Bonding and Molecular Structure

Slides:



Advertisements
Similar presentations
BONDING Ch 8 & 9 – Honors Chemistry General Rule of Thumb:
Advertisements

Bonding By John Patrick Fahy III of Galway. Coulomb’s Law Attractive force is proportional to (+q)(-q)/r^2 +q = magnitude of the positive charge -q =
1 Lewis Dot Formulas of Atoms Lewis dot formulas or Lewis dot representations are a convenient bookkeeping method for tracking valence electrons. Valence.
HONORS REVIEW.  What are valence electrons?  What are valence electrons Electrons in the outermost energy level.
TUTORIAL INSTRUCTIONS: Carefully go through the tutorial, step by step. You may return to any section, as necessary. Once you are satisfied that you understand.
1 CHAPTER 7 Chemical Bonding. 2 Chapter Goals 1. Lewis Dot Formulas of Atoms Ionic Bonding 2. Formation of Ionic Compounds Covalent Bonding 3. Formation.
Copyright © Houghton Mifflin Company. All rights reserved. 12 | 1 Chemical Bonds Forces that hold atoms together Ionic bonds: the forces of attraction.
© 2006 Brooks/Cole - Thomson Writing Lewis Formulas: The Octet Rule The octet rule states that representative elements usually attain stable noble gas.
Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Bonds Forces that hold groups of atoms together and make them function as a unit.
Example: CO3 , carbonate ion
© 2006 Brooks/Cole - Thomson Writing Lewis Formulas: The Octet Rule The octet rule states that representative elements usually attain stable noble gas.
Chemical Bonding Warm-up What determines the reactivity of a metal? What determines the reactivity of a non-metal?
Chapter 8.  Why do atoms bond?  Atoms bond to become more stable.  Atoms are most stable when they have 8 valence electrons. (i.e. the same electron.
Lewis Structures & VSEPR. Lewis Structure Lewis Structures – shows how the _______________ are arranged among the atoms of a molecule There are rules.
Chemical Bonding Chapter 6 Sections 1, 2, and 5. Chemical Bonds A chemical bond is the mutual electrical attraction between the nuclei and valence electrons.
Forces that hold atoms together.  There are several major types of bonds. Ionic, covalent and metallic bonds are the three most common types of bonds.
Zumdahl • Zumdahl • DeCoste
Chapter 11 Chemical Bonds: The Formation of Compounds from Atoms Objectives: Describe the trends in the periodic table Describe the trends in the periodic.
Chapter 5 Compounds and Their Bonds
BONDING General Rule of Thumb: metal + nonmetal = ionic
Covalent Bonding The joining of two or more elements through the sharing of valance electrons to form a molecule Purpose: To form a stable octet between.
Covalent Bonding. Lesson 1:Covalent Bonding Covalent bonds: atoms held together by sharing electrons. Molecules: neutral group of atoms joined together.
I Chemical Bonding. Chemical Bond  attractive force between atoms or ions that binds them together as a unit  bonds form in order to…  decrease potential.
CHAPTER 9 Molecular Structure & Covalent Bonding Theories.
Chapter 6.2 and 6.5 Covalent Compounds.
Bonding General Concepts
Chemical Bonding Chapter 6. Types of Chemical Bonds  Chemical Bond: mutual electrical attraction b/ the nuclei and valence e - of different atoms  Atoms.
Chemical Bonding.  Attractive forces that hold atoms together in compounds. The electrons involved in bonding are usually those in the outermost (valence)
Introductory Chemistry: Concepts & Connections Introductory Chemistry: Concepts & Connections 4 th Edition by Charles H. Corwin Chemical Bonding Christopher.
1 CHAPTER 8 Chemical Bonding. 2 Chapter Goals 1. Lewis Dot Formulas of Atoms Ionic Bonding 2. Formation of Ionic Compounds Covalent Bonding 3. Formation.
Chapter 8. Two Simple Theories of Covalent Bonding  Valence Shell Electron Pair Repulsion Theory __________ R. J. Gillespie ’s  Valence Bond Theory.
Chemical Bonding Chapter 11
Unit 5-K: Chemical Bonding
Section 12.1 Characteristics of Chemical Bonds 1.To learn about ionic and covalent bonds and explain how they are formed 2.To learn about the polar covalent.
Chemical Bonding. Although we have talked about atoms and molecules individually, the world around us is almost entirely made of compounds and mixtures.
Covalent Compounds Chapter 8. Section 1, Covalent Bonds –Remember, ionic compounds are formed by gaining and losing electrons –Atoms can also share electrons.
Covalent Compounds Chapter Covalent Bonds. Covalent Bond The sharing of electrons between atoms Forms a molecule To have stable (filled) orbitals.
Bonding and Molecular Structure: Fundamental Concepts
CHAPTER 8 Basic Concepts in Chemical Bonding. Introduction Attractive forces that hold atoms together in compounds are called chemical bonds. The electrons.
Chemical Bonding. Chemical Bonds A bond is a force that holds groups of two or more atoms together and makes them function as a unit. A bond is a force.
Bonding Dr. Ron Rusay Fall 2001 © Copyright 2001 R.J. Rusay.
1 Chemical Bonds The Formation of Compounds From Atoms Chapter 11 Hein and Arena.
Unit 6: Chemical Bonding Refer to Ch. 8 & 9 for supplemental reading.
Chapter #7 Chemical Bonds.. Chemical Bond An attractive force that holds two atoms together in a complex unit. Electrons combine to form chemical bonds.
Chemical Bonding b Chapter 6. Chemical bond b The force (electrical attraction) that binds two atoms together.
Unit 11 - Bonding Types of Chemical Bonds Electronegativity Bond Polarity and Dipole Moments Stable Electron Configurations Lewis Structures Lewis Structures.
Chemical bonds. Bonding, the way atoms are attracted to each other to form molecules, determines nearly all of the chemical properties we see. Chemical.
Chemical Bonding. Chemical bond: attractive force holding two or more atoms together. Covalent bond results from sharing electrons between the atoms.
Unit 6: Chemical Bonding and Intermolecular Forces
Chemical Bonding.
Bond and Lone Pairs Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. • •• H Cl shared or bond pair lone pair (LP)
Basic Concepts in Chemical Bonding
5.1 Ionic Bonds: Chemical Bonding
Chapter 6 Table of Contents Section 1 Covalent Bonds
Chapter 12 Chemical Bonding.
CHEMICAL BONDING Cocaine Chemistry I – Chapter 8
Chemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals Chapter 10.
Molecular Geometry Cocaine
Chemical Bonding.
CHEMICAL BONDING Cocaine
CHAPTER 6 Chemical Bonding.
7 Chemical Bonding.
A. Types of Chemical Bonds
ChemicalBonding Honors Only Problems and questions —
CHEMICAL BONDING Cocaine Chemistry I – Chapter 8
CHEMICAL BONDING By Mrs. Idha Nurhayati, SPd. Cocaine
Bonding: General Concepts
Covalent Bonding.
CHEMICAL BONDING Cocaine Chemistry I – Chapter 8
Presentation transcript:

Bonding and Molecular Structure CHAPTER 8 Bonding and Molecular Structure

Introduction Bonds: Attractive forces that hold atoms together in compounds Valence Electrons: the outermost electrons -These e- are involved in bonding

Valence Electrons Electrons are divided between core and valence electrons B 1s2 2s2 2p1 Core = [He] , valence = 2s2 2p1 Br [Ar] 3d10 4s2 4p5 Core = [Ar] 3d10 , valence = 4s2 4p5

Valence Electrons -The number of valence electrons of a main group atom is the Group number -For Groups IA-IVA, number of bonding (unpaired) electrons is equal to the group number -For Groups VA -VIIA, number of bonding (unpaired) electrons is equal to 8 - group number

Valence Electrons -Except for H (and sometimes atoms of the 3rd group and higher) -The total number of valence electrons around a given atom in a molecule will be eight: OCTET RULE - (with the exception of hydrogen) atoms in molecules prefer to be surrounded by 8 electrons (or have 4 bonds = 8 electrons)

Lewis Dot Formulas of Atoms IA IIA IIIA IVA VA VIA VIIA VIIIA H He Li Be B C N O F Ne

Ionic Bonding An ion is an atom or a group of atoms possessing a net electrical charge -cations: positive (+) ions These atoms have lost 1 or more electrons -anions: negative (-) ions These atoms have gained 1 or more electrons

Formation of Ionic Compounds Monatomic ions consist of one atom Examples: Na+, Ca2+, Al3+ - cations Cl-, O2-, N3- -anions Polyatomic ions contain more than one atom NH4+ - cation NO2-,CO32-, SO42- - anions

Formation of Ionic Compounds General trend: metals become isoelectronic with the preceding noble gas electron configuration nonmetals become isoelectronic with the following noble gas electron configuration

Formation of Ionic Compounds Reaction of Group IA Metals with Group VIIA Nonmetals

Formation of Ionic Compounds 1s 2s 2p Li  . . F    These atoms form ions with these configurations. Li+  . same configuration as [He] F- same configuration as [Ne]

Formation of Ionic Compounds In general: the reaction of IA metals and VIIA nonmetals: 2 M(s) + X2  2 MX(s) where M is the metals Li to Cs and X is the nonmetals F to I Electronically it looks like: ns np ns np M   M+ __ __ __ __ X    X-   

Formation of Ionic Compounds reaction of IIA metals with VIIA nonmetals: Be(s) + F2(g) BeF2(g)

Formation of Ionic Compounds The valence electrons in these two elements react like: 2s 2p 2s 2p Be [He]  __ __ __  Be2+ __ __ __ __ F [He]     F-    Lewis dot structure representation:

Formation of Ionic Compounds The remainder of the IIA metals and VIIA nonmetals react similarly: M(s) + X2  MX2 M can be any of the metals Be to Ba X can be any of the nonmetals F to I

Formation of Ionic Compounds For the reaction of IA metals with VIA nonmetals: Draw the valence electronic configurations for Li, O, and their appropriate ions

Formation of Ionic Compounds Draw the electronic configurations for Li, O, and their appropriate ions You do it! 2s 2p 2s 2p Li [He]   Li1+ O [He]    O2-    Draw the Lewis dot formula representation of this reaction

Formation of Ionic Compounds Simple Binary Ionic Compounds Table Reacting Groups General Formula Example IA + VIIA MX NaF IIA + VIIA MX2 BaCl2 IIIA + VIIA MX3 AlF3 IA + VIA M2X Na2O IIA + VIA MX BaO IIIA + VIA M2X3 Al2S3

Formation of Ionic Compounds Reacting Groups General Formula Example IA + VA M3X Na3N IIA + VA M3X2 Mg3P2 IIIA + VA MX AlN -H forms ionic compounds when bound to metals (IA and IIA metals For example: LiH, KH, CaH2, and BaH2 -When H is bound to nonmetals, the compounds are covalent in nature

Formation of Covalent Bonds -potential energy of an H2 molecule as a function of the distance between the two H atoms

Covalent Bonding Atoms share electrons If the atoms share: 2 electrons a single covalent bond is formed 4 electrons - a double covalent bond 6 electrons - a triple covalent bond Atoms have a lower potential energy when bound…this is a more favorable situation (why?)

Writing Lewis Formulas: 1. Add the number of valence electrons for all the atoms that are present in the molecule 2. Add or subtract electrons based on the molecule’s (or ion’s) charge 3. Identify the central atom and draw a skeletal structure: -the one that requires the most e- to complete octet -the less electronegative 4. Place a bond between each atom (1 bond = 2 e-) 5. Fill in octet of outer atoms first 6. Finish by completing the octet of central atom – if you run out of e- then multiple bonds must be created between the central atom and atoms bound to it

Writing Lewis Formulas octet rule: representative elements usually attain stable noble gas electron configurations (8 valence e-) in most compounds You must distinguish the difference between: -bonding electrons and nonbonding electrons -shared (paired) and unshared (unpaired) electrons

Formation of Covalent Bonds Lewis dot structures: 1. H2 molecule formation: 2. HCl molecule formation:

Lewis Structures Homonuclear diatomic molecules 1. Two atoms of the same element, H2: 2. Fluorine, F2: Nitrogen, N2:

heteronuclear diatomic molecules Lewis Structures heteronuclear diatomic molecules 1. hydrogen fluoride, HF 2. hydrogen chloride, HCl 3. hydrogen bromide, HBr

Lewis Structures Water, H2O Ammonia molecule , NH3

Lewis Structures Polyatomic ions: ammonium ion NH4+ Notice that the N-atom in this molecule has eight electrons around them (H does not)

Writing Lewis Formulas Sulfite ion, SO32-.

Double and even triple bonds are commonly observed for C, N, P, O, and S H2CO SO3 C2F4

Lewis Structures Example: Write Lewis dot and dash formulas for sulfur trioxide, SO3

Resonance There are three possible structures for SO3: -Two or more Lewis formulas are necessary to show the bonding in a molecule -use equivalent resonance structures to show the molecule’s structure -Double-headed arrows are used to indicate resonance formulas

Resonance Resonance is a flawed method of representing molecules -There are no single or double bonds in SO3

Sulfur Dioxide, SO2 1. Central atom = 2. Valence electrons = ___ or ___ pairs 3. Write the Lewis structure 4. Form double bond so that S has an octet — but note that there are two ways of doing this.

Limitations of the Octet Rule There are some molecules that violate the octet rule: - Be - Group IIIA -Odd number of total electrons. -Central element must have a share of more than 8 valence electrons to accommodate all of the substituents. (i.e. S and P)

Limitations of the Octet Rule Example: Write Lewis formula for BBr3.

Sulfur Tetrafluoride, SF4 Central atom = Valence electrons = ___ or ___ pairs. Form sigma bonds and distribute electron pairs. 5 pairs around the S atom. A common occurrence outside the 2nd period.

Limitations of the Octet Rule Example: Write dot structures for AsF5.

Formal Atomic Charges Atoms in molecules often bear a charge (+ or -) The predominant resonance structure of a molecule is the one with charges on atoms as close to 0 as possible Formal charge = Group number – 1/2 (# of bonding electrons) - (# of Lone electrons) = Group number – (# of bonds) – (# of Lone electrons)

Formal Charge CO2 . . . . . . . .

Which is the most stable resonance form? Formal Charge Thiocyanate Ion, SCN- • S N C • S N C • S N C Which is the most stable resonance form?

Theories of Covalent Bonding Valence Shell Electron Pair Repulsion Theory Commonly designated as VSEPR Principal originator R. J. Gillespie in the 1950’s Valence Bond Theory (Chapter 9) Involves the use of hybridized atomic orbitals L. Pauling in the 1930’s & 40’s

VSEPR Theory electron densities around the central atom are arranged as far apart as possible to minimize repulsions (why?) Five basic molecular shapes: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral

VSEPR Theory Two regions of high electron density around the central atom.

VSEPR Theory 2. Three regions of high electron density around the central atom.

VSEPR Theory 3. Four regions of high electron density around the central atom.

VSEPR Theory 4. Five regions of high electron density around the central atom.

VSEPR Theory 5. Six regions of high electron density around the central atom.

Molecular geometry: arrangement of atoms around the central atom(s) VSEPR Theory Electronic geometry(family): locations of regions of electron density around the central atom(s) Molecular geometry: arrangement of atoms around the central atom(s) Electron pairs are not used in the molecular geometry determination

VSEPR Theory Lone pairs (unshared pairs) of electrons require more volume than shared pairs -there is an ordering of repulsions of lone electrons around central atom Criteria for the ordering of the repulsions: 1. Lone pair to lone pair is the strongest repulsion. 2. Lone pair to bonding pair is intermediate repulsion. 3. Bonding pair to bonding pair is weakest repulsion.

Molecular Shapes and Bonding Symbolism: A = central atom B = bonding pairs around central atom U = lone pairs around central atom For example: AB3U designates that there are 3 bonding pairs and 1 lone pair around the central atom

Linear Electronic Geometry: AB2 Some examples of molecules with this geometry: BeCl2, BeBr2, BeI2, HgCl2, CdCl2

Trigonal Planar Electronic Geometry: AB3 Some examples of molecules with this geometry are: BF3, BCl3

Tetrahedral Electronic Geometry: AB4 Some examples of molecules with this geometry are: CH4, CF4, CCl4, SiH4, SiF4

VSEPR Theory An example of a molecule that has the same electronic and molecular geometries is methane (CH4) -Electronic and molecular geometries are tetrahedral

Tetrahedral Electronic Geometry: AB4

Tetrahedral Electronic Geometry: AB3U Some examples of molecules with this geometry are: NH3, NF3, PH3, PCl3, AsH3 -trigonal pyramidal -electronic and molecular geometries are different. . . . . 107.5°

Tetrahedral Electronic Geometry: AB2U2 Some examples of molecules with this geometry are: H2O, OF2, H2S -bent -electronic and molecular geometries are different 104.5°

VSEPR Theory An example of a molecule that has different electronic and molecular geometries is water (H2O) -Electronic geometry is tetrahedral -Molecular geometry is bent or angular

Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3 Some examples of molecules with this geometry are: PF5, AsF5, PCl5 axial equatorial axial

Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3 If lone pairs are incorporated into the trigonal bipyramidal structure, there are three possible new shapes: One lone pair - Seesaw shape Two lone pairs - T-shape Three lone pairs – linear The lone pairs occupy equatorial positions first: -they are 120o from each other -90o from the axial positions Results in decreased repulsions compared to lone pair in axial position axial equatorial axial

Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3 AB4U molecules have: trigonal bipyramid electronic geometry seesaw shaped molecular geometry polar One example of an AB4U molecule is SF4

Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3

Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3 AB3U2 molecules have: 1. trigonal bipyramid electronic geometry T-shaped molecular geometry polar One example of an AB3U2 molecule is IF3

Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3

Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3 AB2U3 molecules have: trigonal bipyramid electronic geometry linear molecular geometry nonpolar One example of an AB3U2 molecule is BrF2-

Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3

Octahedral Electronic Geometry: AB6, AB5U, and AB4U2 Some examples of molecules with this geometry are: SF6, SeF6, SCl6, etc.

Octahedral Electronic Geometry: AB6, AB5U, and AB4U2 If lone pairs are incorporated into the octahedral structure, there are two possible new shapes: One lone pair - square pyramidal Two lone pairs - square planar The lone pairs occupy any position because they are all 90o from all bonds positions: -Additional lone pairs occupy the position 180º from the first set of lone pairs -This results in decreased repulsions compared to lone pairs in the other positions

Octahedral Electronic Geometry: AB6, AB5U, and AB4U2 AB5U molecules have: octahedral electronic geometry Square pyramidal molecular geometry polar. One example of an AB4U molecule is IF5

Octahedral Electronic Geometry: AB6, AB5U, and AB4U2 AB4U2 molecules have: octahedral electronic geometry square planar molecular geometry and are nonpolar. One example of an AB4U2 molecule is XeF4

Polarity and Electronegativity Figure 8.11

Dipole Moments For example, HF and HI:

Dipole Moments some “nonpolar molecules” that have polar bonds Two conditions to be polar: 1. There must be at least one polar bond present or one lone pair of electrons 2. the molecule must be nonsymmetric Examples: water, CF4, CO2, NH3, NH4+

Polar Molecules Molecular geometry affects molecular polarity -they either cancel or reinforce each other A B A A B A linear molecule nonpolar angular molecule polar

Polar and Nonpolar Bonds Covalent bonds in which the electrons are shared equally are designated as nonpolar covalent bonds -Nonpolar covalent bonds have a symmetrical charge distribution (electron distribution)

Polar and Nonpolar Bonds Polar covalent bonds: electrons are not shared equally -they have different electronegativities H F Electronegativities: 2.1 4.0 Difference = 1.9 very polar bond

Polar and Nonpolar Bonds Compare HF to HI: H I Electronegativities: 2.1 2.5 Difference = 0.4 slightly polar bond more complicated geometries exist…

Bond Polarity Three molecules with polar covalent bonds: -Each bond has one atom with a slight negative charge (-d) -another with a slight positive charge (+d)

Polar or Nonpolar? AB3 molecules: BF3, Cl2CO, and NH3

Polar or Nonpolar? CO2 and H2O Which one is polar?

CH4 … CCl4 Polar or Not? Only CH4 and CCl4 are NOT polar. These are the only two molecules that are “symmetrical.”

Compounds Containing Double Bonds Ethene or ethylene, C2H4, is the simplest organic compound containing a double bond. -has a double bond to obey octet rule Lewis Dot Formula

Double Bonds What is the effect of bonding and structure on molecular properties? s and p Free rotation around C–C single bond No rotation around C=C double bond

# of bonds between similar pairs of atoms Bond Order # of bonds between similar pairs of atoms Double bond Single bond Acrylonitrile Triple bond

Bond Order Consider NO2-: The N—O bond order = 1.5

Bond Order (a) bond strength (b) bond length Bond order is proportional to two important bond properties: (a) bond strength (b) bond length 745 kJ 414 kJ 110 pm 123 pm

Bond Length the distance between the nuclei of two bonded atoms

Bond length depends on size of bonded atoms H—F Bond Length H—Cl H—I Bond distances measured in Angstrom units where 1 Å = 10-2 pm.

Bond length depends on bond order Bond distances measured in Angstrom units where 1 Å = 10-2 pm.

Bond Strength Measure of the energy required to break a bond See Table 9.10 BOND STRENGTH (kJ/mol) H—H 436 KJ C—C 346 KJ C=C 602 KJ CC 835 KJ NN 945 KJ The GREATER the number of bonds (bond order) the HIGHER the bond strength and the SHORTER the bond.