Chemical Bonding Chemical Bonding and Nomenclatureand Nomenclature

By Paul SurkoNew Dimensions High SchoolPoinciana, FL

s

8

I want you to meet a friend of mine?Bonding, the way atoms are attracted to each other to form molecules, determines nearly all of the chemical properties we see. And, as we shall see, the number “8” is very important to chemical bonding.

5.1 What are Molecules?5.1 What are Molecules?Molecules are a

combination of atoms bonded together. Bonding determines the chemical

properties of the molecule (compound).

5.5 Ionic Bonding-Being Like 5.5 Ionic Bonding-Being Like the Noble Gasesthe Noble Gases

All atoms want to have the same number of electrons as the Noble Gases. The Noble Gases have very stable electron configurations. In order to achieve the same electron configuration as the Noble Gases metal atoms will 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+3

Cl becomes Cl- O becomes O-2

Mg becomes Mg+2Na becomes Na+

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

Opposites Attract!Opposites Attract!

Putting Ions TogetherPutting Ions Together

Na+ + Cl- = NaCl

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

Al+3 + S-2 = Al2S

3Ca+2 + N-3 = Ca

3N

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

5.2 The Covalent Bond5.2 The Covalent BondAtoms can form molecules by sharing Atoms can form molecules by sharing

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

5.3 Dot Structures-Octet Rule5.3 Dot Structures-Octet Rule(All atoms want 8 electrons around them.)(All atoms want 8 electrons around them.)

Valence electrons are those in the outermost orbitals. They are the ones that can form bonds.

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

Rules to Write Dot StructuresRules to Write Dot Structures1.1. Write a skeleton molecule with the lone atom in the middle Write a skeleton molecule with the lone atom in the middle (Hydrogen can never be in the middle)(Hydrogen can never be in the middle)2.2. Find the number of electrons needed (N) Find the number of electrons needed (N) (8 x number of atoms, 2 x number of H atoms)(8 x number of atoms, 2 x number of H atoms)3.3. Find the number of electrons you have (valence eFind the number of electrons you have (valence e--'s) (H)'s) (H)4.4. Subtract to find the number of bonding electrons (N-H=B) Subtract to find the number of bonding electrons (N-H=B) 5.5. Subtract again to find the number of non-bonding Subtract again to find the number of non-bonding electrons (H-B=NB)electrons (H-B=NB)6.6. Insert minimum number of bonding electrons in the Insert minimum number of bonding electrons in the skeleton between atoms only. Add more bonding if needed skeleton between atoms only. Add more bonding if needed until you have B bonding electrons.until you have B bonding electrons.7.7. Insert needed non-bonding electrons around (not Insert needed non-bonding electrons around (not between) atoms so that all atoms have 8 electrons around between) atoms so that all atoms have 8 electrons around them. The total should be the same as NB in 5 above.them. The total should be the same as NB in 5 above.

Let's Try it!Let's Try it!1.1.SS

2.2.NN

3.3.HH

4.4.BB

5.5.NBNB

6.6.EE ..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!Let's Try it!1.1.SS

2.2.NN

3.3.HH

4.4.BB

5.5.NBNB

6.6.EE ..H:N:H ●●

HH N H Ammonia NH

3

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

8 – 6 = 2 non-bonding electrons

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

14 - 8 = 6 bonding electrons

8 H

14 N

6 B

2 NB

-

-

..H:N:H

..H:N:H ●●

H

HH

Let's Try it!Let's Try it!1.1.SS

2.2.NN

3.3.HH

4.4.BB

5.5.NBNB

6.6.EE .. ..O::C::O●● ●●

O C O Carbon Dioxide CO2

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

16 – 8 = 8 non-bonding electrons

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

24 - 16 = 8 bonding electrons

16 H

24 N

8 B

8 NB

-

-

O::C::O

.. .. O::C::O ●● ●●

Let's Try it!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

24 – 8 = 16 non-bonding 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:

32 - 24 = 8 bonding electrons

-2

5.6 Polarity-Unequal Sharing 5.6 Polarity-Unequal Sharing of Electronsof Electrons

Even though all atoms want the same number of electrons as the Noble Gases, some want to get or give them more than others. The magnitude of this attraction for electrons is called “Electronegativity”. The more electronegative an atom is, the more it wants the electrons.

Some atoms want to gain electrons so bad, they take them altogether to form negative ions. Some want to lose them so bad that they become positive ions.

Examples of Polar and Non-Examples of Polar and Non-Polar CompoundsPolar Compounds

H2O Water is a bent molecule. The lone pair of electrons

from the Lewis structure distorts its shape and it becomes a very polar molecule.

NaCl Since Na is a metal it gives up its electron to form Na+ and Cl takes the electron completely to form Cl-.

HCl The Chlorine wants the electrons more than the Hydrogen. Thus we have +δHCl-δ.

Cl2 (Cl—Cl) The Chlorine molecules want the electrons

equally so they form a non-polar molecule with NO partial or full charges.

CO2 Carbon Dioxide is a linear molecule. It has no lone pairs

of electrons from the Lewis structure. The two oxygen atoms pull equally and make it a non-polar molecule.

..:O:H ●●

H

.. ..O::C::O●● ●●

5.7 Nomenclature5.7 NomenclatureNaming of CompoundsNaming of Compounds

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

5.7 Nomenclature5.7 NomenclatureNaming of CompoundsNaming of Compounds

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

Metals (Transition Metals) and Non-MetalsOlder System

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

Ferrous Bromide FeBr2

Compare with Ferric Bromide FeBr3

Non-Metals and Non-MetalsUse 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

Let’s Practice!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 IonsPolyatomic Ions(partial list from page 195 (193 2(partial list from page 195 (193 2ndnd edition)) edition))

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

NHNH44++

NONO33--

MnOMnO44--

ClOClO33--

OHOH--

CNCN--

SOSO4 4 2 -2 -

COCO332-2-

CrOCrO442-2-

CC22HH33OO22--

POPO443-3-

Acids Acids (with H in front)(with H in front)Binary acids (without oxygen in formula)

Hydro _________ ic Acid

HCl Hydrochloric acid HBr Hydrobromic acid

Oxy acids (with oxygen in formula)

-ate goes to –ic and –ite goes to -ous

HNO3 Nitric acid HNO2 Nitrous acid

H2SO4 Sulfuric acid H2SO3 Sulfurous acid

H3PO4 Phosphoric acid H3PO3 Phosphorous acid

Lets Practice!Lets Practice!

HFNa2CO3

H2CO3

KMnO4

HClO4

H2S

NaOH

CuSO4

PbCrO4

H2O

NH3

Hydrooxic acid (no……just water)

Nitrogen trihydride (no..just ammonia)

Copper (II) sulfate or Cupric sulfate

Lead (II) chromate or Plubous chromate

Sodium hydroxide

Hyrdogen sulfuric acidPerchloric acid

Potassium permanganate

Sodium carbonate

Hydroflouric acid

Carbonic acid