Chemical Bonding and Nomenclature Adapted from Paul Surko

Chemical Bonding and Nomenclature

Adapted from Paul Surko

Bonding, the way atoms are attracted to each other to form compounds, determines nearly all of the chemical properties we see. The number “8” is very important to chemical bonding.

What is Bonding????????

What are Compounds?Compounds are a

combination of atoms bonded together. Bonding determines the chemical

properties of the compound.

Ionic Bonding-Great with 8All atoms want 8 valence electrons. Metals give up electrons to form positive ions (cations) and non-metal atoms will receive or take additional electrons to become negative ions (anions). IONS are charged particles.

N becomes N-3

Al becomes Al+3Cl becomes Cl-

O becomes O-2

Mg becomes Mg+2

Na becomes Na+

The positive and negative ions are attracted to each other electrostatically.

Properties of Ionic Compounds

Made of cations and anions Exist in crystalline structures (solids) at STP High MPs and BPs Conduct electricity in aqueous and molten

states

Opposites Attract!

Putting Ions TogetherNa+ + Cl- = NaCl

Ca+2 + O-2= CaO Na+ + O-2 = Na2O

Al+3 + S-2 = Al2S

3 Ca+2 + N-3 = Ca3N

2

Ca+2 + Cl- = CaCl2

You try these!

Mg+2 + F- =

NH4

+ + PO4

-3 =

K+ + Cl- =

Al+3 + I- =

Sr+2 + P-3 =

Li+ + Br- =

Sr3P

2

AlI3

MgF2

(NH4)

3PO

4

KCl

LiBr

Not NH43

PO4

NomenclatureNaming of Ionic Compounds with TMs

Binary Compounds have two types of atoms (not diatomic which has only two atoms).

Metals (Groups I, II, and III) and Non-Metals

Metal _________ + Non-Metal _________ideSodium Chlorine

Sodium Chloride NaCl

Metals (Transition Metals) and Non-Metals

Metal ______ +Roman Numeral (__) + Non-Metal ________ide Iron III Bromine

Iron (III) Bromide FeBr3

Compare with Iron (II) Bromide FeBr2

Metals (Transition Metals) and Non-MetalsOlder System

Ferrous Bromine

Ferrous Bromide FeBr2

Compare with Ferric Bromide FeBr3

Metal (Latin) _______ + ous or ic + Non-Metal ________ide

Nomenclature--Naming of Ionic Compounds with TM

Let’s Practice!Name the following.

CaF2

K2S

CoI2

SnF2

SnF4

OF2

CuI2

CuI

SO2

SrS

LiBr

Strontium SulfideLithium Bromide

Copper (I) Iodide or Cuprous Iodide

Sulfur dioxide

Copper (II) Iodide or Cupric Iodide

Oxygen diflourideTin (IV) Flouride or Stannic Flouride

Tin (II) Flouride or Stannous Flouride

Cobalt (II) Iodide or Cobaltous IodidePotassium Sulfide

Calcium Flouride

Polyatomic Ions(partial list from page 195 (193 2nd edition))

Ammonium……………... Nitrate…………………… Permanganate…………. . Chlorate………………… Hydroxide………………. Cyanide…………………. Sulfate…………………... Carbonate………………. Chromate……………….. Acetate………………….. Phosphate……………….

NH4+

NO3-

MnO4-

ClO3-

OH-

CN-

SO4 2 -

CO32-

CrO42-

C2H3O2-

PO43-

Lets Practice!

Na2CO3

KMnO4

NaOH

CuSO4

PbCrO4

NH3ammonia

Copper (II) sulfate or Cupric sulfate

Lead (II) chromate or Plubous chromate

Sodium hydroxide

Potassium permanganate

Sodium carbonate

The Covalent BondAtoms can form molecules by sharing

electrons in the covalent bond. This is done only among non-metal atoms.

Properties of Covalent Compounds

Generally exist as liquids and gases at STP NO crystalline structure Low MPs and BPs Do NOT conduct electricity

Dot Structures-Octet Rule(All atoms want 8 electrons around them.)

Lewis came up with a way to draw valence electrons so that the bonding could be

determined.

Rules to Write Dot Structures1. Write a skeleton molecule with the lone atom in the middle (Hydrogen can never be in the middle)2. Find the number of electrons needed (N) (8 x number of atoms, 2 x number of H atoms)3. Find the number of electrons you have (valence e-'s) (H)4. Subtract to find the number of bonding electrons (N-H=B) 5. Subtract again to find the number of non-bonding electrons (H-B=NB)6. Insert minimum number of bonding electrons in the skeleton between atoms only. Add more bonding if needed until you have B bonding electrons.7. Insert needed non-bonding electrons around (not between) atoms so that all atoms have 8 electrons around them. The total should be the same as NB in 5 above.

Let's Try it!1.S

2.N

3.H

4.B

5.NB

6.E ..H:O:H ●●

H O H Water H2O

2 x 2 = 4 for Hydrogen1 x 8 = 8 for Oxygen4+8=12 needed electrons

8 – 4 = 4 non-bonding electrons

2 x 1 = 2 for Hydrogen1 x 6 = 6 for Oxygen You have 8 available electrons

12 - 8 = 4 bonding electrons

8 H

12 N

4 B

4 NB

-

-

H:O:H

..H:O:H ●●

Let's Try it!1.S

2.N

3.H

4.B

5.NB

6.E ..H:N:H ●●

HH N H Ammonia NH

3

3 x 2 = 6 for Hydrogen1 x 8 = 8 for Nitrogen6+8=14 needed electrons


3 x 1 = 3 for Hydrogen1 x 5 = 5 for Nitrogen You have 8 available electrons


8 H

14 N

6 B

2 NB

-

-

..H:N:H

..H:N:H ●●

H

HH

Let's Try it!1.S

2.N

3.H

4.B

5.NB

6.E .. ..O::C::O●● ●●

O C O Carbon Dioxide CO2

1 x 8 = 8 for Carbon2 x 8 = 16 for Oxygen8+16=24 needed electrons


1 x 4 = 4 for Carbon2 x 6 = 12 for Oxygen You have 16 available electrons


16 H

24 N

8 B

8 NB

-

-

O::C::O

.. .. O::C::O ●● ●●

Let's Try it!1.S

2.N

3.H

4.B

5.NB

6.E .. .. ..O::C: O:●● ●●

OO C O Carbonate CO

3-2

3 x 8 = 24 for Oxygen1 x 8 = 8 for Carbon24+8=32 needed electrons


3 x 6 = 18 for Oxygen1 x 4= 4 for Carbon You have 22 + 2 more available e-'s

24 H32 N

8 B

16 NB

-

-

..O::C:O

.. .. .. O::C: O: ●● ●●

O

..:O: ..

:O:


-2

Nomenclature of Covalently Bonded Compounds--Molecules

Non-Metals and Non-Metals

Use Prefixes such as mono, di, tri, tetra, penta, hexa, hepta, etc.

CO2 Carbon dioxide CO Carbon monoxide

PCl3 Phosphorus trichloride CCl4 Carbon tetrachloride

N2O5 Dinitrogen pentoxide CS2 Carbon disulfide

VSEPR Theory

Valence Shell Electron Pair Repulsion Theory—Geometric Shapes Linear, Bent Trigonal Planar, Trigonal Pyramidal Tetrahedral Trigonal Bipyramidal Octahedral

VSEPR Theory

Why is H2O bent and CO2 linear?

O in water has lone pairs causing bending whereas the C in carbon dioxide does not

VSEPR Theory

Lone pairs on the central atom cause crowding (INCREASED REPULSION) and result in bending

*Remember only the lone pairs on the central atom matter—the lone pairs on the external atoms do not crowd

Polarity of Molecules

Polar Molecule: a molecule that has uneven distribution of charge—dipole moments do NOT cancel

Nonpolar Molecule: a molecule that has even distribution of charge—all of the dipoles cancel

Polarity of Molecules Cont’d

Examples on the Board H2O

CH4

CO2

NH3

*BF Why does water exist as a liquid at STP and

carbon dioxide exists as a gas at STP?

Intermolecular Forces

Intermolecular Forces: forces of attraction that exist between two molecules Hydrogen Bonding: an IMF results from the

attraction between hydrogen and a highly electronegative element like F, N, or O Rather strong force Responsible for water’s high surface tension, holding

together DNA, and varying BPs and MPs

Intermolecular Forces

Dipole-Dipole Force: an IMF that exists between two polar molecules (hydrogen bonding is a special type of dipole-dipole force) Rather strong Used to predict MPs and BPs

Van der Waals Force: an IMF that exists between two nonpolar molecules Very weak Instanteous Predicts the low MPs and BPs of nonpolar molecues

Forces between Ionic Solids

Electrostatic Forces: a force of attraction that exists between ionic compounds due to opposite charges VERY strong Responsible for high MPs and high BPs

Hybridization Theory A theory that suggests that orbitals from

atoms will merge and create bonding orbitals of equivalent energy Sigma: bonding that occurs by overlapping

orbitals end to end Pi: bonding that occurs by overlapping orbitals

side to side

Hybridization Theory

Sigma Bonds () Pi Bond ()

Sigma and Pi Bonds

Single Bond: sigma only Double Bond: 1 sigma and 1 pi Triple: 1 sigma and 2 pi

How many sigma and pi bonds are found in the following: N2

C2H4

Types of Hybridization

sp sp2

sp3

sp3d sp3d2

Note: this theory uses both orbitals involved in bonding and orbitals holding lone pairs so CH4, NH3, and CH4 all have sp3 hybridization

Hybridization Theory

Molecular Orbital Theory

Based on quantum mechanics Treats the electron as a moving object Relates to probability of location not exact

location


Bonding Orbital: area of high electron probability that has lower energy than the orbitals of the separate atoms

Antibonding Orbital: area of high electron probability that has higher energy than the orbitals of the separate atoms


Nonbonding Orbital: an orbital that does not contribute stability nor does it destabilize the molecule

Open parking space near door Open parking space far from door Occupied parking space near door

Paramagnetism of Oxygen

MOT explains the paramagnetism of oxygen

Documents

Chemical Bonding and Nomenclature Adapted from Paul Surko