Chapters 3 & 4 Chemical Bonding

Valence Electrons Outermost electrons s and p electrons for main group elements Responsible for chemical properties of atoms Participate in chemical reactions Valence Electron Core Electrons

Problems Write out the electron configurations for the following elements and identify how many core and valence electrons each has. Mg S Br Kr

Lewis Dot Structures LDS: a representation of an atom using its chemical symbol surrounded by dots that signify valence electrons

Problems Write the Lewis Dot Structures for the following atoms Li Be Br C N Ne

Li: [He]1s1 Na: [Ne]2s1 K: [Ar]3s1

Octet Rule Octet Rule: the tendency for atoms to seek 8 electrons in their outer shells Natural electron configuration of the Noble Gases Done by gaining, losing, or sharing electrons Increases stability H and He seek a “Duet”

Ionic Bonding Ions: atoms that have a charge due to gain or loss of electrons Anion: (-) charged atom Cation: (+) charged atom Ionic Bond: a bond formed through the transfer of one or more electrons from one atom or group of atoms to another atom or group of atoms

Formula Unit

Polyatomic Ion + Non-metal Polyatomic Ion + Polyatomic Ion Ionic Compounds: compounds composed of oppositely charged ions that are held together by their attraction to each other Metal + Non-metal NaCl Metal + Polyatomic Ion NaNO3 Polyatomic Ion + Non-metal NH4Cl Polyatomic Ion + Polyatomic Ion NH4NO3 Net charge on compound equal to zero

Oxyanions SO42- Sulfate SO32- Sulfite PO43- Phosphate PO33- Phosphite NO3- Nitrate NO2- Nitrite ClO4- Perchlorate ClO3- Chlorate ClO2- Chlorite ClO- Hypochlorite

Rules For Naming Ionic Compounds Name the cation by its elemental/polyatomic name If the metal is a transition metal with a variable charge, indicate its charge with a Roman Numeral in parentheses Next, name the anion and change its ending to “-ide” If the anion is polyatomic, do not change the ending to “-ide” Do NOT use prefixes (mono, di, tri etc.) to indicate how many of each atom are present

Problems Write the name for the following compounds: KI MgBr2 Al2O3 FeCl2 CaSO4 Ba(NO2)2 Cu(NO3)2

Write the Formula for the following ionic compounds: Sodium Fluoride Calcium Sulfite Calcium Chloride Iron (III) Oxide Cobalt (II) Hydroxide Ammonium Bromide Ammonium Carbonate Aluminum Carbonate

Iron (II) Chloride Iron (III) Chloride

Covalent Compounds Covalent Compounds: compounds composed of atoms bonded to each other through the sharing of electrons Electrons NOT transferred No + or – charges on atoms Non-metal + Non-metal Also called “molecules” Examples: H2O CO2 Cl2 CH4

or H-H Duet or

Naming Covalent Compounds Name the first non-metal by its elemental name Add a prefix to indicate how many 1 6 2 7 3 8 4 9 5 10 Name the 2nd non-metal and change its ending to “-ide” Add a prefix to indicate how many

Problems Write the name of the following compounds: CO NI3 N2O SF6 B2O3

Write the formula for the following compounds: Phosphorous Pentachloride Nitrogen Monoxide Dinitrogen Tetroxide Tetraphosphorous Decoxide

Problems KCl Na2S H2O SO2 K3PO4 FeCl3 (NH4)2SO4 SCl2 Cu(OH)2 P2O5

Sodium Iodide Aluminum Sulfate Phosphorous Pentabromide Magnesium Nitride

Naming Acids Acids that do not contain oxygen HCl HF Begin the name with “hydro” Name the anion, but change the ending to “-ic” Add “acid” on the end HCl HF

Acids that contain oxygen Do not put “hydro” at the beginning Begin the name with the anion If the anion has the ending “-ate,” change this to “-ic acid” If the anion has the ending “-ite,” change this to “-ous acid” HClO4 HClO3 HClO2 HClO

Problems Name the following HBr(g) HBr(aq) HNO2(aq) HNO3(aq) HI (aq) HI (g) H2CO3 (aq) H3PO4 (aq) H3PO3 (aq) HCN (aq)

Molecular Structures

Ball & Stick Models Space-Filling Models Water Methane

Ethanol

Lewis Dot Structures Count the total number of valence electrons in the molecule. Ex: PCl3 Use atomic symbols to draw a proposed structure with shared pairs of electrons. Atoms don’t tend to bond to other atoms of the same element when they can avoid it Exception: Carbon

Place any leftover electrons on the central atom Place lone pair electrons around each (except H) to satisfy the octet rule, beginning with the terminal atoms Place any leftover electrons on the central atom If the number of electrons around the central atom is less than 8, change single bonds to the central atom to multiple bonds (double or triple). Ex: CH2O

Problems Draw the LDS’s for the following molecules: Cl2O C2H4 C2H6O

What Things Like To Do Halogens Like to be terminal Like to have one single bond and 3 lone pairs (non-bonding electrons) Carbon Likes to have 4 single bonds and no lone pairs A double bond counts as two singles A triple bond counts as three singles Likes to be central Likes to bond to other carbons

Silicon Likes to do what carbon does Oxygen Likes to have two single bonds and 2 lone pairs Sulfur Likes to do what oxygen does May expand its octet Nitrogen Likes to have 3 single bonds and one lone pair

Phosphorous Likes to do what nitrogen does May expand its octet Hydrogen Likes to be terminal with only one single bond No lone pairs! Boron Likes 3 bonds and no lone pairs (sextet)

Problems Draw the Lewis Structures for the following molecules: SH2 C3H8 Si2H6 PI3 CH3OH C2H2 BF3 CCl2O

N2H4 CH2OS C2H6O CO BrHO

Electronegativity The measure of the ability of an atom to attract electrons to itself Increases across period (left to right) and Decreases down group (top to bottom) fluorine is the most electronegative element francium is the least electronegative element

Electronegativity Scale

Types of Bonding Non-Polar Covalent Bond: Difference in electronegativity values of atoms is 0.0 – 0.4 Electrons in molecule are equally shared Examples: Cl2, H2, CH4 ENCl = 3.0 3.0 - 3.0 = 0 Pure Covalent

Polar Covalent Bond: Difference in electronegativity values of atoms is 0.4 – 2.0 Electrons in the molecule are not equally shared The atom with the higher EN value pulls the electron cloud towards itself Partial charges Examples: HCl, ClF, NO ENCl = 3.0 ENH = 2.1 3.0 – 2.1 = 0.9 Polar Covalent

Ionic Bond: Difference in EN above 2.0 Complete transfer of electron(s) Whole charges ENCl = 3.0 ENNa = 1.0 3.0 – 0.9 = 2.1 Ionic

Problems Predict the type of bonding in the following compounds using differences in EN values of the atoms. Indicate the direction of the dipole moment if applicable KBr HF BrI FI

Valence Shell Electron Pair Repulsion Theory VSEPR theory: Electrons repel each other Electrons arrange in a molecule themselves so as to be as far apart as possible Minimize repulsion Determines molecular geometry

Defining Molecular Shape Electron pair geometry: the geometrical arrangement of electron groups around a central atom Look at all bonding and non-bonding e-’s Molecular Geometry: the geometrical arrangement of atoms around a central atom Ignore lone pair electrons

2 e- groups surrounding the central atom e- pair geometry: linear MG: linear AXE designation: AX2E0 A: Central Atom X: Bonding pairs E: Non-bonding pairs Example: BeCl2

3 e- groups 3 Bonds, 0 Lone Pairs 2 Bonds, 1 Lone Pair e- PG: Trigonal Planar (Triangular planar) MG: Trigonal Planar AX3E0 BF3 2 Bonds, 1 Lone Pair e- PG: Trigonal Planar (Triangular planar) MG: Bent/angular AX2E1 GeCl2

4 e- groups 4 bonds, 0 Lone Pairs 3 bonds, 1 Lone Pair e- PG: Tetrahedral MG: Tetrahedral AX4E0 CH4 3 bonds, 1 Lone Pair e- PG: Tetrahedral MG: Triangular Pyramidal AX3E1 NH3

2 bonds, 2 Lone Pairs e- PG: Tetrahedral MG: Bent/Angular AX2E2 H2O

Drawing LDS With Correct Geometry

Molecular Polarity

Problems BF3 CH2O CBr4 CHCl3 CH2Cl2 Draw the 3D Lewis Dot Structures, using wedges and dashes when applicable, for the following molecules and then identify the net dipole, if any. BF3 CH2O CBr4 CHCl3 CH2Cl2