Bonding and Naming

Bonding and Naming1. What is a chemical bond?

2. How do atoms bond with each other?

3. How does the type of bonding affect properties of compounds?

• How can all matter in the universe exist from only 92 elements?• Why can you dissolve salt in water but not melt it easily?• Why do the carbon atoms in graphite slide off my pencil lead while

diamonds are forever?• What is the role of carbon in the molecular diversity of life?• How can two toxic and violently reactive elements combine to

produce table salt (essential for life)?• Why is sodium lauryl sulfate found in my shampoo)?

Basics

• Bond– something that binds, attaches, or restrains

• Chemical Bond– the force that binds atoms to each other

• Why?– to Achieve Stability

(Noble Gas Configuration – s2p6)

Octet Rule

• Valence electrons– Electrons in the outermost energy level of an

atom• Rules

– Octet—Atoms bond to achieve 8 e- in outer energy level

– Duet—H bonds to achieve 2 e- in outer energy level

Lewis Structures (Electron Dot Diagrams)

• Using X to mean any element, draw Lewis Structures for elements in Groups 1,2,13,14,15,16,17,18

Electron Dot DiagramsGroup 1Group 2Group 13Group 14Group 15Group 16Group 17Group 18

x

Application of Octet Rule to Ionic Compounds

• Note: The element’s symbol represents the core electrons.

• Dots represent the valence electrons

Na ·

1s22s22p6 3s1

Application of Octet Rule

• Ionic Compounds– Metals go to cations

• Na Na+ + 1 e-

– Nonmetals go to anions• Cl + 1 e- Cl-

• Formation of NaCl

Na Cl[ ] + [ ]-

Practice

• Draw Lewis Structures for the following compounds:

• magnesium oxide

• calcium fluoride

• aluminum oxide

Answers

Mg O

2+ [ ]2-

Answers

Ca FF[ ] [ ]- -2+

Answers

Al

Al

O

O

O

3+

3+

2-

2-

2-

[ ][ ][ ][ ]

[ ]

Types of Chemical Bonds

• Three basic types of bonds– Ionic

• Electrostatic attraction between ions

– Covalent and covalent network• Shared electrons

– Metallic• Metal atoms bonded to

several other metal atoms

Types of Chemical Bonds

• Type of bond properties of compound

• We will begin with ionic bonding

Ionic Compounds

• also called salts• held together by attraction between (+) and

(-) charges in a lattice (electrostatic attraction)

• these charged particles are formed by donating (and receiving) electrons

• have neutral charge overall: the number of + charges = number of – charges

• strength of the bond is amplified through the crystal structure

Ionic Compounds

• repeating 3-D structure in crystal lattice = unit cell

Ionic Compounds

• Energy considerations– Forming a crystal lattice loss of a lot of PE

• ↑ stability

– “lattice energy” = E released when 1 mol of an ionic compound is formed from gaseous ions.

Making NaCl

http://www.youtube.com/watch?v=2mzDwgyk6QM

Making NaCl

Lattice Energy

Definition:

lattice energy = energy released when a crystal containing 1 mole of an ionic compound is formed from gaseous ions.

• Related to charge on the ion and distance between nuclei

higher charge stronger pull

smaller means closer together stronger pull

Lattice Energies of Some Ionic Compounds

Compound Lattice Energy

(kJ/mol)

Compound Lattice Energy

(kJ/mol)

KI -632 KF -808

KBr -671 AgCl -910

RbF -774 NaF -910

NaI -682 LiF -1030

NaBr -732 SrCl2 -2142

NaCl -769 MgO -3795

Lattice Energies of Some Ionic Compounds

• What happens as the ionic radius increases? (e.g. Check out the Na compounds or the F compounds)

• What happens as the amount of charge on a single ion increases? Both ions? (look at bottom right hand corner of table)

Lattice Energies of Some Ionic Compounds

• What happens as the ionic radius increases? Lattice energy decreases – ions are further away from each other.

• What happens as the amount of charge on a single ion increases? Both ions? Lattice energy increases, especially if both ions have multiple charges.

Lattice Energies of Some Ionic Compounds

• Predict whether the lattice energy of CsCl is larger or smaller than that of KCl.

• Predict whether the lattice energy of Na2O will be larger of smaller than that of MgO.

• What do you think will be the effect of lattice energy on melting point?

Lattice Energies of Some Ionic Compounds

• Predict whether the lattice energy of CsCl is larger or smaller than that of KCl.Smaller – Cs+ is larger than K+. (CsCl: -657 kJ/mol, KCl: -701 kJ/mol)

• Predict whether the lattice energy of Na2O will be larger of smaller than that of MgO.Smaller – Na+ is 1+, while Mg2+ has double the charge. (Na2O: -2481 kJ/mol)

• What do you think will be the effect of lattice energy on melting point?The higher the lattice energy, the more energy it takes to melt it, therefore melting point increases as lattice energy increases.

Characteristics

• solid at room temperature• hard, brittle• high melting point, boiling point• usually soluble in water• conduct electricity when melted or

dissolved in water (charges can move)• do not conduct electricity when solid

Dissolving NaCl in Water

Naming Chemical Compounds

Terms for describing how atoms bond

• Chemical Formula– Relative #’s of different elements using symbols and

subscripts– C2H4

• Empirical Formula– Chemical Formula Reduced to lowest whole #’s of

each element– C2H4→CH2

• Formula Unit– Lowest whole number ratios of Ionic Compounds– NaCl, MgBr2, etc.

Terms for describing how atoms bond

• Molecular Formula– Chemical formula of one molecule – e.g. C2H4

• Structural Formula of Molecule– Shows how atoms are bonded together in a

molecule H H

H H

C C

Naming Ions

• Monatomic– 1 atom + or –

• Polyatomic– Many atoms covalently bonded with an overall + or –

charge (especially anions)

Naming Cations

1. Monatomic Cations

– Use Atomic Symbol + Charge

e.g. Na+ = sodium Ion

(element’s name + ion)

Naming Cations

2. Polyatomic cations (Memorize these!)

NH4+ ammonium ion

Hg22+ mercury(I) ion

How Do We Know Charges?

• Apply Octet Rule (Groups 1, 2, 16, 17)

• Memorize Ion List

• Use Roman numberals to name elements with >1 charge– Copper(I) Ion =Cu+

– Copper(II) Ion=Cu2+

Naming Anions

• monatomic anion:

anions provide the second part of the

name of an ionic compound.

e.g. Cl- chloride ion

use suffix “ide” + “ion”

Can you name the rest of the monatomic anions?

Naming Polyatmic Anions

• Polyatomic (with oxygens)

– ite NO2- Nitrite SO3

2- Sulfite

– ate NO3- Nitrate SO4

2- Sulfate

• ClO-=Hypochlorite=least oxygens• ClO2

-=Chlorite

• ClO3-=Chlorate

• ClO4-=Perchlorate=most Oxygens

Naming Polyatmic Anions

• If the anion contains hydrogen, prefix name with “hydrogen”

–HCO3- = hydrogen carbonate

–HSO4- = hydrogen sulfate

–HS- = hydrogen sulfide

Let’s try all the phosphates that contain hydrogen:

Naming Binary Ionic Compounds

• cation + anion

Na+ + Cl-

• NaCl = sodium chloride

Naming Binary Ionic Compounds

• Steps

– Write Ions Pb4+ O2-

– Balance Charges (LCM) 4+ 2x2-– Empirical Formula/Formula Unit PbO2

– Name Lead (IV) Oxide

Practice

• Calcium Chloride• Iron (II) Oxide• Iron (III) Sulfide

• CaCl2• FeO

• Fe2S3

Ionic Compounds Containing Polyatomic Ions

• Few positive, many negative polyatomic ions• Use Atomic Symbols + Charge

– polyatomic cation• NH4

+ ammonium ion

• Hg22+ mercury(I) ion

– polyatomic anion• OH- hydroxide ion• CO3

2- carbonate ion

Note: Charge Refers to

whole groups of atoms,

not just th

e last o

ne

Polyatomic Ion Formulae

• Write ions– NH4

+ CO32-

• Balance Charges (LCM)– 2 x 1+ 2-

• Write Formula with lowest whole # ratios– (NH4)2CO3 ammonium carbonate

– Show >1 Polyatomic ion with parentheses

Polyatomic Ion Formulae

• Practice:

aluminum sulfate

magnesium hydroxide

copper(II) acetate

Polyatomic Ion Formulae

• Practice:

aluminum sulfate Al2(SO4)3

magnesium hydroxide Mg(OH)2

copper(II) acetate Cu(C2H3O2)2

Oxidation #’s

• Oxidation #=Charge of an Ion– K+ = oxidation # of +1– O2- = oxidation # of -2

• We use oxidation numbers to figure out the formulas of ionic compounds.

• The sum of oxidation numbers for the formulas of an ionic compounds must = 0.

Metallic Bonds

• held together by the attraction of free-floating valence electrons (-) for the positive ions in a lattice structure: “electron-sea”

Metallic Bonds1. What is a regular, repeating three-dimensional arrangement of

atoms called?

2. Do the separate electrons that are shown belong exclusively to a single atom? What word is used to describe such electrons?

3. Are the electrons shown the only ones actually present? Explain.

4. Why are the central atoms shown as positively charged?

5. How does the number of separate electrons shown for the group 1A metal atoms compare to the number of atoms? Explain why in terms of valence electrons.

Metallic Bonds - KEY1. What is a regular, repeating three-dimensional arrangement of atoms

called? Crystal lattice

2. Do the separate electrons that are shown belong exclusively to a single atom? What word is used to describe such electrons? no, they are delocalized

3. Are the electrons shown the only ones actually present? Explain.No, they are valence electrons from the metal atoms.

4. Why are the central atoms shown as positively charged?The delocalized (valence) electrons come from the neutral atoms, thus leaving the atoms with a positive charge – i.e. cations.

5. How does the number of separate electrons shown for the group 1A metal atoms compare to the number of atoms?Explain why in terms of valence electrons. They are equal – have only one valence electron.

Metallic Bonds6. How does the number of separate electrons shown for the group 2A

metal atoms compare to the number of atoms?

7. What holds the metal atoms together in such an arrangement?

8. What term is used to describe this model of metallic bonding?

9. How well do metals tend to conduct electricity? How does the model of metallic bonding account for that property?

10.Do metals tend to be brittle, or are they malleable and ductile? How does the model of metallic bonding account for that property?

Metallic Bonds - KEY6. How does the number of separate electrons shown for the group 2A metal atoms

compare to the number of atoms? twice as many electrons

7. What holds the metal atoms together in such an arrangement?Delocalized electrons are simultaneously attracted to > 1 metal cation.

8. What term is used to describe this model of metallic bonding? electron sea

9. How well do metals tend to conduct electricity? How does the model of metallic bonding account for that property? Metals tend to conduct electricity well. The delocalized electrons are not held strongly by individual atoms and are thus able to move easily throughout the metal.

10. Do metals tend to be brittle, or are they malleable and ductile? How does the model of metallic bonding account for that property? malleable and ductile. The delocalized electrons are able to move around the positive metal core atoms and keep the crystal from breaking.

Characteristics of Metals

• Characteristics of Metals– high m.p. and b.p.– Good conductors of electricity and heat

conductors in the solid state– malleable, ductile– shiny, reflective, usually gray (or grey)

Characteristics of Metals

As a metal is struck by a hammer, the atoms slide through the electron sea to new positions while continuing to maintain their connections to each other.

Warmup – Naming Ionic Compounds – binary, monatomic

1.Name the following ionic compounds, and write the pairs of ions that make up these compounds:

a) KCl

b) SrO

c) PbF2

Warmup – Naming Ionic Compounds – binary, monatomic

2.Write the chemical formula for the following ionic compounds, and write the ions that make up these compounds:

a) lithium bromide

b) barium sulfide

c) chromium (III) oxide

Warmup – Naming Ionic Compounds, polyatomic ions

3. Write the chemical formula for the following ionic compounds, and write the ions that make up these compounds:

a) cesium oxalate

b) calcium hydroxide

c) potassium sulfate

d) ammonium hydrogen phosphate

Warmup – Naming Ionic Compounds

4. Write the chemical formula for the compounds made up of the following ions, then name the compounds:

a) Cs+, O2-

b) Pb2+, Br-

c) Fe3+, F-

d) Na+, CO3-

Ionic BondingElectron Dot Diagrams

5. How many of each ion will you need to form each of the following ionic compounds?Draw electron dot diagrams to demonstrate the ionic bonding of each compound.

a) lithium and bromineb) potassium and sulfurc) aluminum and chlorined) gallium and oxygene) bismuth (V) and sulfur

Alloys

• Mixture of elements that has metallic properties (solid solutions)

• Substitutional alloys—atoms of similar sizes– Brass (Cu +Zn)– Bronze (Cu + Sn + Pb)– Pewter (Sn + Sb + Pb)

• Interstitial—much smaller atoms fill spaces between larger atoms– Carbon Steel

Alloys

• It is more difficult for layers of atoms to move over each other, especially in interstitial alloys.

http://www.frankswebspace.org.uk/ScienceAndMaths/chemistry/alloys.htm

Review

• Ionic compounds – crystal lattice of oppositely-charged ions, held together by electrostatic charges

• Metals – lattice of positively-charged ions in a sea of electrons

Covalent Bonding

• Held together by shared electrons (covalent bond)

• Both electrons spend time around each nucleus but spend most of their time in the middle

Diatomic Elements

• Memorize: N O F Cl Br I H

• all exist diatomically in nature (more stable): N2 O2 F2 Cl2 Br2 I2 H2

• 7- rule: Go to Element #7, travel across to Group 7A, then down in the shape of a 7.

Another way to remember them:• Hairogens: H2 (H), N2 and O2 (air), F2, Cl2,

Br2, I2 (halogens)

Diatomic Molecules

• 7 Elements bond with themselves (increased stability)– H→H2

– N→N2

– O→O2

– F→F2

– Cl→Cl2– Br→Br2

– I→I2

Covalent Bonding

• Characteristics of covalently-Bonded compounds (molecules)

– relatively low m.p. and b.p.– do not conduct electricity under any

circumstances– generally not soluble in water, but soluble in

alcohol

Covalent Network Solids

• form covalent bonds in all directions continuous network of strong covalent bonds no individual molecules

• extremely hard, very high m.p. and b.p.• nonvolatile, insoluble in all solvents• brittle, nonconductors of heat and

electricity• diamond (C), quartz (SiO2)

Compare structures of networks/lattices:

a) Covalent network (quartz, SiO2)b) Salt (NaCl)c) Metal (Cu)

Chemical Bonds• Metallic ‘bond’

– Metals and Alloys– ‘Sea’ of electrons

• Ionic bond– Ionic Compounds or Salts– Metal + Non-metal: NaCl, MgSO4

– Electrons exchanged between atoms

• Covalent bond– Molecules and covalent network

solids– Non-metals: H2O, CH4

– Electrons shared among atoms

Formation of a Covalent Bond

• Covalent Bond– When orbitals from two atoms overlap– 2 electrons of opposite spin in the overlap– As the amount of overlap ↑, the energy of the

interaction ↓• When minimum energy is reached, bonding

distance occurs• Attraction and repulsion of electrons and nuclei are

exactly balanced

– At some distance, nuclei repel, increasing energy again

Formation of a Covalent Bond

Energy Considerations

• What does “stable” mean? Changes that lower potential energy are favored.

• In covalent bonds:• Shared electrons loss of PE

­ stability• Bond energy = energy required to break a

chemical bond and form neutral atoms

Relationship between bond length and bond energy in

molecules • Bond length = average distance between

two bonded atoms (distance of minimum potential energy)

• As Ebond ↑ ­, Lengthbond ¯, because the closer the atoms are, the more attraction between nuclei and electron clouds. ­

­­­ ­\ harder to separate.

Naming Molecular Compounds

• 2 systems – prefixes, vs. oxidation numbers (Honors/AP)

• See p. 248 in textbook for list of prefixes to memorize

Naming Molecular CompoundsA. Prefixes, roots, suffixes1. Begin with element with lowest electronegativity.

Nitrogen

2. Add appropriate prefix (unless it is mono-). Dinitrogen

3. End second element with "ide" (as for ionic compounds…). Oxide

4. Use appropriate prefix for second element. tetroxide dinitrogen tetroxide

N2O4

Naming Molecular CompoundsPractice:

CCl4

CO

CO2

As2S3

P2O5

P4O10

Naming Molecular CompoundsPractice:

CCl4 carbon tetrachloride

CO carbon monoxide

CO2 carbon dioxide

As2S3 diarsenic trisulfide

P2O5 diphosphorus pentoxide

P4O10 tetraphosphorus decoxide

Oxidation Numbers (p. 2 of Honors Supplement)

Oxidation number = the charge on an ione.g. K+ has an oxidation number of 1+ O2- has an oxidation number of 2-

We use oxidation numbers to figure out the formulas of ionic compounds.

The sum of oxidation numbers for the formulas of an ionic compounds must = 0.

Oxidation Numbers

We can also use oxidation numbers for molecular compounds, by pretending the atoms are ions = “apparent charge”.

Oxidation Numbers

The oxidation number of:

• an element in the uncombined state is 0.

• a monatomic ion equals the charge on the ion.

• hydrogen is generally +1; in hydrides, -1.

• oxygen is generally -2; in peroxides, -1.

Oxidation NumbersThe oxidation number of:

• the more electronegative element in a binary covalent compound is negative, while that of the other element is positive.

• elements other than oxygen and hydrogen in a neutral compound is such that the sum of the oxidation numbers for all atoms in the compound is 0.

• elements other than oxygen and hydrogen in a polyatomic ion is such that the sum of the oxidation numbers for all atoms in the ion equals the charge on the ion.

Use these rules to assign oxidation numbers to each element in each of the given formulas

e.g. H2O H: 2 x +1, O: -2

N2 0

Assign oxidation numbers to each element:

1. Cl22. Cl-

3. Na

4. Na+

5. KCl

6. H2S

7. CaO

8. H2SO4

Assign oxidation numbers to each element:

9. NO3-

10. Cr2O72-

11. NH4Cl

12. NH3

13. NO2

14.CaH2

(calcium hydride)

15.Na2O2

(sodium peroxide)

Assign oxidation numbers to each element: KEY

1. Cl2 0

2. Cl- 1-

3. Na 0

4. Na+ 1+

5. KCl K: 1+, Cl: 1-

6. H2S H: 2 x 1+, S: 2-

7. CaO Ca: 2+, O: 2-

8. H2SO4 H: 2 x 1+, S: 6+, O: 4 x 2-

Assign oxidation numbers to each element: KEY

9. NO3- N: 5+, O: 3 x 2-

10. Cr2O72- Cr: 2 x 6+, O: 7 x 2-

11. NH4Cl N: 5+, H: 4 x 1-, Cl: 1-

12. NH3 N: 3+, H: 3 x 1-

13. NO2 N: 4+, O: 2 x 2-

14.CaH2

(calcium hydride)

Ca: 2+, H: 2 x 1-

15.Na2O2

(sodium peroxide)

Na: 2 x 1+, O: 2 x 1-

Oxidation Numbers - warmup

• Write the oxidation numbers for all elements in the following compounds:

P2O5

NO2-

NO3-

HNO3

K2CrO4

K2Cr2O7

Oxidation Numbers - warmup

• Write the oxidation numbers for all elements in the following compounds:

P2O5 P: 2 x 5+, O: 5 x 2-

NO2- N: 3+, O: 2 x 2-

NO3- N: 5+, O: 3 x 2-

HNO3 H: 1+, N: 5+, 3 x 2-

K2CrO4 K: 2 x 1+, Cr: 6+, O: 4 x 2-

K2Cr2O7 K: 2 x 1+, Cr: 6+, O: 7 x 2-

Using Oxidation Numbers to Name Molecular Compounds

1. Name the following molecular compounds using oxidation numbers:

a) CCl4 carbon (IV) chloride

b) CO

c) CO2

d) P2O5

e) PCl5f) SO2

Using Oxidation Numbers to Name Molecular Compounds

1. Name the following molecular compounds using oxidation numbers:

a) CCl4 carbon (IV) chloride

b) CO carbon (II) oxide

c) CO2 carbon (IV) oxide

d) P2O5 phosphorus (V) oxide

e) PCl5 phosphorus (V) chloride

f) SO2 sulfur (IV) oxide

Using Oxidation Numbers to Name Molecular Compounds

2. Write formulas for the following molecular compounds:

a) carbon(IV) iodide Cl4

b) sulfur(VI) oxide

c) nitrogen(IV) oxide

d) arsenic(III) sulfide

e) phosphorus (III) fluoride

Using Oxidation Numbers to Name Molecular Compounds

2. Write formulas for the following molecular compounds:

a) carbon(IV) iodide Cl4

b) sulfur(VI) oxide SO3

c) nitrogen(IV) oxide NO2

d) arsenic(III) sulfide As2S3

e) phosphorus (III) fluoride PF3

Compounds That Become Acids When Dissolved in Water

• General Formula: HX

H+ X-

monatomic

or polyatomic

anion

Compounds That Become Acids When Dissolved in Water

Three Rules:

1. When X ends in “ide”

(e.g. chloride, cyanide)

“hydro_______ ic acid”

e.g. hydrochloric acid,

hydrocyanic acid

Compounds That Become Acids When Dissolved in Water

Three Rules:

2. When X ends in “ite”

(e.g. chlorite, sulfite)

“______ous acid”

e.g. chlorous acid

sulfurous acid

Compounds That Become Acids When Dissolved in Water

Three Rules:

3. When X ends in “ate”

(e.g. chlorate, sulfate)

“______ ic acid”

e.g. chloric acid

sulfuric acid

Compounds That Become Acids When Dissolved in Water

Your turn:

HBr

HNO2

HNO3

Compounds That Become Acids When Dissolved in Water

Your turn:

HBr hydrobromic acid

HNO2 nitrous acid

HNO3 nitric acid

Warmup

Write formulas for the following molecular compounds:• sulfur trioxide• phosphorus pentachloride• nitrogen dioxide• tetraphosphorus decoxide• oxygen difluoride

Warmup• Write Lewis structures (dot diagrams) for the

following elements:• carbon• hydrogen• fluorine• sulfur• nitrogen• oxygen• phosphorus• bromine

Molecules and Lewis Structures

Lewis structures show

• All atoms in the molecule

• How atoms are connected (# of bonds)

• Any unshared electron pairs (lone pairs)

Making Lewis Structures

1. Find total # of valence e-’s for ALL atoms (for ions, consider the charge).

2. Write atom symbols, beginning with the central atom:

a) Carbon in center – CH4

b) Most electropositive in center – SO42-

c) Non metal (other than H or O) in center – H2PO4

-

d) Hydrogen and oxygen are usually on the outside

Making Lewis Structures

3. Add valence e’-s to each atom, beginning with the bonding electron for each atom added to the central atom.

4. Shared e- pairs represent bonds and are counted in the valences of both elements.

5. Adjust so that each element has 8 valence electrons (2 for H), note exceptions…

Lewis Structure Practice(Electron Dot Diagrams)

Molecules

CH3Br

BrI

H2S

PH3

IonsClO-

SO42-

H2PO4-

NH4+

HydrocarbonsC2H6

C2H4

C2H2

C6H6

Warmup – naming ionic compounds

How many of each ion is in each of the following compounds?

a) AlBr3

b) PbCl4

c) RbNO3

d) MgSO4

e) K3PO4

Warmup – naming ionic compounds

How many of each ion is in each of the following compounds?

a) AlBr3 1 Al3+, 3 Br-

b) PbCl4 1 Pb4+, 4 Cl-

c) RbNO3 1 Rb+, 1 NO3-

d) MgSO4 1 Mg2+, 1 SO42-

e) K3PO4 3 K+, 1 PO43-

Warmup – naming ionic compounds

• Write the formula for each of the following compounds:

a) ammonium phosphate

b) cesium oxide

c) copper(I) fluoride

d) silver nitride

e) beryllium nitrate

Warmup – naming ionic compounds

Write the formula for each of the following compounds:

a) ammonium phosphate (NH4)3PO4

b) cesium oxide Cs2O

c) copper(I) fluoride CuF

d) silver nitride Ag3N

e) beryllium nitrate Be(NO3)2

Exceptions to the Octet Rule

Central atom has < 8 valence e-’s

BeF2

BF3

Central atom has > 8 e-’s

PF5

SF6

XeF4

Answers

C

HHH

Br4 e-

7 e-

1 e-

1 e-

1 e-+14 e-

CH3Br

Answers

7 e-

6 e-

1 e-

ClO+

14 e-

ClO-

Answers

7 e-

7 e-Br

I

+

14 e-

BrI

Answers

6 e-6 e-

6 e-

6 e-

6 e-

+ 2 e-

SO42-

Warmup – Lewis Structures

Draw Lewis Structures for the following molecular compounds:• carbon tetrachloride• sulfite ion• ammonia (NH3)

• water• chlorine (Cl2)

• oxygen (O2)

• nitrogen (N2)

Molecular GeometryFrom 2D Lewis Structures to 3D: Valence Shell Electron Pair Repulsion (VSEPR Theory)

The influence of Unshared Pairs on Geometry

When describing the shape of a molecule, consider the arrangement of the atoms,

which is influenced by the unshared (lone) pairs

Molecular GeometryTerminology

• Electron domain = region where electron pairs reside

= #bonded pairs + #lone pairs

• Bonding domain = bonded pair

• Nonbonding domain = lone pair

• Annotate the diagrams in your notes with these terms

2-3 electron domains

VSEPR Theory – Determining Molecular Shapes

4 electron domains

5 electron domains

5 electron domains (Honors)

6 electron domains

Back to Activity

• Observe the models, then draw and name the shapes in table B

Note: double and triple bonds = 1 lone pair

• Next, you will build models of the compounds in table A, then draw and name the structures

• We will discuss polarity later this week.

Hybridization (Honors)

• Atomic orbitals are mixed new, identical hybrid orbitals

• helps to explain VSEPR

• # of hybrid orbitals = # atomic orbitals mixed, including lone pairs

Hybridization

Hybrid Orbital Summary

Electron-domain geometry must be known before hybridization is assigned.

• To assign hybridization:• Draw a Lewis structure.• Assign the electron-domain geometry using VSEPR

theory.• Specify the hybridization required to accommodate the

electron pairs based on their geometric arrangement.

• Name the geometry by the positions of the atoms.

Warmup

1. What is meant by the term “polar”?

Use an example that makes sense to you.

2. Using the PT of theElectronegativities, calculate the difference in electronegativities between

a) Na and Cl

b) Br and Cl

c) B and Cl

3. Identify the bonds in Q2 as ionic, polar covalent or

nonpolar covalent.

Warmup

1. What is meant by the term “polar”?

Use an example that makes sense to you.

2. Using the PT of theElectronegativities, calculate the difference in electronegativities between

a) Na and Cl (2.1) ionic

b) Br and Cl (0.2) NP covalent

c) B and Cl (1.0) P covalent

3. Identify the bonds in Q2 as ionic, polar covalent or

nonpolar covalent.

Ionic, polar covalent or nonpolar covalent?

• The nature of the bond between any two atoms is determined by the difference in their electronegativities (see chart on ho)

• The greater the difference, the more ionic the bond (on a continuum)

ionic covalent

polar (P) NP

0.41.74.0 0

Ionic range: 4.0 - 1.8

Cl – Na: 3.0 - 0.9 = 2.1 (ionic)

Polar covalent range: 1.7 - 0.5

Cl – Al: 3.0 - 1.5 = 1.5 (polar covalent)

Nonpolar covalent range: 0.4 - 0.0

Cl – Br: 3.0 - 2.8 = 0.2 (nonpolar covalent)

ionic covalent

polar (P) NP

0.41.74.0 0

Cl - Na Cl - Al Cl - Br

• ionic bonds - the less electronegative atom donates 1 or more e-’s to the more electronegative atom

• covalent bonds - e-’s from both atoms are sharedo Polar covalent bonds - e-’s are shared, but

are not shared equally between two atomso nonpolar covalent bonds- e-’s are shared

equally between two atoms

Ionic, polar covalent or nonpolar covalent?

Polar vs. Non-polar:a matter of symmetry

Determining Molecule Polarity

1. Draw electron dot diagram (Lewis structure) of the molecule, let lines represent bonds

2. Compare the electronegativities of each of the bonded atom pairs

3. Determine whether each bond is polar (P) or nonpolar (NP)

4. Draw an arrow parallel to each bond directed towards the more electronegative atom

5. Is the molecule symmetrical? (If you cut it through the xy, xz and yz planes, would it split into mirror images that are identical?)

• Yes - molecule is NP • No - molecule is P

Determining Molecule Polarity

Examples

HCl

CCl4CH3Cl

NH3

BF3

H2O

Back to Activity

• Last column of Data Table B

• Last two columns of Data Table A

Sigma Bonds and Pi Bonds (Honors) canceled this year

• Sigma (s) bonds: electron density lies on the axis between the nuclei.

• All single bonds are s bonds.

Sigma Bonds and Pi Bonds(Honors) canceled this year

• What about overlap in multiple bonds?

• Pi (p) bonds: electron density lies above and

below the plane of the nuclei.

• A double bond consists of one s bond and

one p bond.

• A triple bond has one s bond and two p

bonds (above and below the plane of the nuclei;

in front of and behind the plane of the nuclei).

Sigma Bonds and Pi Bonds canceled this year

Formation of a Double Bond canceled this year

Formation of a Double Bond canceled this year

Formation of a Double Bond canceled this year

Triple Bondcanceled this year

Warm Up

For each of the following three compounds:

CBr4, CH2O, and CCl2F2

a)Draw the Lewis structures

b)Name the shape

c)Calculate the polarity of each bond

d)Predict the polarity of each molecule

e) (Honors only) What is the hybridization of the central atom?

Intermolecular forces (van der Waals forces)

• Dispersion (London) forces

• Dipole force

• Hydrogen bonding

1. Dispersion (London) forces• named after Fritz London, 1900-1954)

a. weakest intermolecular force

b. results from the constant motion of electrons uneven distribution of electrons at any particular moment:

“temporary dipole” which may dipole in nearby molecule.

c. acts on all molecules all the time

d. only intermolecular force acting among noble gas atoms and nonpolar molecules

1. Dispersion (London) forces

e. ­ with ­ number of electrons:

note m.p., b.p.

e.g. halogens

F2, Cl2 gases at room T

Br2 liquid at room T

(more e-’s than F2 and Cl2)

I2 solid at room T (most e-’s)

1. Dispersion (London) Forces

2. Dipole Force (polar molecules)

a. the attraction between two polar molecules:

(-) end of one polar molecule attracts the (+) end of another polar molecule

b. more polar stronger dipole force

c. closer together stronger dipole force

2. Dipole Force (polar molecules)

3. Hydrogen bonding

a. always involves Husually involves O, F or N (small, high electronegativity)

b. strongest intermolecular forceHow strong? 5% of the strength of a covalent bond

c. higher b.p. and higher viscositye.g. H2O

3. Hydrogen bonding

Comparing Bond Types Properties Ionic: salts Metallic Covalent:

moleculesCovalent network

Description of bond

General appearance

Malleable vs. brittle

Conduct electricity?

Under what circumstances

?

m.p., b.p.

State of matter at

room temp?

Comparing Bond Types Properties Ionic: salts Metallic Covalent:

moleculesCovalent network

Description of bond

Electrostatic attraction

Sea of electrons

Shared electrons

Shared electrons in

networkGeneral

appearancecrystalline Shiny, prob.

grayRange of

colorsRange of

colors

Malleable vs. brittle

brittle malleable brittle Prob. brittle

Conduct electricity?

yes yes no no

Under what circumstances

?

dissolved in H2O or molten

all N/A N/A

m.p., b.p. High High Low Very high

State of matter at

room temp?

Always solid Mainly solid Gas/solid solid

Comparing ALL bond types:

Which is stronger?covalent network > metallic > ionic > covalent (molecules) > H bond > dipole > dispersion

e.g. Compare melting points:SiO2 > Fe > NaCl > C12H22O11 > H2O > HCl > H2

sand > iron > salt > sugar > ice > hydrogen chloride > hydrogen gas