Acid NO3- Nitrate H2SO3 Sulfurous Acid SO32- Sulfite H2SO4 Sulfuric Acid SO42- Sulfate H3PO3 Phosphorous Acid PO33- Phosphite H3PO4 Phosphoric Acid PO43- Phosphate H2CO3 Carbonic Acid CO32- 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 how 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 Na+ and Cl- = NaCl Zn2+ and Br- = ZnBr2 K+ and OH- = KOH Ca2+ and OH- = Ca(OH)2 Fe2+ and SO42- = FeSO4 Fe3+ and SO42- = Fe2(SO4) 3 Ca2 + and PO43- = Ca3(PO4)2 NH4+ and Cl- = NH4Cl NH4+ and CO32- = (NH4)2CO3

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  ZnCl2 + H2

Reactants and Products In the chemical reaction Zn + 2 HCl  ZnCl2 + H2 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  ZnCl2 + H2 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 BaCl2 + K3PO4  Ba3 (PO4)2 + KCl There are 3 Ba on the right so we need coefficient of 3 in front of BaCl2 There are 2 PO4 on the right so we need a coefficient of 2 in front of K3PO4. 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 BaCl2 + 2 K3PO4  Ba3 (PO4)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 BaCl2 + 2 K3PO4  Ba3(PO4)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 Symbols (s) Solid (l) Liquid (g) Gas (aq) State symbols are often added to chemical equations. CaCO3 (s) + 2 HCl (aq)  CaCl2 (aq) + CO2 (g) + H2O (l) Symbols (s) Solid (l) Liquid (g) Gas (aq) 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 CaO + CO2  CaCO3 2 H2 + O2  2 H2O FeCl2 + Cl2  FeCl3 N2 + O2  2 NO

Decomposition In a dcecomposition, a single compound is broken down into two or more simplier substances Examples 2 KClO3  2 KCl + 3 O2 ZnCO3  ZnO + CO2 Cu(OH)2  CuO + H2O

Single Replacement In a single replacement, one substance (usually an element) takes the place of another in a compound Examples Zn + H2SO4  ZnSO4 + H2 Cl2 + 2 KBr  2 KCl + Br2 Mg + CuCl2  MgCl2 + Cu

Double Replacement In a double replacement, two substances exchange places in their respective compounds Examples AgNO3 + NaCl  AgCl + NaNO3 3 CaCl2 + 2 K3PO4  Ca3(PO4)2 + 6KCl BaCl2 + Na2SO4  BaSO4 + 2NaCl

Diatomic Molecules Hydrogen Fluorine Oxygen Iodine Chlorine Bromine Certain elements exist as diatomic molecules in nature H2 Hydrogen N2 Nitrogen F2 Fluorine O2 Oxygen I2 Iodine Cl2 Chlorine Br2 Bromine

Diatomic Molecules Hydrogen Have Nitrogen No Fluorine Fear Oxygen Of Certain elements exist as diatomic molecules in nature H2 Hydrogen Have N2 Nitrogen No F2 Fluorine Fear O2 Oxygen Of I2 Iodine Ice Cl2 Chlorine Cold Br2 Bromine Beer