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Naming Compounds Writing Formulas and Equations

Naming Compounds Writing Formulas and Equations. Ban DHMO?. DHMO is dihydrogen monoxide. Evaluate the following link and decide if this potentially dangerous chemical should be prohibited. www.dhmo.org. Naming Compounds. The chemical formula represents the composition of each molecule.

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Naming Compounds Writing Formulas and Equations

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  1. Naming CompoundsWriting Formulasand Equations

  2. Ban DHMO? DHMO is dihydrogen monoxide. Evaluate the following link and decide if this potentially dangerous chemical should be prohibited. www.dhmo.org

  3. 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.

  4. 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. 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.

  6. 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

  7. 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-

  8. 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

  9. 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

  10. 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 asdinitrogen 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

  11. Binary compounds of Hydrogen Water H2O is not called dihydrogen monoxide The binary compounds of hydrogen are special cases. They were discovered before a convention was adopted and hence their original names have stayed. • Hydrogen forms binary compounds with almost all non-metals except the noble gases. • Examples • HF - hydrogen fluoride • HCl - hydrogen chloride • H2S - hydrogen sulfide

  12. 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”

  13. 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

  14. 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

  15. 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

  16. 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

  17. 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

  18. 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

  19. 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 of6 in front of the KCl on the right The balanced equation is 3 BaCl2 + 2 K3PO4 Ba3 (PO4)2 + 6 KCl

  20. 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

  21. State Symbols • State symbols are often added to chemical equations. CaCO3(s) + 2 HCl (aq) CaCl2(aq) + CO2(g) + H2O (l)

  22. 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

  23. 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

  24. 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

  25. 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

  26. 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 + 6KClBaCl2 + Na2SO4 BaSO4 + 2NaCl

  27. Diatomic Molecules • Certain elements exist as diatomic molecules in nature

  28. Diatomic Molecules • Certain elements exist as diatomic molecules in nature

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