Naming inorganic Compounds and Writing Formulas Part 1 Naming inorganic Compounds and Writing Formulas

Common and systematic names IUPAC (international Union of Pure and Applied Chemistry) Each compound has a name indicating its composition and its properties A lot of chemicals have common names as well as the proper IUPAC name Chemicals that should always be named by common name and never named by the IUPAC method are: H2O water, not dihydrogen monoxide NH3 ammonia, not nitrogen trihydrid

To keep in mind! 3

Rules to assign Oxidation Numbers (O.N.) Classification of molecules Rules to assign names and write formula 4. Chemical reactions  balancing equations

1. Oxidation number Oxidation number of an element represents the positive or negative character of an atom of that element in a particular bonding situation

Rules for Assigning Oxidation Numbers (O.N.) The oxidation number of an atom in the elemental state is zero Example: Cl2 and Al both are 0 The oxidation number of a monatomic ion is equal to its charge Example: In the compound NaCl, the sodium has an oxidation number of 1+ and the chlorine is 1- The algebraic sum of the oxidation numbers in the formula of a compound is zero Example: the oxidation numbers in the NaCl add up to 0 The oxidation number of hydrogen in a compound is 1+, except when hydrogen forms compounds called hydrides where it is 1- Examples: H is 1+ in H2O, but 1- in NaH The oxidation number of oxygen in a compound is 2-, except in peroxides when it is 1- Example: In H2O the oxygen is 2-, in H2O2 it is 1- The algebraic sum of the oxidation numbers in the formula for a polyatomic ion is equal to the charge of that ion Example: in the sulfate ion, SO42-, the oxidation numbers of the sulfur and the oxygens add up to 2-. The oxygens are 2- each, and the sulfur is 6+ 6 6

Assigning Oxidation Numbers (O.N.) The oxidation number is equal to the charge Ionic compounds monatomic ions Barium cloride BaCl2 (Ba2+ 2Cl-) +2(Ba) -1(Cl) Sodium solfate Na2SO4 (2Na+ SO42-) +1(Na) -2(SO42-) +6(S) -8(O) = -2 polyatomic ion The algebraic sum of the O.N. in the formula is equal to the charge of the ion  O.N. of the sulfur and the oxygens add up to 2- The O.N in the formula add up to 0 Covalent compounds Hydrogen oxide H2O2 -1 (O) +1 (H) Hydrogen H2 0 (H) The O.N. of an atom in the elemental state is zero 7

Potassium dichromate K2Cr2O7 (2K+ Cr2O72-) Determining the oxidation state The algebraic sum of the oxidation numbers in the formula of a compound is zero Potassium dichromate K2Cr2O7 (2K+ Cr2O72-) N.O. Cr = ? (Cr = +2, +3, +6) 0 = +1(K) + x(Cr) -2(O) 0 = +1(2) + x(2) -2(7) 0 = +2 + x(2) -14 x = -14 +2/2  Cr = + 6 8

Classifying the inorganic compounds BINARY COMPOUNDS composed of just two elements TERNARY COMPOUNDS containing three or more elements

Classifying the inorganic compounds IONIC OXIDE Me + nMe IONIC compounds SALTS COVALENT OXIDE BINARY COMPOUNDS nMe + nMe COVALENT compounds SALTS H+ nMe HYDRACID

Classifying the inorganic compounds Cation + Anion IONIC OXIDE Me + nMe IONIC compounds SALTS COVALENT OXIDE BINARY COMPOUNDS nMe + nMe COVALENT compounds SALTS H+ nMe HYDRACID

Cation name (Me) + Anion name (n-Me) ending in –IDE Naming binary compounds IONIC COMPOUNDS (Me + n-Me) Cation name (Me) + Anion name (n-Me) ending in –IDE Ionic OXIDE K2O (2K+ O2-) Potassium oxide Me + nMe  IONIC COMPOUNDS Binary SALTS BaCl2 (Ba2+ 2Cl-) Barium cloride

Monatomic Ions

When the METALS can assume more than one oxidation state…… Naming binary compounds IONIC COMPOUNDS When the METALS can assume more than one oxidation state……

Cations formed from Transition Metals

Fe (+2, +3) (older name) (systematic-stock name) +2  FeO ferrous oxide iron (II) oxide +3 Fe2O3 ferric oxide iron (III) oxide suffix - OUS (lower) suffix - IC (higher) Roman numerals

Classifying the inorganic compounds Me + nMe nMe + nMe IONIC compounds HYDRACID IONIC OXIDE SALTS COVALENT compounds COVALENT OXIDE 1H+ nMe BINARY COMPOUNDS

Naming binary compounds HYDRACID (1H + nMe) prefix Hydro + anion name (nMe) ending in –IC + acid -IDE  -IC Anion Acid 1H+nMe  HYDRACID HF (H+ F- ) fluoride hydrofluoric acid hydrogen fluoride

Classifying the inorganic compounds IONIC OXIDE Me + nMe IONIC compounds SALTS COVALENT OXIDE BINARY COMPOUNDS nMe + nMe COVALENT compounds SALTS H+ nMe HYDRACID

COVALENT COMPOUNDS (nMe + n-Me) Naming binary compounds COVALENT COMPOUNDS (nMe + n-Me) 1° n-Me name + 2° n-Me name ending in –IDE PREFIX in case of two or more atoms Prefix (mono)- 1 di- 2 tri- 3 tetra- 4 penta- 5 hesa- 6 (use the prefix!) COVALENT COMPOUNDS Covalent OXIDE CO2 carbon dioxide N2O3 dinitrogen trioxide SALTS CCl4 carbon tetracloride PCl5 phosphorus pentachloride nMe + nMe 

Si B P H C S I Br N Cl O F The elemet that occurs first in this series is written and name first, the neme of the second element retains the -IDE ending

I prefissi valgono pure per gli ossidi ionici?!!! Naming oxides I prefissi valgono pure per gli ossidi ionici?!!! stock CrO3 chromium (VI) oxide chromium trioxide MnO2 manganese (IV) oxide manganese dioxide SO3 sulfur (VI) oxide sulfur trioxide SO2 sulfur (IV) oxide sulfur dioxide tradizionale

Classifying the inorganic compounds OXYACID HYDROXIDE TERNARY SALTS H + polyatomic anion Me + OH- Me + polyatomic anion TERNARY COMPOUNDS

TERNARY SALTS AND OXYACIDS CONTAIN POLYATOMIC IONS 24

TERNARY SALTS Cation name (Me) + polyatomic anion ending in Naming ternary compounds TERNARY SALTS Cation name (Me) + polyatomic anion ending in (higher) (lower) ox. State –ATE or –ITE Depending on the charge of the n-Me Na3PO4 (3Na+ PO43-) sodium phosphate MgSO4 (Mg2+ SO42-) magnesium sulfate S = -6 MgSO3 (Mg2+ SO32-) magnesium sulfite S = -4 25

Depending on the charge of the n-Me Naming ternary compounds OXYACID Name of the polyatomic anion changing its ending as follows: Anion Acid -ATE  -IC -ITE  -OUS Depending on the charge of the n-Me followed by –”acid” H + polyatomic anion  OXYACID H2SO4 (2H+ SO42-)  sulfate sulfuric acid HClO2 (H+ ClO2-)  clorite chlorous acid 26

Naming ternary compounds In case of multiple oxidation states  Prefix and suffix Anion name Acid name for the lowest n.ox HYPO ………….–ITE HYPO….………OUS 2° n.ox ....……………….-ITE ..….…………….OUS 3° n.ox ...………………-ATE ….…………………IC for the higher n.ox PER ….………..-ATE PER….……………IC Es. Cl n.ox. 1, 3, 5, 7 n.ox. +1 hypochlorite ClO-  HCIO hypochlorous acid n.ox. +3  chlorite ClO2-  HClO2 chlorous acid n.ox. +5  chlorate ClO3-  HClO3 chloric acid n.ox. +7  Perchlorate ClO4-  HClO4 perchloric acid 27

Classifying the inorganic compounds OXYACID HYDROXIDE TERNARY SALTS H + polyatomic anion Me + OH- Me + polyatomic anion TERNARY COMPOUNDS 28

HYDROXIDE Cation name (Me) + “hydroxide” Naming ternary compounds Me + OH-  HYDROXIDE NaOH (Na+ OH-) sodium hydroxide

Features and Balancing Part 2 Chemical reaction: Features and Balancing

One substance changes chemically into another substance Chemical equation One substance changes chemically into another substance Reactants starting materials Products resulting substance Additional information in a chemical reaction: Physical state of the substance (solid, liquid, or gas) Identifies the solvent, if there is one (a solvent is the solution the materials are dissolved in, such as water) Experimental conditions such as heat, light, or electrical energy added

Types of Chemical Reactions There are four main types of chemical reactions Precipitation reactions Reactions with Oxygen Acid-base reactions Oxidation-reduction reactions

Precipitation reactions A chemical change that produces an insoluble product that will form a solid. Usually the solid can be seen “falling out” of the solution, hence, called precipitation. At other times the solid makes the solution turn from clear to cloudy. 2. Reactions with oxygen Many substances react with oxygen. If the substance contains carbon then carbon dioxide is usually produced. If the substance contains hydrogen, then water is usually produced.

4. Oxidation Reduction Reactions 3. Acid-base Reactions This involves an acid combining with a base to form a salt HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(aq) 4. Oxidation Reduction Reactions Involves the transfer of negative charge from one reactant to another Zn(s) + Cu2+ (aq) → Zn2+ (aq) + Cu (s) An example would be

Features of a chemical reaction CaCO3(s) →∆ CaO(s) + CO2(g) The products are on the right of the arrow. Reactants are on the left of the arrow. The arrow indicates the reaction occurs in one direction.

Features of a chemical reaction CaCO3(s) →∆ CaO(s) + CO2(g) “l” would indicate the substance was a liquid. “s” indicates the chemical is a solid substance “g” indicates the substance is a gas

Features of a chemical reaction Law of Conservation of Mass CaCO3(s) →∆ CaO(s) + CO2(g) The main feature of a chemical equation is that it is balanced with the same number of elements in compounds on both sides of the arrow

Balancing a chemical equation Apply the Law of Conservation of Mass to get the same number of atoms of every element on each side of the equation CaCO3(s) → CaO(s) + CO2(g) On the reactant side On the product side 1 mol Ca 1 mol C 3 moles of O The equation is balanced!

Stoichiometric Coefficients Law of Conservation of Mass  In a balanced reaction, both sides of the equation have the same number of elements. The stoichiometric coefficient is the number written in front of atoms, ion and molecules in a chemical reaction to balance the number of each element on both the reactant and product sides of the equation. This stoichiometric coefficients establish the mole ratio between reactants and products. 

Steps to balancing a chemical equation Step 1 Count the number of moles of atoms of each element on both product and reactant side H2(g) + O2(g)  H2O(l) On the reactant side: 2 moles of H 2 moles of O On the product side: 2 moles of H 1 mole of O

Steps to balancing a chemical equation Step 2 Determine which elements are not balanced H2(g) + O2(g)  H2O(l) The oxygen atoms are not balanced in this equation.

Steps to balancing a chemical equation Step 3 Balance one element at a time H2(g) + O2(g)  H2O(l) First O: H2(g) + O2(g)  2H2O(l) Then H: 2H2(g) + O2(g)  2H2O(l)

Steps to balancing a chemical equation Step 4 After you believe you have successfully balanced the equation, check to make sure you have the same number of atoms on both sides of the equation. 2H2(g) + O2(g)  2H2O(l) 4 moles of H 2 moles of O 4 moles of H 2 moles of O

The equation is not balanced Fe3O4 + H2 Fe + H2O On the reactant side On the product side 3 moles Fe 1 mole Fe 2 moles H 4 moles O 1 mole O The equation is not balanced

Fe3O4 + H2 Fe + H2O 1) Fe: Fe3O4 + H2 3 Fe + H2O 2) O: Fe3O4 + H2 3 Fe + 4 H2O 3) H: Fe3O4 + 4 H2 3 Fe + 4 H2O

The equation is balanced Fe3O4 + 4 H2 3 Fe + 4 H2O On the reactant side On the product side 3 moles Fe 3 mole Fe 8 moles H 4 moles O 1 mole O The equation is balanced