Chapter 7 Chemical Formulas and Bonds 7-1 Ionic Bonding Octet Rule Atoms in compounds tend to have a noble gas configuration. Typically 8 valence electrons. Ionic Bonds Force of attraction between oppositely charged ions. Formula units(f.u.) – chemical formula for an ionic compound.

Cations / Anions Cations Anions Positively charged ions Drops to lower energy level Noble gas or pseudo noble gas configurations Transition and Group 4A elements have multiple charges except Zinc and Silver. (Zn 2+ and Ag 1+ ) Anions Negatively charged ions Fills the valence shell Noble gas configuration Halide, ions formed from the halogen atoms.

Lewis Dot Diagrams Using electron dot notation, sketch the exchange of valence electrons in atoms as they form ionic bonds. Ca + Cl Na + S

Class work Show the electron transfer for the following elements and write the formula unit for the binary ionic compound: 1) Zn + I 2) Al + S 3) K + P 4) Mg + N 5) Ba + Se 6) Fe 2+ + O

Chapter 7.2 Covalent Bonding Force of attraction for electrons, that results in a pair of electrons being shared by two atoms. Molecule (Molecular Substance) Group of atoms that are held together by a covalent bond. Molecular Formula Describes the composition of atoms in a single molecule of a compound. Number and type of each atom in the molecule.

Structural Formula Lewis structures Formula that specifies which atoms are bonded to each other in a molecule. Places atoms in a pattern in which they can share electrons to satisfy the octet rule. Every shared pair of electrons are depicted as a dash ( — ). Shared Pair, are 2 electrons that are shared between 2 atoms. (bonding pairs) All unshared electrons must be paired. Unshared Pair, are 2 electrons that are NOT shared between 2 atoms. (non-bonding pairs)

Single Covalent Bonds Draw the Lewis Structure for the molecule that has 1 N and 3 H.

Systematic over Trial-Error Count the total number of valence electrons (N) needed to account for the atoms (based on the column of the atom in the periodic table) and charge (add one electrons for each negative charge; subtract one electron for each positive charge). Draw the framework with single bonds. Some knowledge of the way the atoms are connected may be required. Using lone pairs, complete octets around the noncentral atoms. Count the number of electrons depicted (two for each bond and two for each lone pair). If this number is less than N, then add electrons to the central atom until the total number of electrons depicted is N. If the octet rule is not satisfied for the central atom and lone-pair electrons are nearby, use those electrons to make double or triple bonds to the central atom.

Practice Single Covalent Bonds Write the Lewis Structure for the following formulas: Cl2 OF2 C2H6 C3H7Cl

Multiple Covalent Bonds Some structures need more than a single pair of electrons to reach the octet. Double covalent bonds Covalent bond between 2 atoms in which there are a total of 4 electrons being shared. (2pair) Triple covalent bonds Covalent bond between 2 atoms in which there are a total of 6 electrons being shared. (3pair) CO2 and HCN

Polyatomic Ion Structures Polyatomic ions, are covalently bonded atoms that form a charge due to the gain or loss of electrons. When drawing the structures of polyatomic ions: (–) ions must have extra electrons in the structure that is equivalent to its charge. (+) ions must lose the number of electrons in the structure that is equivalent to its charge.

Example NH4 1+, the ammonium ion. SO4 2-, sulfate ion.

Classwork Lewis Structure CO SO3 ClO3 1- NO2 1-

Exceptions to the Octet Incomplete Octet Some atoms that have less than eight electrons. Hydrogen only needs two and Boron needs six electrons. Expanded Octet Atoms that can have more than eight electrons. Common expanded octets: Cl, Br, I, S, and P IMPORTANT!!! Only one atom may exceed the standard eight valence. Must be the central atom only!!

Draw the Lewis Structure for the following: BF3 IF4 1-

Coordinate Covalent Bond Like a single covalent bond, but a single atom is sharing 2 electrons with another atom, in which it doesn’t increase the # of electrons in the atom sharing the 2 electrons. Helps to show ownership of electrons. Indicated by using an arrow rather than a dashed line. () The arrow must point away from the atom that has less electrons than what it should.

Practice SO3 NO3 1-

Polar Vs. Nonpolar Covalent Bonds Non-Polar – electrons are equally shared Balanced Distribution of Electrical Charge Polar – electrons are not equally shared Unbalanced Distribution of Electrical Charge Partial Charge  +  -

Determining Bond Type The Difference in electronegativity determines the bond type, page 241, according to the following scale: Example Na-Cl Na = 0.9 Cl = 3.0 Difference = 3.0 – 0.9 = 2.1 Ionic 2.0 4.0 0.4 Ionic Polar Covalent Non

Identifying Bond Types Determine if the following are going to form; Ionic, Polar Covalent or Non-polar Covalent Bonds: C-N Ca-F Br-Br H-Br

7.3 Nomenclature Systematic way of writing and naming compounds. Purpose So we don’t have to memorize all the common names.

Formula Writing for Ionic Compounds Binary Ionic Compounds Ionic Bond between two atoms. Total charge of all compounds must equal zero! Total (+) = Total (-) Identify the charges of both + and – ions. If charges equal no subscripts needed. If charges are not equal, cross the charges without (+/-) signs as subscripts. If the subscripts can be reduced, reduce to the lowest whole number ratio.

Criss Cross Method Binary Ionic Ba + I BaI2

Practice Formulas Write the formula unit for the following Binary Ionic Compounds: (Zn + O) (Fe 3+ + Cl) (Ba + N) (Sn 4+ + S) (Cu 1+ + P)

Writing Ternary Ionic Compounds Polyatomic Ions Combinations of 2 or more non-metals that form common ions. Most are negatively charged except NH4 1+ Must be placed in ( ) if more than 2 are needed. Ternary Ionic Compounds Ionic bond between an element and a polyatomic ion.

Criss Cross Method Ternary Ionic Na and PO4 3- Na3PO4

Practice Writing Ternary Ionic Write the formula unit for the following Ternary Ionic Compounds: (Zn + OH-) (Fe 3+ + ClO3 1- ) (Ba + NO3 1- ) (Sn 2+ + SO3 2- ) (NH4 1+ + P)

Naming Ionic Compounds Binary Ionic Compounds Metals, retain the name of the element. Multiple charged metals must use either stock or classical naming system. (STOCK is preferred) Non-metals, root of name + ide ending. Chlorine = Chloride Nitrogen = Nitride Phosphorus = Phosphide Sulfur = Sulfide Carbon = Carbide Oxygen = Oxide Selenium = Selenide Example MgS Magnesium Sulfide

Naming Binary Ionic Naming ionic compounds with multiple charged metals. Use roman numerals after the metal to indicate the charge of the metal. Cr2O3 Chromium (III) Oxide PbS Lead (II) Sulfide

Practice Naming Binary Ionic NaCl FeCl2 CaF2 KI Al2O3 SnO

Naming Ternary Ionic Compounds Same rules as the binary ionic compounds. Except look up the name of the polyatomic ion, on the ion sheet given, and write the name as it appears. Example NaNO3 Sodium Nitrate

Practice Naming Ternary Ca3(PO4)2 AgNO2 Zn(OH)2 KC2H3O2

Summary of Writing Ionic Compounds Identify the ending of the compound. -ide, check to see if the first word is NOT ammonium or the last is NOT Hydroxide or Cyanide. If not, then it’s a binary ionic compound. If yes, then it’s a ternary ionic compound. -ate, or –Ite Ternary Ionic Compound.

Summary of Naming Ionic Compounds Identify the number of elements. 2 elements, then it’s a binary ionic: Name the metal + nonmetal –ide 3 or more elements, then it’s a Ternary ionic: Name the metal + polyion OR Name the polyion + nonmetal –ide OR Name the polyion + polyion

Practice Write formulas for the following compounds: Silver Sulfide Magnesium Nitrate Copper(II) Nitride Chromium(III) Sulfite Name the following compounds: K2SO4 Fe(NO3)3 Ca3P2 SnO2

Writing/Naming Molecular Compounds Mono- 1 Di- 2 Tri- 3 Tetra- 4 Penta- 5 Hexa- 6 Hepta- 7 Octa- 8 Nona- 9 Deca- 10 Molecular Compounds Contain covalent bonds. Writing Formulas Write the symbol for the name of each atom. Use a subscript equivalent to the meaning of the prefix. Table of Prefixes, page 246.

Example Write the formula for Carbon Monoxide. CO No prefix on the first word, indicates 1 atom. Prefix mono-, indicate 1 atom (Used only on second word) Write the formula for Carbon Dioxide. CO2

Practice Writing Formulas Write the formulas for the following compounds: Dinitrogen Pentoxide Phosphorus Trichloride Sulfur Trioxide Carbon Tetrachloride

Naming Molecular Compounds Use the same prefixes on page 246. First word Prefix + Name of element, only uses a prefix if there are 2 or more atoms. Second word MUST HAVE a prefix followed by name ending in –ide.

Practice Naming Write the name for the following formulas: P4O10 N2O5 CF4 N2O3 SO2

Binary Acids / Oxyacids All acids contain the element H. H is always written 1st in any formula representing acids. Binary Acids Contain Hydrogen + a non-metal. Oxyacids Contain Hydrogen + a polyatomic ion

Binary Acids Naming: Example Must begin with the prefix Hydro- Must end in the suffix –ic. Place the nonmetal between the prefix and suffix. Example HCl Hydrochloric Acid

Binary Acids Writing: Example Always place the H first. Identify the nonmetal in the acid. Cross the charges on the H 1+ and the nonmetal as subscripts. Example Hydrosulfuric Acid H +1 S 2- H2S

Binary Acids Practice Write the name for each of the following acids: HI H2S H3N Write the formula for the following acids: Hydrofluoric acid Hydrophosphoric acid Hydroselenic acid

Oxyacids Naming: Example Identify the polyatomic ion after the H. Write the name of the non-metal in the polyatomic ion. If the polyatomic ion ends in: -ate  change to –ic -ite  change to –ous Example H2SO4 Sulfate  Sulfuric Acid H2SO3 Sulfite  Sulfurous Acid

Oxyacids Writing: Example Identify the root of the polyatomic ion. If the acid ends in: - ic  use the –ate ending - ous  use the –ite ending Place the H in front of the polyatomic ion and cross the charges if needed. Example Phosphoric Acid Phosphoric = Phosphate (PO4) 3- H3PO4

Oxyacids Practice Write the name for the following acids: HNO3 H2CO3 HC2H3O2 Write the formula for the following acids: Nitrous acid Perchloric acid Phosphorous acid

Hydrates Ionic compounds that contain absorbed water. Example MgSO4  5H2O Magnesium Sulfate Pentahydrate

Summary of Naming Compounds 1st element H, then its an acid: 2 or more elements after the H Oxyacid 1 element after the H Binary acid 2 elements, 1st element is NOT hydrogen. Both are non-metals Molecular compound, use prefixes. 1st is a metal Binary Ionic, check to see if the 1st element is multiple charged. 3 or more elements, 1st element is NOT hydrogen. Ternary Ionic Contains at least 1 polyatomic ion. 1st element a metal, check to see if it is a multiple charge.

Summary of Writing Formulas Most formulas use only 2 words. Check the 2nd word ending first! Ending is –ic or –ous Compound is an acid. Has –ic ending or –ous without a Hydro- prefix Oxyacid Contains a Hydrogen + Polyatomic Ion. Has –ic ending with a Hydro- prefix Binary Acid Contains a Hydrogen + Non-metal.

Writing Formulas Continued Ending is –ide. If it ends in Hydroxide or Cyanide. Ternary Ionic, check charges and cross. Reminder!! More than 1 polyatomic ion, use ( ). If it isn’t Hydroxide or Cyanide. Binary Compound Check if prefixes are used on the 2nd word. Prefix  then Binary molecular, just write subscripts = to prefixes. No Prefix  Binary Ionic, check charges and cross if needed. If it ends in –ite or –ate. More than 1 polyatomic ion needed, use ( ).