CHEM 200/202Professor Jing GuOffice: EIS-210

All emails are to be sent to:chem200@mail.sdsu.edu

My office hours will be held in online on Monday 9-11 AM or by appointment.

ANNOUNCEMENTS• Chapter 6 homework due Monday March 16.

• Class (https://SDSU.zoom.us/j/281477051)

• Go to the help room GMCS-212 (https://SDSU.zoom.us/s/239381878)

• Go to the Supplemental Instruction sessions(http://bit.ly/virtualCHEM200)

LECTURE OBJECTIVES• Chapter 7: Chemical Bonding and Molecular Geometry

• Chapter 7.1 Ionic Bonding

• Explain the formation of cations, anions, and ionic compounds.

• Predict the charge of common metallic and nonmetallic elements, and write their electron configurations.

• Chapter 7.2 Covalent Bonding

• Describe the formation of covalent bonds.

• Define electronegativity and assess the polarity of covalent bonds.

• Chapter 7.3

• Write Lewis symbols for neutral atoms and ions.

• Draw Lewis structures depicting the bonding in simple molecules.

TYPES OF CHEMICAL BONDS• Ionic Bonding:

• Transfer of valence electrons from metal to non-metals

• Covalent Bonding:

• Sharing of valence electrons between atoms

• Metallic Bonding:

• Sharing of valence electrons between metals - not covered in this course

MODELS OF CHEMICAL BONDING

Ionic Bonding

Covalent Bonding

Metallic Bonding

IONIC VERSUS COVALENT

• Ionic bonds involve the transfer of valence electrons.

• Covalent bonds involve the sharing of valence electrons.

Ionic Bonding

Covalent Bonding

GENERAL PROPERTIES OF IONIC COMPOUNDS

• Compounds with ionic bonds share similar properties due to the way that ions interact.

• Physical behavior : hard, rigid, and brittle (e.g. NaCl)

• Electrical Conductivity: typically only conduct electricity well when dissolved or melted.

• Melting & Boiling points: very high melting and boiling points; vaporize ionic compounds form ion pairs.

IONIC BONDING MODEL• Involves the transfer of electrons from a metal

to a non-metal.

• This transfer results in the formation of ions of opposing charge, which bond each other in a solid.

• Na=[Ne]3s1 → Cl = [Ne]3s23p5

• Na+=[Ne] and Cl- = [Ne]3s23p6 = [Ar]

• The ions formed have a filled outer shell (valence shell) of eight electrons (octet rule)

Ionic Bonding

IONIC COMPOUNDS• Predict the formulas for the ionic compounds made from

the pairs below:

• Ca and F

• Al and O

• Sn and Cl

• Sr and N

GENERAL PROPERTIES OF COVALENT COMPOUNDS

• Compounds with covalent bonds share similar properties due to the way that their atoms interact.

• Physical behavior : varied properties due to strong bonds between atoms, but weak intermolecular attraction. Network covalent compounds are very very strong due to the extensive bonding.

• Electrical Conductivity: poor electrical conductors in all states.

• Melting & Boiling points: low melting and boiling points, except network covalent compounds.

diamond

COVALENT BOND FORCES

Covalent bonds are a balance of attractive forces between the electrons and nuclei and repulsive forces between electrons and between nuclei.

H2 COVALENT BOND

Figure 7.4 The potential energy of two separate hydrogen atoms (right) decreases as they approach each other,and the single electrons on each atom are shared to form a covalent bond. The bond length is the internucleardistance at which the lowest potential energy is achieved.

It is essential to remember that energy must be added to break chemical bonds (an endothermic process), whereasforming chemical bonds releases energy (an exothermic process). In the case of H2, the covalent bond is very strong;a large amount of energy, 436 kJ, must be added to break the bonds in one mole of hydrogen molecules and cause theatoms to separate:

H2(H) ⟶ 2H(H) Δ) = 436 kJ

Conversely, the same amount of energy is released when one mole of H2 molecules forms from two moles of H atoms:

2H(H) ⟶ H2(H) Δ) = −436 kJ

340 Chapter 7 | Chemical Bonding and Molecular Geometry

This OpenStax book is available for free at http://cnx.org/content/col11760/1.9

The distance between atoms in a covalent bond is determined by the minimization of energy.

As the nuclei get closer the attractive energy favors the bond.

Getting too close increases the repulsion forces.

COVALENT VS POLAR COVALENT BONDS

• In covalent bonds the bonding electrons are shared between the atoms. But what happens if the electrons aren’t shared evenly?

• Polar covalent bonds arise when one atom pulls the shared electrons closer to itself than the other bound atom can.

• The atom which draws the electrons towards itself gains a partial negative charge (δ-), while the other atom gains a partial positive charge (δ+).

Polarity arrow indicates the

direction in which electrons are being

pulled.

ELECTRONEGATIVITY (EN)• The ability of an atom within a molecule to attract valence electrons

towards itself.

• High EN - more electronegative

• Low EN - more electropositive

• Periodic trends in Electronegativity:

• EN tends to decrease down a group.

• EN tends to increase across a period.

• EN tends to do the opposite of the ionization energy

PAULING ELECTRONEGATIVITY VALUES

BOND CLASSIFICATION• Bonds can be classified by

their electronegativity (EN) differences.

• Polar covalent bonds share characteristics of both polar and ionic bonds.

• The characteristics of the bond will generally follow the EN differences - greater difference means more ionic character.

QUESTIONWhich if the most polar covalent bond?

Answer:A: N－HB: O－HC: H－FD: C－FE: All of these are non-polar

PAULING ELECTRONEGATIVITY VALUES

PROBLEM• Are the bonds in each of the following

substances ionic, non-polar covalent, or polar covalent? Arrange the substances with polar covalent bonds in order of increasing bond polarity.

• S8

• RbCl

• PF3

• SCl2

• F2

• SF2

• Rank the members of each set of compounds in order of increasing ionic character of their bonds. Use polar arrows to indicate the bond polarity of each.

• HBr, HCl, HI

• H2O, CH4, HF

• SCl2, PCl3. SiCl4

LEWIS ELECTRON-DOT SYMBOLS

• The Lewis dot system is a means of representing the valence electrons and the bonds formed by atoms.

• For main group elements, the group number identifies the number of valence electrons.

• Each dot represents a valence electron; dots are placed around the “nucleus” on all four sides, and only become paired when more than four electrons are present.

Nitrogen5 valence electrons N•••••

N• •• •

•N•• •••

N••• ••

QUESTIONWhich is the correct Lewis electron-dot

symbol for carbon?

C••• •

•C• ••C• ••• •

C• •••

Answer:

- A- B- C

- D

LEWIS ELECTRON-DOT SYMBOLS

Dots represent the valence electrons around the symbol of the element. No more than one pair of electrons per

side, pairing only occurs on the 5th electron.

FORMATION OF LI+ AND F-

Li 1s22s1 + F 1s22s22p5 Li+ 1s2 + F- 1s22s22p6

Li

1s 2s 2p

F

1s 2s 2p

+

Li+

1s 2s 2p

F- 1s 2s 2p

+

.+ F : ::Li . Li+ + F - ::

::

Electron Configurations

Orbital Diagrams

Lewis electron-dot symbols

LEWIS THEORY OF CHEMICAL BONDING• Lewis structures - represent valence electrons around the

atom; only pair electrons when more than 4 valance electrons.

• A covalent bond in a molecule forms when two atoms in the molecule share one, or more, electron pairs.

• Octet Rule: when main-group elements form bonds, they lose, gain or share valence electrons so as to “have” a total of eight.

• Lewis structures depict the connectivity of the molecule.

• Classifies valence electrons in the molecule:

• Bonding electrons - shared Lone pairs - unshared

CO O••••

••••

CONVERTING FORMULAS INTO LEWIS STRUCTURES

1. Start with molecular formula.

2. Place atoms - typically atom with least electronegative atom in center.

3. Valence electrons - sum the total # of valence e- for all atoms

4. Draw single bonds between atoms - subtract 2e- for each single bond, 4e- for double bonds...

5. Ensure each atom has 8e- (2e- for H) - use lone pairs.

LEWIS STRUCTURES

• Hydrogen atoms are terminal.

• Halogen atoms are terminal. Exceptions: the heavier halides (Cl, Br, I) are central atoms only if they bind to more than one O or F atom.

• The less electronegative atoms will be the central atom.

• More than one central atom: the less electronegative atom will be surrounded by more atoms than the more electronegative atom.

• Structures with the maximum number of bonds are preferred.

The following rules allow us to predict the most stable Lewis structure: