Chapters 3 & 4 Chemical Bonding Chapter 3

Chapters 3 & 4 Chemical Bonding Chapter 3

Chapters 3 & 4 Chemical Bonding Chapter 3 and 4 Goals/HW

Octet Rule Naming Ionic and covalent compounds Introduction to molecular structures Lewis Dot Structures Molecular geometries Electronegativity and bond polarity Molecular polarity Homework Chapter 3= 33, 35, 41, 43, 55, 61 Chapter 4 = 27, 29, 31, 35, 43, 53, 55, 57, 71, 73 Octet Rule

Octet Rule: the tendency for atoms to achieve 8 electrons in their valence shell Natural electron configuration of the Noble Gases Done by gaining, losing, or sharing electrons Increases stability, less chance to react afterwords H and He seek a Duet Na = [Ne]3s1 Na+ = [Ne] Cl = [Ne]3s23p5 Cl- = [Ne]3s23p6 = [Ar] To gain or to lose electrons? Goal is to achieve a noble gas electron

configuration Metals lose electrons Nonmetals gain electrons Its easier for a nonmetal (Ex: chlorine) to gain 1 electron vs losing 7 electrons. Its easier for a metal (Ex: magnesium) to gain 2 electron vs losing 6 electrons. 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 Oxidation state oxidation state/number represents the number of electrons lost or gained by an atom of an element in that compound It relates to the charge of the atom as well If a sodium atom loses one electron it will have a positive charge and the oxidation state will be +1

If an oxygen atom gains two electrons it will have a negative charge and the oxidation state will be -2 Oxidation state Formula Unit Oxidation state 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

PO3 3- Phosphite NO3- Nitrate NO2-

Nitrite ClO4- Perchlorate ClO3- Chlorate ClO2- Chlorite

ClO- Hypochlorit e Rules For Naming Ionic Compounds 1) Name the cation by its elemental/polyatomic name Na+Cl- = NaCl = sodium chloride 2) If the metal is a transition metal with a variable charge, use a Roman Numeral in parentheses for its charge FeCl2 = 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 NaNO3 = sodium nitrate 5) Do NOT use prefixes (mono, di, tri etc.) to indicate how many of each atom are present Iron (II) Chloride FeCl2 Iron (III) Chloride FeCl3 Problems

Write the name for the following compounds: 1) KI potassium iodide 2) MgBr2 3) Al2O3 magnesium bromide aluminum oxide iron(II)chloride 4) FeCl2 calcium sulfate

5) CaSO4 barium nitrite 6) Ba(NO2)2 copper(II) nitrate Write the Formula for the following ionic compounds: NaF 8) Sodium Fluoride 9)

Calcium Sulfite 10) Calcium Chloride 11) Iron (III) Oxide CaSO3 CaCl2 Fe2O3 12) Cobalt (II) Hydroxide Co(OH)2

13) Ammonium Bromide NH4Br 14) Ammonium Carbonate (NH4)2CO3 15) Aluminum Carbonate Al2(CO3)3

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- = NaCl = sodium chloride 2) If the metal is a transition metal with a variable charge, use a Roman Numeral in parentheses for

its charge FeCl2 = 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 NaNO3 = 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: CO2 Cl2 CH4

Dihydrogen Monoxide (DHMO)!!!!!! Properties Deadly, colorless, odorless liquid Hazards!!! Can mutate DNA!! Death due to accidental inhalation of DHMO, even in small quantities. Prolonged exposure to solid DHMO causes severe tissue damage DHMO is a major component of acid rain Contributes to soil erosion.

Dihydrogen Monoxide H2O, commonly known as water Can mutate DNA!! water can get ionized to react with DNA Death due to accidental inhalation of DHMO, even in small quantities--- when someone drowns Prolonged exposure to solid DHMO causes severe tissue damage--- Ice DHMO is a major component of acid rain Contributes to soil erosion--- water corrodes the landscape over time

Covalent Bonding 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

mono 6 hexa 3) If only1one atom, dont put mono 2 3 4 5 di tri tetra penta

7 8 9 10 hepta octa nona deca 4) Name the 2nd non-metal and change its ending to -ide

5) Add a prefix to indicate how many Problems Write the name of the following compounds: carbon monoxide 1) CO 2) NI3 3) N2O 4) SF6 nitrogen triiodide dinitrogen monoxide sulfur hexaflouride

diboron trioxide 5) B2O3 Write the formula for the following compounds: 6) Phosphorous Pentachloride PCl5 7) Nitrogen Monoxide NO

8) Dinitrogen Tetroxide 9) Tetraphosphorous Decoxide N2O4 P4O10 Problems 1) KCl 2) Na2S 3) H2O 4) SO2 5) K3PO4 6) FeCl3

7) (NH4)2SO4 8) SCl2 9) Cu(OH)2 10) P2O5 Potassium chloride Sodium sulfide Dihydrogen monoxide Sulfur dioxide Potassium phosphate Iron(III) chloride Ammonium sulfate Sulfur dichloride

Copper(II) hydroxide Copper(II) hydroxide 8) Sodium Iodide NaI 9) Aluminum Sulfate Al2(SO4) 10) Phosphorous Pentabromide

PBr5 11) Magnesium Nitride Mg3N2 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 = hydrochloric acid HF = hydrofluoric acid EXCEPTION, if in the gas phase, treat like a regular covalent compound for naming with no prefixes HCl(g) = hydrogen chloride

Acids that contain oxygen/oxyanions 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

HClO4 HClO3 HClO2 HClO perchloric acid chloric acid chlorous acid hypochlorous acid

Problems Name the following 1) 2) 3) 4) 5) 6) 7) 8) 9)

10) HBr(g) HBr(aq) HNO2(aq) HNO3(aq) HI (aq) HI (g) H2CO3 (aq) H3PO4 (aq) H3PO3 (aq) HCN (aq)

Hydrogen bromide Hydrobromic acid Nitrous acid Nitric acid Hydroiodic acid Hydrogen iodide Carbonic acid Phosphoric acid Phosphorous acid hydrocyanic acid Molecular Structures

Ball & Stick Models Water H2O Methane CH4 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: PCl3 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 dont 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). 6) Ex: CH2O

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 Cl Cl C Cl Cl 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 Cl Cl Cl Likes to bond to other carbons Cl C C C C C Cl Cl Silicon

Likes to do what carbon does Cl Cl Si Cl Cl Oxygen Likes to have two single bonds and 2 lone pairs H O H Sulfur Likes to do what oxygen does May expand its octet

H S H Nitrogen Likes to have 3 single bonds and one lone pair H H N C C H H N H H O O S O H O

H Phosphorous Likes to do what nitrogen does H May expand its octet H C H H H H P H Hydrogen Likes to be terminal with only one single bond

H No lone pairs! N C C H H Boron Likes 3 bonds and no lone pairs (sextet) H H B H O H P C H

H C H H Problems Draw the LDSs for the following molecules: 1)Cl2O 2)C2H4 3)C2H6O

Problems Draw the Lewis Structures for the following molecules: 1) SH2 2) C3H8 3) 5)

CH3OH 6) C2H2 Si2H6 7) BF3 9) N2 H4

10) CH2OS 11) C2H6O 12) P4 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: Cl2, H2, CH4 ENCl = 3.0 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, ClF, 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

Draw the lewis dot structure then 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) BrI 4) O2 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: AX2E0 A: Central Atom X: Bonding pairs

E: Non-bonding pairs Example: BeCl2, CO2 O C O 3 e- groups 3 Bonds, 0 Lone Pairs 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 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 Draw the 3D Lewis Dot Structures, using wedges and dashes when applicable, for the following molecules and then identify the net dipole, if any. 1) BF3 2) CH2O 3)

CBr4 4) CHCl3 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, 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

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