Unit 4:BONDING

Why do elements form bonds????

I. Energy and Bonds

• Elements form bonds to become more stable

• Forming bonds releases energy

• Breaking bonds absorbs energy

• Therefore:– Forming a bond is…– Breaking a bond is…

II. Types of Bonds

• Bond = attractive force that hold elements attractive force that hold elements togethertogether

• There are 3 major types of bonds formed between elements

• Each type of bond has different attractions and different properties

Identifying Bond Types…

• A] Metallic bonds: present within a metal

• B] Ionic bonds: metal + nonmetal– Cations and anions form neutral substances

– Electrons are given/taken to form ions

• C] Covalent: nonmetals sharing electrons– No actual charges formed

III. Lewis Dot Structures

• Lewis dot diagrams show number and relative placement of valence electrons

• Uses element symbol and dots in pattern

1 2

8 3

5 6

7 4

A. Single Elements

• Count number of valence electrons [look at “s” and “p” electrons]

• Place in pattern around element symbol

• Ex.]

B. Ions and Ionic Compounds

• Determine the charge of the ion within the compound [look at oxidation numbers on PT]– Positive ions have NO valence electrons!– Negative ions have 8 valence electrons!

• Arrange with opposite charges connecting

Ex.] NaCl MgCl2

C. Covalent Structures• Determine the number of valence electrons

for each element involved

• Choose a central atom [least popular]

• Organize remaining atoms symmetrically

• Form bonds to provide each element with 8 valence electrons

• May use multiple bonds for each element to see 8

Ex.] H2 O2 N2

D. The Octet Rule

• Octet = eight valence electrons on an atom– valence electrons are those in “s” and “p”

sublevels

• Elements with 8 valence electrons are very stable and usually not reactive!

– What group has all 8 valence electrons naturally?

– Which group of metals is most reactive?

– Which nonmetals are most reactive?

Octets, continued…

• Since all elements want 8 electrons, each atom will gain or lose electrons to “see” 8 valence electrons– Metals lose electrons– nonmetals gain electrons

• Ex.] NaCl MgO

Exceptions to Octet Rule

Some need less than 8:– H, He, B

Some can take more than 8, creating an expanded octet:

– S, P, etc.

IV. Metallic Bonds

• Metallic Bonds = special bonds between the atoms within a metal sample

• Have fixed nuclei with mobile electrons

• “Sea of Mobile Electrons”

• Give metals special properties: – Malleability -- Good Conductivity– Ductility

– http://micro.magnet.fsu.edu/electromag/java/rutherford/

Diagram of Metallic Bonding

http://www.drkstreet.com/resources/metallic-bonding-animation.swf

metallic bonding online demo

V. Covalent Bonds

• Covalent bonds are formed between NONMETALS who share electrons

• Some nonmetals can form more than one bond between the same 2 elements

• Different types of covalent bonding, due to symmetry and electronegativity values

• No formal charge, or ions, formed

Properties of Covalent Compounds

• All phases present at STP

• Low boiling and melting points

• Low density• High vapor pressure,

or volatility• Poor conductors of

heat and electricity

                                

a. Nonpolar Covalent Bonds

• Nonpolar = equal sharing

• Electrons shared within a bond are “seen” equally by both atoms

• Between same atoms ONLY!

b. Multiple Bonds

1]Single Bonds = 2 e- shared between 2 atoms

• 1 e- from each element• Not very strong• Longest covalent bond• Examples:

– some diatomics

Multiple Bonds, cont.

2] Double Bonds = 4 e- shared between 2 atoms

• 2 e- from each element • Stronger and shorter

than single bonds• Examples:

– O2

Multiple bonds cont’.

3] Triple bonds = 6 e- shared between 2 atoms

• 3 e- from each element• Strongest and shortest

bond type• Examples:

– N2

ONLINE RESOURCES

Online resources for understanding bonding geometries of common compounds:

https://phet.colorado.edu/en/simulation/molecule-shapes

https://phet.colorado.edu/en/simulation/molecule-shapes-basics

http://www.pbslearningmedia.org/asset/lsps07_int_covalentbond/

Great Polar covalent bonds link: http://www.chem1.com/acad/webtext/chembond/cb04.html

c. Polar Bonds

Polar bondsPolar bonds are formed between atoms having differences in EN

o Atoms of different EN will have different attractions for the bonding electrons

o The atom with the higher EN will have a stronger attraction for the bonding electrons

• Most polar bonds are ‘polarized’ meaning that the electrons spend more time closer to the atom with the higher EN and less time near the atom of a lower EN

Polar Bonds cont’

• Molecules with polar bonds will have “Dipoles”

• Dipoles = a charge imbalance within a bond created by different attractions for the bonding electrons

d. Coordinate Covalent Bonds

Coordinate covalent bonds

bonds formed when only one element contributes electrons to the bond

Only in special cases:

d. Network Covalent Bonds

Network covalent bonds = these are very strong bonds formed within a network solid between atoms of the same element or molecule

• Special cases:

VI. Molecular Structures

• Molecular shapes depend upon the distribution of electrons number of bonds formed

• Shapes are 3-Dimensional

• http://www2.chemistry.msu.edu/faculty/reusch/virttxtjml/models.htm#start

a. Nonpolar Molecules• Nonpolar bondsNonpolar bonds are

formed only between atoms having the same EN

• Only diatomic elements have true nonpolar bonds

• All bonding electrons are shared equally between atoms of the same EN

• Ex. Diatomic molecules…

Nonpolar, cont’

• Even polar bonds can create nonpolar molecules…

• Nonpolar molecules are SYMMERTICAL!

• Electrons are evenly distributed throughout the molecule, making it nonpolar!

Symmetrical =

Nonpolar

•Ex.] CF4

b. Polar Molecules

• Polar Molecules have an asymmetric pull of electrons throughout the molecule

• Nonbonding electrons from lone pairs also create an asymmetric pull within the molecule

Asymmetric =

Polar

•Ex.] H2O

Polar or Nonpolar Molecule???

Examples:

a. CO2   

b. OF2    

c. CCl4    

d. CH2Cl2   

e. HCN

c. Molecular Shapes, in 3D!

• Atoms are 3-dimensional substances that create 3-D structures when bonding

• Both the bonds and the lone pair [nonbinding] electrons play a role in determining the shape of a molecule

Bonding/Molecular Shape Terms

• DomainDomain = placement of electrons around an atom

• Bonding DomainBonding Domain = includes all electrons participating in a bond; counts as one area of space

• Nonbonding DomainNonbonding Domain = space occupied by a lone pair of electrons [nonbonding]

Additional [secret] information:the VSEPR Theory

• VSEPR = Valence Shell Electron Pair Repulsion

– This theory explains why the electrons within the bonds and the nonbonding electrons move as far apart as possible, creating a structure in 3-dimensional space

• Nonbonding pairs sometimes have a greater effect than single bonds… let’s see!

B. Shapes and Bond Angles

http://intro.chem.okstate.edu/1314F97/Chapter9/VSEPR.html

Orhttp://en.wikipedia.org/wiki/Molecular_geometry

1. Linear

1 or 2 bonding domains 180o bond angle Symmetric if same elements, or distributed evenly Asymmetric if different atoms Examples: Diatomics, CO2, HCl

2. Trigonal Planar

3 bonding domains 120o bond angle Symmetric if all same elements Flat molecule! Examples: BF3, SO3

3. Trigonal Pyramidal

3 bonding domains, 1 nonbonding domain 107o bond angle Asymmetric due to lone pair electrons Examples: NH3, PCl3

4. Bent

2 bonding domains, 2 nonbonding domains 104.5o bond angle Asymmetric due to two lone pairs of electrons Examples: H2O, SCl2

5. Tetrahedron

4 bonding domains 109.5o bond angle Symmetric if all the same atoms bonded Asymmetric if different atoms Examples: CH4, CCl4

6. Trigonal Bipyramidal

5 bonding domains Expanded octet of 10 electrons 120o and 90o bond angle Symmetric if all the same atoms bonded Asymmetric if different atoms Examples: PF5

7. Octahedral

6 bonding domains Expanded octet of 12 electrons 90o bond angle Symmetric if all the same atoms bonded Asymmetric if different atoms Examples: SF6

VII. Ionic Bonds

• Ionic bonds = a bond formed due to the transfer of electrons between metals and nonmetals

• Attractions [bonds] occur between ions [charged atoms that have gained/lost electrons]

• CationsCations = positive ions; have _______e-– Metals form cations

• AnionsAnions = negative ions; have ______e-– Nonmetals form anions

Properties of Ionic Bonds

• High melting/boiling points

• Hard, but brittle crystals [solids]

• Dissolve in polar solvent

• Conduct electricity as liquid or in solution, but NOT as a solid

Properties cont.

• Ionic substances have high heats of vaporization

• Low vapor pressure; not very volatile

• Most dissolve in water to form (+) and (-) ions, or electrolytes

                                

A. Electronegativity Differences

Large differences in EN = Ionic BondsIonic Bonds When there is a larger difference in EN, the element with

the higher EN will most likely to “see” the bonding electrons more, or share them less

Ionic bonds have the greatest differences in EN! • reinforced by the fact that one of the elements will actually

TAKE the electrons instead of sharing them

Covalent Bonds and EN

•   Even though nonmetals have relatively low EN in general, they do do have slight differences

•  The only time there is no EN difference between atoms is for Diatomic elements

•  This means that the electrons in the bond(s) between the diatomic elements will be shared equally

Rankings of EN DifferencesSee Figure 6-11, page 107, and figure 6-14, page 109

0 0.3 1.7 >1.9

Diatomic Nonpolar Polar Ionic Elements Covalent Covalent Bonds

 

Polar bonds in Molecules• Arrows point to the element with the highest EN• Use lower case Greek letter delta to represent

partial charges: δ+ or δ-• Partially negative = more Electronegative atom!

b. Polyatomic Ions

• See Table E!!!• Have BOTH covalent and ionic properties• Covalent bonds hold the atoms together

within the ion• Overall, the structure has lost/gained

electrons to have a charge• Share electrons within, has brackets and

charges for the Lewis Structure

E. Resonance Structures

• Lewis dot structures with double-bond electrons that rotate from one pair to another

• Overall structure = hybrid of all resonance structures

Other Resonance structures

• NO3-1

• C6H6

• SO3-2

• CO3-2

Review:Bond Strengths

Network Covalent

Ionic

Covalent Triple Bond

Covalent Double Bond

Covalent Single Bond

More Stable Molecules = Stronger Bonds

Larger EN Differences = Stronger Bonds

Stronger BondsStronger Bonds

EqualEqual

More Stable MoleculesMore Stable Molecules

VIII. Intermolecular Forces

• INTRAmolecular forces: forces between atoms– Bonds = forces between the atoms

• INTERmolecular forces: forces between molecules– Four major variations

– Depends on the type of molecules or ions involved

1] Molecule-Ion Attractions:

Definition:

Invisible force of attraction holding polar molecules and ions together in a solution

Need: polar molecules as solvent and ionic compound [create (+) and (-) ions]

Molecule-ion forces…

• The strongest of all the intermolecular forces!– Positive ion attracted to partially negative end

of polar molecule– Negative ion attracted to partially positive end

of polar molecule– Orientation of polar molecules important!!!!– Ex.] Solution of NaCl(aq)

2] Dipole-Dipole Attractions

• Definition:

Partially positive and partially negative ends of polar molecules develop attractive forces

Need: polar molecules as liquid

Dipole-Dipole forces…• Occur within a sample of polar molecules

• Attraction occurs between partially positive ends of several of same polar molecules– Partially positive end of the molecule– near the

atom with lower EN • [bonding electrons pulled away from it]

– Partially negative end of molecule—near the atom with the highest EN

• [pulls bonding electrons towards it]

• Ex.] HCl, HBr, HI

3] Hydrogen Bonding

• Definition:

Special type of dipole-dipole forces occurring between polar molecules containing hydrogen and fluorine, oxygen, or nitrogen

Ex.] HF, H2O, and NH3

Hydrogen Bonding…• Stronger than Dipole-Dipole, but weaker than

actual bonds forming• Hydrogen [partially (+)] strongly attracted to

F, O, or N [partially (-)] end of molecule• F, O, and N have high EN, small radii, and strong

pull on bonding electrons• Responsible for:

– Abnormally high boiling point of water– Larger volume of water in liquid phase

4] Weak, London Dispersion, or Van der Waals Forces

• Definition:

Weak attractive forces present between nonpolar molecules

Need:

Nonpolar, symmetric molecules

Weak, LD, or VdW forces…

• Weakest attractive forces

• Created when nonpolar atoms/molecules have small, temporary dipoles formed via distribution of electrons

• Change as…– Distance between molecules increase, forces

decrease– Mass of molecules increase, forces increase

“Special effects” of weak/LD/VdW forces…

• Reason why diatomic elements of group 17 have increasing boiling points from top to bottom– Remember phases of group 17:

– gas, gas, liquid, solid, solid

• Cause hydrocarbons of fossil fuels to have increasing boiling points as their size and mass increase

– Methane is a gas, gasoline is a liquid, grease is a solid at the same temperatures

Strengths of IMF’s:Strongest to Weakest

1] Molecule-ion

2] Hydrogen bonding

3] Dipole-Dipole

4] London Dispersion/Van der Waals/Weak