CCl4 MgCl2 Guess at the names for these two compounds Predict whether each is ionic or molecular compound Understanding the difference in bonding, explain what the formula means for each. Draw pictures to help.

Ch. 7: Chemical Formulas and Compounds 7.1 Chemical Names and Formulas

Chemical Formulas molecular compound number of atoms of each element contained in a single molecule of the compound NO2, CH2Cl2,

Chemical Formulas ionic compound represents one unit simplest ratio of the compound’s anions and cations MgO, Mg(OH)2, NH4Br

Monatomic Ions ions formed from a single atom Not all representative elements easily form ions Some atoms form covalent bonds instead Others form ions without noble gas configurations d and p-block metals many form 2+ or 3+, some +1 or +4

Monatomic Ions type of ion can be determined by looking at the number of valence electrons in the neutral atom atom is more stable with a full shell zero valence eight valence determine which is easier for the atom to achieve

Monatomic Ions cation anion positive if the atom looses electrons B Mg Rb Al Na anion negative if the atom gains electrons N F S O Br

Naming Monatomic Ions cations anions positive ion written first in formula by element’s name add a Roman Numeral if it can form more than 1 type of ion that is called the STOCK SYSTEM anions negative ion written second in formula by element’s root name + -ide ending

Practice: Monatomic Ions Identify the ion created by each atom and give the name for it: Ga Se Ca I N Li

Binary Ionic Compounds compounds made of two different ions Crossing Over write two ion types with correct charges make the anion’s charge, the cation’s subscript make the cation’s charge, the anion’s subscript simplify the ratio if possible combine the names of the two ions

Practice Naming Al and O Mg and Br Cu and Br

Practice Writing Formulas iron (III) sulfide cadmium oxide potassium nitride tin (IV) sulfide

Cmpds with Polyatomic Ions the cation or anion could be a group of covalently bonded atoms instead of one atom most polyatomic ions are oxyanions: polyatomic ions containing oxygen more than one polyatomic anion can exist for one element sulfate ion: SO42- sulfite ion: SO32-

Practice Naming KNO3 Cu(OH)2 NH4Br CaSO4

Practice Writing Formulas sodium permanganate iron (II) chlorite ammonium acetate calcium carbonate

Write the formula for barium nitrate zinc phosphate

Traditional vs. Stock Stock System using Roman Numerals in names of metal ions with more than one charge more important and used more often Traditional add –ic or –ous ending to metal name lower charge: -ous, higher charge: -ic Shown on the chart below Stock System chart

Examples stannous chloride SnCl2 chromic sulfate Cr2(SO4)3 plumbic nitrate Pb(NO3)4 aurous oxide Au2O ferric phosphate FePO4 cuprous chlorate CuClO3

Examples CuCl2 cupric chloride PbO plumbous oxide Pt3(PO4)2 plantinous phosphate AuBr aurous bromide PbO plumbous oxide SnO2 stannic oxide Fe(NO3)2 ferric nitrate

Acids Acids will be focused on greatly later on in the year Can be identified by H being first in the formula All acids contain hydrogen and an anion Can be polyatomic anion : HNO3 Can be monatomic anion: HBr Name is based on the anion root

Binary Acids Only H and monatomic anion Naming Example: Hydro- prefix Root of element name -ic ending Example: HBr, hydrobromic acid H2S, hydrosulfic acid

Ternary Acids Naming Examples: Polyatomic ion name If it has –ate ending, change to –ic If it has –ite ending, change to –ous Examples: HNO3: from nitrate, so nitric acid H2SO3: from sulfite, so sulfurous acid

Practice HBr HClO3 H2C2O4 H3AsO3 H2Se H2S Binary Ternary ternary

Writing Formulas for Acids Determine if it is binary (has hydro-prefix) or ternary (no hydro- prefix) Find the anion contained in the acid (binary- element, ternary- polyatomic ion) add enough H+1 ions to balance the charge of the anion to the front

Practice sulfuric acid nitrous acid hydrofluoric acid hydrosulfuric acid phosphoric acid formic acid

Write the formula for sodium sulfate and determine the number of atoms in one unit.

Ch. 7: Chemical Formulas and Names Oxidation States 7.2

Oxidation States Also called oxidation numbers For an individual atom of one type Used to indicate approximate electron distribution in covalent bonding What contains covalent bonding? molecular compounds polyatomic ions Not “real” charges since electrons are shared

Oxidation States Shared electron are counted as “belonging” to the more electronegative element Since it has a greater attraction for the electrons The electrons are actually located closer to that atom Periodic Trend for electronegativity: increases as you go up and go to the right

Assigning Oxidation States Atoms in a pure element have an oxidation state of zero. Na, O2, Fe, etc. The more electronegative element is assigned a negative oxidation state while the less electronegative has a positive state CF4: C is less electronegative: + F is more electronegative: -

Assigning Oxidation States Fluorine always has a -1 oxidation state Since it is the most electronegative element Oxygen usually has a -2 Except in peroxides where it has a -1 Hydrogen usually has a +1 But when it is paired with a metal, it has a -1: ex. LiH

Assigning Oxidation States Sum of the oxidation states must equal zero in a neutral compound N2O5 O: -2 x 5 = -10 N: +5 x 2 = +10 Sum of states in a monatomic/polyatomic ion must equal the charge of the ion NO3-1 O: -2 x 3 = -6 N: + 5 x 1 = +5

Practice UF6 H2SO4 ClO31- CO2

Determine the Oxidation States Br2 NH3 CaSO3 HSO3- B2H6

Using Oxidation States in Naming Can use Stock System (Roman Numerals) to names molecular compounds too The Roman Numeral in between names is the oxidation state of the first element Do NOT simplify the ratio for molecular compounds Only use either prefixes OR Roman Numerals, not a combination

Stock System Naming PCl3 PCl5 N2O Cl: -1, P: +3 Phosphorus (III) chloride PCl5 Cl: -1, P: +5 Phosphorus (V) chloride N2O O: -2, N: +1 Nitrogen (I) oxide

Stock System Naming CO BrF5 CO2

Writing Formulas Sulfur (IV) fluoride S: +4 F: -1 SF4 Nitrogen (II) oxide N: 2+ O: -2 NO Lead (IV) oxide Pb: +4 Pb2O4 PbO2 Sulfur (IV) fluoride S: +4 F: -1 SF4 Chlorine (III) fluoride Cl: +3 ClF3

Writing Formulas Hydrogen (I) oxide carbon (IV) sulfide Phosphorus (V) oxide Boron (III) hydride Hydrogen (I) oxide carbon (IV) sulfide

Name the following compounds Na2CO3 H2SO4 N2O4

Ch. 7: Chemical Formulas and Compounds 7.3 Uses of Chemical Formulas

Meaning of Formulas the subscripts in a formula give you: simplest ratio of atoms OR the number of atoms in a molecule What type of compound has a formula that is always the simplest ratio? What type of compound forms a molecule? ALWAYS provides you with a ratio of moles of atoms Example: C13H18O2 13 moles of C : 18 moles of H : 2 moles of O

Example If I have 2.00 moles of C13H18O2, how many moles of each atom would I have?

Formula Mass mass of a molecule, ion, or formula unit sum of mass of all atoms in the chemical formula in amu Ex: H2O 18.01528 amu formula mass = molecular mass for molecular compound

Example Find formula mass of potassium chlorate. KClO3

Molar Masses mass of one mole of pure substance numerically equal to formula mass units: g/mol Find molar mass of barium nitrate. Ba(NO3)2

Molar Mass in Conversions can be used as a conversion factor between grams and moles What is the mass in grams of 2.50 mol of oxygen gas?

Example Ibuprofen, C13H18O2, is the active ingredient in many pain relievers. Find molar mass:

Example If the tablets in a bottle contain a total of 33 g of ibuprofen, how many moles are in one bottle? How many molecules of ibuprofen are in the bottle?

Example What is the number of moles of carbon in that bottle? What is the total mass in grams of carbon in the bottle?

I have a sample of 67.9 grams of Na2CO3 Find the number of moles Find the number of molecules

Percentage Composition percentage by mass of each element in the compound divide each part of your molar mass into totals for each element divide these element subtotals by the total molar mass of compound multiply by 100

Example 1 Find the percent composition for calcium nitrite Ca(NO3)2

Example 1 Find the total molar mass Divide each mass by total molar mass and multiply by 100

Example 2 Find percent composition for C13H18O2

Hydrates ionic compound with water molecules attached to it name of ionic compound + prefix for the number of water molecules + hydrate Ex. CaSO4.2H2O calcium sulfate dihydrate

Practice Naming BaCl2.2H2O LiClO4.3H2O MgCO3.5H2O barium chloride dihydrate LiClO4.3H2O lithium perchlorate trihydrate MgCO3.5H2O magnesium carbonate pentahydrate

Finding Percent of Water find percent composition but group all of the atoms in water molecules together To find the % of water find total molar mass (including waters) divide mass of just waters by total multiply by 100

Example 172.172 g/mol Find percent of water in CaSO4.2H2O total molar mass = 172.172 g/mol

Example Divide mass of water by total and multiply by 100

Consider N2O4 and NO2 name both of these using prefixes find the percent composition of N and O for both compounds

Ch. 7: Chemical Formulas and Compounds 7.4 Determining Chemical Formulas

Empirical Formula formula containing the simplest whole-number ratio of atoms not necessarily the CORRECT molecular formula Why? Ex. BH3 and B2H6 have same empirical formula have different molecular formulas

Calculating Empirical Formulas from Percent Composition convert percentage to grams by assuming there is 100 g total of sample convert grams to moles for each element using molar mass identify the smallest mole value divide each mole value by that smallest value

Example 1 Quantitative analysis shows that a compound contains 32.38% Na, 22.65% S, and 44.99% O. Find the empirical formula. Assuming you have 100 g of sample total 32.38 g Na 22.65 g S 44.99 g O

Example 1 Na2SO4 sodium sulfate Convert each of those to moles Divide by the smallest / 0.7064 ≈ 2 / 0.7064 ≈ 1 / 0.7064 ≈ 4 Na2SO4 sodium sulfate

Calculating Empirical Formula from mass composition don’t have to assume you have 100 g since you have an actual amount follow all other steps the same way

Example 2 Analysis of a 10.150 g sample of a compound containing only P and O, is known to contain 4.433 g of P. Find the empirical formula. Find the mass of all components mtotal = mP + mO so mO = mtotal – mP mass of O : 10.150 - 4.433 = 5.667g

Example 2 P2O5 diphosphorus pentoxide Convert all mass values to moles using molar mass Divide by the smallest mole value To get rid of the decimal, multiply by integer / 0.1431 ≈ 1 / 0.1431 ≈ 2.5 P2O5 diphosphorus pentoxide

Molecular Formulas the molecular formula is the actual formula for a compound could be same as the empirical formula but doesn’t have to be same to find molecular formula: need empirical formula need actual molar mass (or formula mass) compare the molar mass of empirical formula to actual molar mass

Example 1 In the last example the empirical formula was found to be P2O5. Experimentation shows that the molar mass is actually 283.89 g/mol. Find the molecular formula.

Example 1 Find molar mass of empirical formula 2(30.9738) + 5(15.9995) = 141.9446 Divide the actual molar mass by the empirical formula’s molar mass 283.89 / 141.9446 ≈ 2 Multiply this number by each subscript in empirical formula P4O10

Example 2 Use the percent composition to write the empirical formula: 41.39 % C, 3.47 % H, and 55.14 % O / 3.443 = 1 / 3.443 = 1 / 3.443 = 1

Example 2 empirical formula: CHO If the actual molecular mass is 116.07 g/mol, what is the molecular formula? Find the molar mass of CHO 1(12.011) + 1(1.00794) + 1(15.9994) = 29.018 Compare it to the molecular mass 116.07 / 29.018 = 4 Multiply that by subscripts 4(CHO) = C4H4O4

Example 3 The empirical formula of a compound is C2H5 and its formulas mass is 87 amu. What is the molecular formula?