Chapters 3 4 Chemical Bonding Octet Rule Octet

Chapters 3 & 4 Chemical Bonding

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 – added electron(s) – Cation: (+) charged atom – lost electron(s) • 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

Ionic Compounds: compounds composed of oppositely charged ions that are held together by their attraction to each other • Metal + Non-metal – Na. Cl • Metal + Polyatomic Ion – Na. NO 3 • Polyatomic Ion + Non-metal – NH 4 Cl • Polyatomic Ion + Polyatomic Ion – NH 4 NO 3 • Net charge on compound equal to zero

Oxyanions SO 42 - Sulfate SO 32 - Sulfite PO 43 - Phosphate PO 33 - Phosphite NO 3 - Nitrate NO 2 - Nitrite Cl. O 4 - Perchlorate Cl. O 3 - Chlorate Cl. O 2 - Chlorite Cl. O- Hypochlorite

Rules For Naming Ionic Compounds 1) Name the cation by its elemental/polyatomic name Na+Cl- = Na. Cl = sodium chloride 2) If the metal is a transition metal with a variable charge, use a Roman Numeral in parentheses for its charge Fe. Cl 2 = iron(II)chloride 3) Next, name the anion and change its ending to “ide” Cl- = chloride 4) If the anion is polyatomic, do not change the ending to “-ide” Na. NO 3 = sodium nitrate 5) 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: potassium iodide 1) KI 2) Mg. Br 2 3) Al 2 O 3 4) Fe. Cl 2 5) Ca. SO 4 6) Ba(NO 2)2 7) Cu(NO 3)2 magnesium bromide aluminum oxide iron(II)chloride calcium sulfate barium nitrite copper(II) nitrate

Write the Formula for the following ionic compounds: Na. F 8) Sodium Fluoride 9) Calcium Sulfite Ca. SO 3 10) Calcium Chloride Ca. Cl 2 11) Iron (III) Oxide Fe. O 3 12) Cobalt (II) Hydroxide Co(OH)2 13) Ammonium Bromide NH 4 Br 14) Ammonium Carbonate (NH 4)2 CO 3 15) Aluminum Carbonate Al 2(CO 3)3

Iron (II) Chloride Iron (III) Chloride

Chapter 3 review • Octet rule • Ions and ionic bonding • Ionic compounds - memorize polyatomic ions from the tables • How to name ionic compounds - know the rules

Rules For Naming Ionic Compounds 1) Name the cation by its elemental/polyatomic name Na+Cl- = Na. Cl = sodium chloride 2) If the metal is a transition metal with a variable charge, use a Roman Numeral in parentheses for its charge Fe. Cl 2 = iron(II)chloride 3) Next, name the anion and change its ending to “ide” Cl- = chloride 4) If the anion is polyatomic, do not change the ending to “-ide” Na. NO 3 = sodium nitrate 5) Do NOT use prefixes (mono, di, tri etc. ) to indicate how many of each atom are present

Chapter 4 Covalent Compounds Dalton Trans. , 2016, 45, 15481 -15491

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: – H 2 O – CO 2 – Cl 2 – CH 4

or Duet or H-H

Naming Covalent Compounds 1) Name the first non-metal by its elemental name 2) Add a prefix to indicate how many 1 2 3 4 5 mono di tri tetra penta 6 7 8 9 10 hexa hepta octa nona deca 3) Name the 2 nd non-metal and change its ending to “-ide” 4) Add a prefix to indicate how many

Problems Write the name of the following compounds: carbon monoxide 1) CO 2) NI 3 nitrogen triiodide 3) N 2 O dinitrogen monoxide 4) SF 6 Sulfur hexaflouride 5) B 2 O 3 Diboron trioxide

Write the formula for the following compounds: 6) Phosphorous Pentachloride PCl 5 7) Nitrogen Monoxide NO 8) Dinitrogen Tetroxide N 2 O 4 9) Tetraphosphorous Decoxide P 4 O 10

Problems 1) KCl 2) Na 2 S 3) H 2 O 4) SO 2 5) K 3 PO 4 6) Fe. Cl 3 7) (NH 4)2 SO 4 8) SCl 2 9) Cu(OH)2 10) P 2 O 5

8) Sodium Iodide 9) Aluminum Sulfate 10) Phosphorous Pentabromide 11) Magnesium Nitride

Naming Acids • Acids that do not contain oxygen 1) Begin the name with “hydro” 2) Name the anion, but change the ending to “-ic” 3) Add “acid” on the end • • HCl HF

• Acids that contain oxygen 1) Do not put “hydro” at the beginning 2) Begin the name with the anion 3) If the anion has the ending “-ate, ” change this to “-ic acid” 4) If the anion has the ending “-ite, ” change this to “-ous acid” • • HCl. O 4 HCl. O 3 HCl. O 2 HCl. O

Problems • Name the following 1) 2) 3) 4) 5) 6) 7) 8) 9) 10) HBr(g) HBr(aq) HNO 2(aq) HNO 3(aq) HI (g) H 2 CO 3 (aq) H 3 PO 4 (aq) H 3 PO 3 (aq) HCN (aq)

Molecular Structures

Ball & Stick Models Water H 2 O Methane CH 4 Space-Filling Models

Octet Rule or Duet or H-H

Making Lewis Dot Structures 1) Count the total number of valence electrons in the molecule. Ex: PCl 3 2) Use atomic symbols to draw a proposed structure with shared pairs of electrons • Pick a central atom – atom that wants to make the most bonds • Atoms don’t tend to bond to other atoms of the same element when they can avoid it • Exception: Carbon

3) Place lone pair electrons around each outside atom (except H) to satisfy the octet rule, beginning with the terminal atoms 4) Place any leftover electrons on the central atom 5) 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: CH 2 O

Covalence chart Atom H C or Si P or N O or S F, Cl, Br, or I B Covalence number 1 4 3 2 1 3 Non bonding electrons 0 0 1 2 3 0

What Certain Atoms Like To Do • Halogens – Like to have one single bond and 3 lone pairs (non-bonding electrons) – F, Cl, Br, I • 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 LDS’s for the following molecules: 1)Cl 2 O 2)C 2 H 4 3)C 2 H 6 O

Problems Draw the Lewis Structures for the following molecules: 1) SH 2 5) CH 3 OH 2) C 3 H 8 6) C 2 H 2 3) Si 2 H 6 7) BF 3 4) PI 3 8) CCl 2 O

9) N 2 H 4 10) CH 2 OS 11) C 2 H 6 O 12) P 4 13) C 6 H 6

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 1) Non-Polar Covalent Bond: • • • Difference in electronegativity values of atoms is 0. 0 – 0. 4 Electrons in molecule are equally shared Examples: Cl 2, H 2, CH 4 ENCl = 3. 0 - 3. 0 = 0 Pure Covalent

2) 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, Cl. F, NO ENCl = 3. 0 ENH = 2. 1 3. 0 – 2. 1 = 0. 9 Polar Covalent

3) 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 1) KBr 2) HF 3) Br. I 4) 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 – Atoms and lone pairs count as electron groups • 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: AX 2 E 0 • A: Central Atom • X: Bonding pairs • E: Non-bonding pairs – Example: Be. Cl 2, CO 2

3 e groups • 3 Bonds, 0 Lone Pairs – e- PG: Trigonal Planar (Triangular planar) – MG: Trigonal Planar – AX 3 E 0 – BF 3 • 2 Bonds, 1 Lone Pair – e- PG: Trigonal Planar (Triangular planar) – MG: Bent/angular – AX 2 E 1 – Ge. Cl 2

4 e- groups • 4 bonds, 0 Lone Pairs – – e- PG: Tetrahedral MG: Tetrahedral AX 4 E 0 CH 4 • 3 bonds, 1 Lone Pair – – e- PG: Tetrahedral MG: Triangular Pyramidal AX 3 E 1 NH 3

• 2 bonds, 2 Lone Pairs – e- PG: Tetrahedral – MG: Bent/Angular – AX 2 E 2 – H 2 O

Drawing LDS With Correct Geometry

Molecular Polarity

Problems Draw the 3 D Lewis Dot Structures, using wedges and dashes when applicable, for the following molecules and then identify the net dipole, if any. 1) BF 3 2) CH 2 O 3) CBr 4 4) CHCl 3 5) CH 2 Cl 2

Chapter 4 review • Covalent compounds- what is a covalent compound? Know the rules for naming. • Lewis dot structures – how to draw molecules • Electronegativity – What is it and how does it determin the type of bond? Pure covalent, and ionic bonds depend on the electronegativity difference between the each atom in question • Valence shell electron pair repulsion theory • Molecular Geometry and molecular polarity

Chapter 4 Homework Problems • 27, 29, 31, 35, 37, 53, 55, 57, 63, 69, 71, 73, 77, 81, 83, 89, 97