Modern Chemistry Chapter 6- Chemical Bonding

• chemical bond- a mutual electrical attraction between the nuclei and the valence electrons of different atoms that binds the atoms together

• ionic bond- a chemical bond that results from the electrical attraction between anions and cations

Na + Cl NaCl

sodium loses its one valence electron to form the cation Na+

This allows the atom to become like neon with eight electrons in it outer energy level.

chlorine gains that electron to form the anion Cl-

This allows the atom to become like argon with eight electrons in its outer energy level.

Na+ is attracted to Cl- (opposites attract) and an ionic bond forms

• covalent bond- a chemical bond that results from the sharing of electron pairs between two atoms

• nonpolar covalent bond- a covalent bond in which the bonding electrons are shared equally by the atoms forming the bond

-because their electronegativities are essentially equal

• polar covalent bond- a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons

-because one atom has a greater electronegativity than the other

• Nonpolar covalent bond

H· ·H

Equal electronegativities of the hydrogen atoms cause the pair of electrons to be shared equally and a nonpolar covalent bond to form.

• Polar covalent bond

H · ·F+ -

Fluorine’s greater electronegativity causes the shared electrons to move closer to it and creates areas of slight positive and negative charge forming a polar covalent bond

Determining Bond Type• Find the absolute difference in the

electronegativities of the bonding atoms.• The greater the difference, the greater the

% ionic character which makes it more like an ionic bond.

– IF the absolute difference is• < 0.3 the bond is nonpolar covalent• > 0.3 and < 1.7 the bond is polar covalent• > 1.7 the bond is ionic

Do section review problems #1-4 on page 177.

Section Review page 1771- Electron pairs are shared in covalent bonds and electrons

are transferred between atoms in ionic bonds.

2-The difference in the electronegativities of bonding atoms determines the bond type.

For problem #3 use the electronegativity chart on page 161.

3a- Li = 1.0 F = 4.0 4.0-1.0 = 3.0 ionic bond

3b- Cu = 1.9 S = 2.5 2.5-1.9=0.6 polar covalent

3c- I = 2.5 Br = 2.8 2.8-2.5=0.3 polar covalent

4- c (0.3) < b (0.6) < a (3.0)

Covalent Bonding & Molecular Compounds

• molecule- a neutral group of atoms that are held together by covalent bonds

• molecular compound- a chemical compound whose simplest units are molecules

• chemical formula- indicates the relative numbers of atoms of each kind in a chemical compound by using element symbols and numerical subscripts

• molecular formula- shows the types and numbers of atoms combined in a single molecule of a molecular compound

• diatomic molecule- a molecule containing only two atoms

Formation of a Covalent Bond**-see figure 5 on page 179-**

• As two atoms come near one another, the nuclei of each atom are attracted to the electrons of the other atom.

• This causes a decrease in the potential energy of the atoms.

• As a bond between the atoms forms, the potential energy of the system reaches its lowest point. At this point, the two atoms “share” at least one pair of electrons which are then able to move freely between the nuclei of the two atoms in overlapping orbitals.

• If the atoms get closer to one another, repulsion between the nuclei increases and the potential energy increases.

Bond Length & Bond Energy

• bond length- the average distance between two bonded atoms

• bond energy- the energy required to break a chemical bond and form neutral, isolated atoms

– see figure 7 on page 181– see table 1 on page 182

What is the general relationship between bond length and bond energy (strength) as seen in table 1?

The Octet Rule• octet rule- chemical compounds

tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its outermost energy level

• EXCEPTIONS:– hydrogen atoms are complete with

two electrons (H2)– boron atoms are complete with 6

electrons (BF3)– some elements show expanded

valence involving “d” orbitals (PF5 & SF6)

Electron-Dot Notation

• electron-dot notation- is an electron configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the element symbol

¨

: F :

.

Lewis Structures

• Lewis structures- are formulas in which atomic symbols represent nuclei and inner shell electrons, dot pairs adjacent to a single atom represent unshared electron pairs, and dashes between the atomic symbols represent covalent bonds between two atoms

¨ ¨

:F-F: ¨ ¨

Rules for Drawing Lewis Structures

1- Determine the types and numbers of atoms in the molecule.2- Arrange the atoms to form a skeleton structure for the

molecule. If carbon is present, it is the central atom. Otherwise, the least electronegative atom (except for hydrogen) is central. Add the electron dot structure for each atom in the molecule. 3- Connect the atoms with lines to represent covalent bonds between shared electron pairs.4- Add unshared pairs of electrons so each atom (other than hydrogen) has eight electrons.

• structural formula- indicates the kind, number, arrangement, and bonds but not the unshared pairs of electrons of the atoms in a molecule

F-F

Covalent Bonds

single bond- is a covalent bond in which one pair of electrons is shared between two atoms

H· ·H H-H

-Do practice problems #1-4 on page 186

Drawing Lewis Structures

• page 186 #1

NH3 N (5 electrons) + 3 H (1 electron each)

H H

H N H H N H

Drawing Lewis Structures

• Page 186 #2 H2S

H S H S

H H

Page 186 #3 & 4

3- SiH4 H

H Si H

H

4- PF3

F P F

F

multiple covalent bonds- are double (two shared pairs of electrons) or triple (three shared pairs of electrons) bonds

:O=O: :N=N: ¨ ¨

-Do practice problems #1-2 on page 188.

-Do Section Review problems #2, 4, & 5 on page 189.

Drawing Lewis Structures

• page 188 #1 CO2

O C O O C O

Drawing Lewis Structures

• page 188 #2 HCN

H C N H C N

Section review page 1892-State the octet rule.

Atoms will gain, lose, or share electrons so that each atom has an octet of electrons.

4- a) I Br c) H C C Cl

b) H d) Cl

H C H Cl Si Cl

Br Cl

e) F O

F

Section review page 189

5- H N N H

H H

H N N H

Lewis Structure Practice• Draw the Lewis structures of the following

molecular compounds. Also use the model kits to build the molecule. The highlighted formulas represent molecules that contain multiple bonds.

NH3 CO2 N2 O2

HBr H2CO3 C2H6

C2H4

C2H2 PF3

NH3 H N H

H

HBr Br H

C2H2 H C C H

CO2 O C O

H2CO3 H C O O H

O

PF3 F P F

F

N2 N N

C2H6 H H

H C C H

H H

O2 O O

C2H4 H C C H

H H

Drawing Lewis Structures- review

1- Determine which element in the formula will be the central atom(s) of the structure.

2- Make a probable skeleton arrangement of the atoms.

3- Put the correct number of dots to equal the valence electrons of each atom.

4- Draw lines between single electrons of adjacent atoms.

5- If there are extra dots around adjacent atoms, draw multiple bond lines.

6- Make sure each atom has 8 electrons either in unshared pairs or shared bonds. Remember, hydrogen has just two electrons.

Lewis Structure QuizDraw the Lewis structures for the following

compounds.

1- H2O

2- PF3

3- SiO2

4- SeBr2

5- CS2

Ionic Bonding & Ionic Compounds

• ionic compound- is composed of positive and negative ions that are combined so that the number of positive and negative charges are equal

• formula unit- is the simplest collection of atoms from which an ionic compound’s formula can be established

eg.NaCl

Formation of Ionic Bonds

• An atom of an element with low electronegativity approaches another with high electronegativity.

• The highly electronegative atom then transfers an electron from the atom with low electronegativity.

• This creates an anion and a cation.• The attraction between the ions forms an ionic

bond.

Na + Cl Na+ + Cl -

Ionic Crystals

• Ionic compounds tend to form an orderly arrangement known as a crystal lattice which then forms crystals.

***see figure 14 on page 191***

Comparing Ionic & Molecular Compounds

• ionic compound molecular compoundhigh melting point lower melting point

high boiling point lower boiling point

extreme hardness lower hardness (usually)

brittle less brittle

Polyatomic Ions

• polyatomic ion- a covalently bonded group of atoms with a positive or a negative charge

• Review the list of polyatomic ions given to you by the teacher.

Metallic Bonding

• metallic bond- a chemical bond resulting from the attraction between metal atoms and the surrounding sea of electrons

• The ability of the electrons to move freely between the nuclei of the metal atoms accounts for the unique properties of metals.

– This accounts for their being good conductors of heat and electricity.

Properties of Metals

• luster- the ability to absorb light energy and immediately re-emit it at the same or similar frequency which makes them reflective.

• malleable- the ability of a substance to be hammered or beaten into thin sheets

• ductility- the ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire

Molecular Geometry

• VSEPR theory (valence shell electron pair repulsion)- allows us to predict the shape of molecules. It states that repulsion between the sets of valence level electrons surrounding an atom causes these sets to be oriented as far apart as possible

VSEPR Theory• When determining the shape of a molecule using VSEPR, use the

following steps:

Draw the Lewis structure of the compound. Find the central atom(s). Use the letter “A” for this

atom. Count the number of atoms bonded to the central

atom. Use the letter ”B” and a subscript for the number of atoms bonded to the central atom “A”.

Count the number of unshared electron pairs around the central “A” atom.

Use the letter “E” and a subscript for the number of unshared electron pairs around the central atom

Use this “ABE” designation to find the molecular shape using table 5 on page 200 of the textbook.

• Do practice problems #1 & 2 on page 201.

Practice problems page 201

1a- F-S-F S = A F = B = 2 E = 2

AB2E2 bent or angular

1b- Cl-P-Cl P = A B = Cl = 3 E = 1

Cl

AB3E trigonal-pyrimidal

H

H-C-N-H

H H

This molecule has two central atoms (C & N) so it has two molecular shapes that are combined.

Hybridization

• hybridization- is the mixing or two or more atomic orbitals of similar energies on the same atom to produce new hybrid atomic orbitals of equal energies

• hybrid orbitals- orbitals of equal energy produced by the combination of two or more orbitals on the same atom

• Hybridization explains the unique qualities of a carbon atom with its “sp3” orbitals.

Intermolecular Forcesintermolecular forces- the forces of attraction

between molecules

1. dipole-dipole forces- forces of attraction between polar molecules

2. hydrogen bonding- intermolecular force in which a hydrogen atom bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule

3. London Dispersion forces- intermolecular attraction resulting from the constant motion of electrons and the creation of instantaneous dipoles

Chapter 6 Review

• Do the following review problems from pages 209-211 of the textbook.

#6, 15, 19, 21, 34,

43, & 48

End of Chapter Practice

#6 H (2.1) I (2.5) 0.4 = polar covalent

S (2.5) O (3.5) 1.0 = polar covalent

K (0.8) Br (2.8) 2.0 = ionic

Si (1.8) Cl (3.0) 1.2 = polar covalent

K (0.8) Cl (3.0) 2.2 = ionic

Se (2.4) S (2.5) 0.1= nonpolar covalent

C (2.5) H (2.1) 0.4 = polar covalent

#15 H = 1 F = 7 Mg = 2 O = 6

Al = 3 N = 5 C = 4

End of Chapter Practice

#19 Li Ca Cl

O C P

Al S

End of Chapter Practice

F

#21 F C F H Se H

F

Br

I N I Br Si Br

I Br

Cl

H C H

H

End of Chapter Practice

#34 AB2 = linear

AB3 = trigonal planar

AB4 = tetrahedral

AB5 = trigonal bipyramidal

AB6 = octahedral

End of Chapter Practice

#43 AB3E = trigonal pyramidal

AB2E2 = bent or angular

AB2E = bent or angular

Honors Chemistry Chapter 6 Test Review

40 multiple choice questions: valence electrons, chemical bonds (how they occur) atoms & potential energy leading to stability and bond

formation polar & nonpolar covalent bonds difference in electronegativity & % ionic character using electronegativities, determine if a bond is ionic,

polar or nonpolar covalent definition of molecule, molecular formula (&examples),

bond length, octet & octet rule elements meeting octet rule naturally how to draw a Lewis structure, identifying a Lewis

structure, bonding in Lewis structures

Honors Chemistry Chapter 6 Test Review

formula of an ionic compound represents… lattice energy, crystal lattice compare properties of ionic & molecular compounds and

the strength of their bonds electrons & charge of polyatomic ions metallic bonds & their electrons properties of metals & the cause properties of ionic crystals VSEPR definition & use intermolecular forces (dipole-dipole, hydrogen bonding,

London dispersion), their relative strength and properties

Chemistry Chapter 6 Test Review

25 multiple choice questions definitions and functions of valence electrons & chemical bonds bonding & potential energy polar & nonpolar covalent bonds & their characteristics use difference in electronegativity to determine bond type define molecule, molecular formula, & octet which elements satisfy the octet rule by themselvs which elements form multiple covalent bonds what is necessary to draw a Lewis structure & recognize a correct Lewis structure properties of ionic vs. covalent compounds excess (or deficit) electrons in polyatomic ions valence electrons in metallic bonds properties of metals and why they occur definition of VSEPR intermolecular forces and why they occur, especially dipole-dipole forces polar molecules

VSEPR Lab

H2O H O AB2E2

H

CO2 O C O AB2

H AB4

CH3NH2 H C N H

H H AB3E

VSEPR Lab

H2CO H C O AB3

H

CH4 H

H C H AB4

H

VSEPR Lab

C2H6 H H AB4

H C C H

H H AB4

C2H2 H C C H AB2

AB2

VSEPR Lab

HCOOH H C O H AB3

AB2E2

O