Chemical Names and Formulas

I. Ions - have a charge (not neutral) - stable with eight valence electrons. (p+) + (e-) = charge (must be written number first and then the sign) a. Monatomic Ions – i. Cations - metals 1. Electrons – neutral atom lost electrons 2. Charge – positive – has less electrons than protons. 3. Size – smaller than the neutral atom 4. Names – same as atom, just add the word ion 5. Representative (A) Metals: Group 1A=1+ and Group 2A=2+ 6. Transition (B) Metals: no trend, Roman Numerals state charge  Roman Numerals – in parenthesis, after the element name, represents the value (quantity) of the positive charge Review: 1 I 1+ 2 II 2+ 3 III 3+ 4 IV 4+ 5 V 6 VI Examples: Iron (II) Fe 2+ Iron (III) Fe 3+ Copper (I) Cu 1+ Copper (II) Cu 2+ Lead (II) Pb 2+ Lead (IV) Pb 4+  Exceptions, you are expected to know: Silver Group 1B = 1+ Zinc Group 2B = 2+ Cadmium Group 2B = 2+  Classical Names: Suffix is attached to Latin name -ous : represents the lower charge -ic : represents the higher charge Example: Fe 2+ Ferrous Fe 3+ Ferric ii. Anions - nonmetals 1. Electrons – neutral atom gained electrons 2. Charge - negative - has more electrons than protons. 3. Size – larger than the neutral atom 4. Names – change suffix to “ide” ending Examples:

Symbol p+ n0 e- AN AM Charge Ion Name

53 3+ Cr 24

80 1- Br 35

225 Ra 88

b. Polyatomic Ions – groups of atoms acting as a unit with a charge (page 858) Examples: sulfate ion SO4 2- sulfite ion SO3 2- ammonium ion NH4 1+ i. Naming – (-ite) ending means one less oxygen (-ate) ending means 1 more oxygen *does not tell the actual # of oxygens Examples: (-ite) less SO3 NO2 ClO2 (-ate) more SO4 NO3 ClO3 2- 1- 1- Note: Charges do not change, only the number of oxygen ii. Others – ammonium cation NH4 1+ cyanide anion CN 1- hydroxide anion OH 1- iii. When polyatomic formula begins with hydrogen the charges combine by addition. Examples: H (1+) + CO3 (2-) ==> HCO3 (1-) hydrogen + carbonate ==> hydrogen carbonate II. Compounds – two or more atoms that chemically combine – neutral Octet Rule – objective, to have 8 valence electrons to be stable like noble gas. a. Molecular Compound – non-metal plus a non-metal i. called a molecule - (NM+NM) – share electrons in covalent bond ii. low melting and boiling points (below 300ºC) iii. may be solid, liquid, or gas at RT Examples: carbon dioxide CO2 and water H2O b. Ionic Compound – metal plus a non-metal – transfer electrons from metal to nonmetal in ionic bonds i. called a formula unit (M+NM) ii. high MP and BP(usually above 300ºC) iii. usually a crystalline solid at RT Examples: sodium chloride NaCl and calcium chloride CaCl2 c. Chemical Formulas - universal language - shows the types and quantities of atoms in compound. d. Percent Composition (Review) - what a substance is made of – percent by mass of each element in a compound, regardless of sample quantity (size). Formula: % Composition = total mass of one element in the compound x 100 total formula mass of the whole compound e. Law of Definite Proportions (Review) - in a specific chemical compound, the masses of the elements are always in the same proportions. Sample quantity does not matter. (proportional = ratio) Example One: every sample of water is 11.1% hydrogen & 88.9% oxygen 2amu hydrogen x 100 = 11.1% 16amu oxygen x 100 = 88.9% 18amu water 18amu water Note: Does not matter how much of H2O that you have. Hydrogen will always be 11.1% & oxygen will always be 88.9% of the sample. Example Two: 100g of MgS has 56.87g of sulfur & 43.13g of magnesium 43.13g Mg = 0.7584g Mg = a ratio of 0.7584:1 56.87g S 1g S f. Law of Multiple Proportions (more than one ratio) - when two elements form more than one type of compound, different masses of one element combine with the same masses of the other element, in simple whole number ratios. i. One element is equal in both compounds. ii. The other element is in a simple whole number ratio. Example: H2O (water) and H2O2 (hydrogen peroxide) There is the same quantity of hydrogen in each compound. Oxygen atoms are in a ratio of 2:1. Look at the mass... H2O H = 1g x 2 = 2g O = 16g x 1 = 16g H2O2 H = 1g x 2 = 2g O = 16g x 2 = 32g 32g/16g = 2:1 ratio, no units. They cancel out. III. Ionic Compound Formulas – smallest representative unit, subscripts are in simplest whole number ratios a. Binary Ionic Compounds – two monatomic ions combine i. always list the cation first then the anion (+)(-) ii. ion charges cancel so the compound is neutral - net charge is zero iii. least common multiple of subscripts to cancel charges Examples: make potassium chloride from K (1+) and Cl (1-) 1K + 1Cl ==> KCl 1Ca + 2Br ==> CaBr2 Calcium Bromide = Ca(2+) & Br(1-)

2Fe + 3O ==> Fe2O3 iron (III) oxide = Fe(3+) & O(2-)

Shortcut: Criss-Cross” method – exchange charge for subscript 3+ 2- Fe O = Fe2O3 2+ 2- Ca S = Ca2S2 must reduce to CaS Note: Always reduce to the least common multiple. Note: NEVER list charges in a neutral compound. b. Ternary Ionic Compounds - 3 or more elements combine to make an electrically neutral compound. i. Element plus a polyatomic ion or 2 polyatomic ions. ii. Cation always listed first and then the anion iii. Steps – 1) write the symbol and possible charges of each ion 2) balance the charge to a net 0 for the compound Example: Ca (2+) and NO3 (1-) make... Ca(NO3)2 Note: Use parenthesis to show you need more than one polytatomic ion (The whole polyatomic ion is needed more than once, not just part of it!!) Example: Write the chemical formula...Calcium Hydroxide IV. Ionic Compound Names a. Binary Ionic Compounds i. Cation first – no name change - use a Roman Numeral is the metal is a transition metal ii. Anion second – ends in “ide” Examples: name these ionic compounds CaCl2, Na2O, NaCl, BeS, Zn3P2 b. Ternary Ionic Compounds – contain a polyatomic ion Examples: Name these compounds... NaC2H3O2 and K2Cr2O7 V. Molecular Compound Formulas - 2 elements: NM + NM a. two anions – both want electrons – cannot try to balance charges to net 0 b. Diatomic Molecules: cannot exist as single atoms in nature i. seven of them: N2 O2 F2 Cl2 Br2 I2 H2 c. sometimes the two nonmetals can combine in more than one way Examples: CO (poisonous gas) and CO2 (we breath out). They cannot both be named of Carbon Oxide. Need a way to distinguish them. VI. Molecular Compound Names – Use prefixes to represent the subscript (the number of atoms needed for each element to create a neutral compound). 1 mono- CO = carbon monoxide 2 di- 3 tri- CO2 = carbon dioxide 4 tetra- 5 penta- 8 octa- 6 hexa- 9 nona- 7 hepta- 10 deca- You Try to name the following: SO2 and N2O3 Note: Do not put the prefix mono- on the first element. Do not put two of the same vowels together when adding prefixes to elements that begin with vowels. Note: Second anion still ends in (–ide). Do NOT put the (-ide) ending on the first anion. Note: Diatomic Gases – name does not change… I2 Iodine Gas H2 Hydrogen Gas N2 Nitrogen Gas Cl2 Chlorine Gas O2 Oxygen Gas F2 Fluorine Gas Br2 Bromine Gas (Next Unit Exam – skip for now.) VII. The Mole (Chapter 3) a. Molar Mass (Formula Mass) = 1 Mole b. Particles: 6.02x1023 molecules, formula units, atoms, ions = 1 Mole c. Conversions VIII. Calculating Empirical Formulas a. An empirical formula is the lowest whole number ratio of the elements in a compound b. tells the kinds and relative counts of atoms in molecules and formula units. c. will be reduced even if it is a polyatomic. Example: Hydrogen Peroxide H2O2 Empirical Formula = HO Lowest whole number ratio is 1:1 Question: A compound is 25.9% Nitrogen and 74.1% Oxygen. What is it’s empirical formula? Step 1) Assume 100 grams. Convert the mass to moles for each compound. Step 2) If the moles are whole numbers then use them as subscripts on the element to which it refers. Step 3) If the moles have decimals then divide them both by the smaller of the two numbers calculated in step 1. Step 4) If the moles are whole numbers then use them as subscripts on the element to which it refers. Step 5) If the moles still have decimals then multiply by the lowest whole number that will convert the decimal to a whole number. Step 6) Write the empirical formula using the whole number of moles for each element as that elements subscript. IX. Calculating Molecular Formulas when know the empirical formula and the formula mass. Step 1) Given the empirical formula CH and the formula masses 26 grams (Compound A) and 78 grams (Compound B). Step 2) Calculate the empirical formulas mass. Step 3) Divide the formula mass that is given for each compound by the empirical formula mass calculated in step 2. Step 4) Multiply the empirical formulas subscripts by the number calculated in step 3, for each compound. Example: Given the empirical formula CH2O. Given the formula masses 30 grams (methanol), 60 grams (ethanoic acid) and 180 grams (glucose), find the compounds molecular formulas. Now go from percent composition all the way to the molecular formula. (Note: You must go through the empirical formula.) Example: Given the FM of 102 grams. Given the percent composition is 58.8% C, 9.8% H, and 31.4% O. What is the compounds molecular formula.