Chapters 4 5 Chemical Bonding Valence Electrons Outermost

Chapters 4 & 5 Chemical Bonding

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

Problems • 1) 2) 3) 4) 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 • Ne

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

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

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 2) If the metal is a transition metal with a variable charge, indicate its charge with a Roman Numeral in parentheses 3) Next, name the anion and change its ending to “ide” 4) If the anion is polyatomic, do not change the ending to “-ide” 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: 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

Write the Formula for the following ionic compounds: 8) Sodium Fluoride 9) Calcium Sulfite 10) Calcium Chloride 11) Iron (III) Oxide 12) Cobalt (II) Hydroxide 13) Ammonium Bromide 14) Ammonium Carbonate 15) 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: – 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 6 7 8 9 10 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: 1) CO 2) NI 3 3) N 2 O 4) SF 6 5) B 2 O 3

Write the formula for the following compounds: 6) Phosphorous Pentachloride 7) Nitrogen Monoxide 8) Dinitrogen Tetroxide 9) Tetraphosphorous Decoxide

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” 1) 2) 3) 4) 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 Methane Space-Filling Models

Ethanol

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. • 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 (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

Problems Draw the LDS’s for the following molecules: 1) Cl 2 O 2) C 2 H 4 3) C 2 H 6 O

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 • Nitrogen – Likes to have 3 single bonds and one lone pair

• Phosphorous – Likes to do what nitrogen does • Hydrogen – Likes to be terminal with only one single bond – No lone pairs!

Problems 1) SH 2 2) C 3 H 8 3) Si 2 H 6 4) PI 3 5) CH 3 OH 6) C 2 H 2

7) CCl 2 O 8) N 2 H 4 9) CH 2 OS 10) C 2 H 6 O 11) CO 12) Br. HO

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 • 1) 2) 3) 4) 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 Br. I 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: AX 2 E 0 • A: Central Atom • X: Bonding pairs • E: Non-bonding pairs – Example: Be. Cl 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 Lewis Dot Structures for the following molecules and then identify the direction of polarity, if any. 1) 2) 3) 4) 5) NF 3 CH 2 O CBr 4 CHCl 3 CH 2 Cl 2