Chemical Bonding

What is a Bond?

• Force that holds atoms together

• Results from the simultaneous attraction of electrons (-) to the nucleus (+)

Breaking/Forming Bonds

• When a bond is broken energy is absorbed– Endothermic

• When a bond is formed energy is released– Exothermic

• The greater the energy released during the formation of the bond, the greater its stability– Stable bonds require a great deal of energy to

break

Lewis Dot Diagrams

• Use dots to represent the number of valence electrons

• How to write: – Write the symbol. – Put one dot for each valence electron – Electrons go on the 4 sides, no more than 2

per side

Dot Diagram Examples:

• Draw dot-diagrams for the following1.Mg

2.C

3.Ne

Dot Diagrams - Ions

• For ions, use brackets and place the charge outside the brackets

• Examples: 1. Na+

2. O2-

3. H+

Octet Rule

• Atoms will gain or lose electrons in order to have a full valence shell – like the nobles gases

• “Take the shortest route”

• Metals lose electrons to form positive ions (Cations)

• Nonmetals gain electrons to form negative ions (Anions)

Exceptions

• 1st principle energy level only holds 2 electrons• Transition elements can lose valence (s) and

inner (d) electrons – this is why they have multiple oxidation states

• Some atoms may be stable with less than an octet – many compounds with B

• Some atoms may be stable with more than an octet – elements beyond period 2, especially P and S, the additional electrons are added to the d sublevel

• Molecules with an odd number of electrons – they will be unstable

Types of Bonds

• Ionic - Electrons are transferred from a metal to a nonmetal

• Covalent - Electrons are shared between 2 nonmetals– Polar Covalent – electrons are shared unequally– Nonpolar Covalent – electrons are shared equally

• Metallic - Electrons are mobile within a metal, “Sea of Electrons”

Identifying Bond Type

• Ionic – metal and a nonmetal

• Covalent – 2 nonmetals– Nonpolar Covalent (equal)

• Same atoms (diatomics, triatomics)

– Polar Covalent (unequal)• Unequal electronegativity

• Metallic – metals

Identifying Bond Types

• Indicate the type of bond present in each: 1. HCl

2. CCl43. MgCl24. O2

5. Hg

6. H2O

Ionic Bonds• Transfer of 1 or more electrons

from a metal to a nonmetal• The greater the electronegativity

difference between atoms, the greater the ionic character

• Example: Sodium Chloride (NaCl)

Na ClX

Na electron transferred to Cl

Monatomic Ions

• One atom in an ion

• Look at the valence electrons to determine the charges

• Examples: K+, O2-

Polyatomic Ions

• More than one atom in the ion• Reference Table E• Charge belongs to the entire ion, not an

individual atom• Within the polyatomic ion the atoms are

held together by covalent bonds• When writing it, place ( ) around the entire

ion, with the charge outside• Examples: (NH4)+, (H3O)+, (CO3)2-

Writing Ionic Formulas

• You need an equal amount of positive and negative charges, so that the compound is neutral

• Ionic Formulas are always written as empirical formulas (reduced)

Examples

1. Na1+ + Cl1-

2. Mg2+ + Cl1-

3. Ca2+ + CO32-

4. Al3+ + O2-

Criss Cross Method

1. Write the symbol for the cation and anion

2. Write each ion’s charge as a superscript

3. Criss-cross the charges to become subscripts of the other ion

Do not put (+) or (-) charges in the final formula

4. Reduce to least common multiple (empirical formula)

Ionic Formulas

• Write the formula for the compound formed from the following ions:

1. Mg2+ + Cl-

2. Ca2+ + CO32-

3. Al3+ + O2-

4. Calcium ion + hydroxide ion

Naming Ionic Compounds

• Name the cation first, the anion second• Cation keeps its name, anion changes its

ending to –ide (Chlorine → Chloride)• Do not change the ending of polyatomic

ions• Examples:

1. NaCl

2. CaCO3

3. MgF2

Stock System – only used for positive ions

• Some cations have more than one positive oxidation states

• A roman numeral is used to indicate the charge of the positive ion

Stock System Examples

1. Iron (II) Chloride

2. Iron (III) Oxide

3. Copper (II) Oxide

4. a. What charge does copper have in copper II sulfate?

b. What is the formula for copper II sulfate?

Ionic Salts

• Salts are ionic compounds made up of cations and anions

• The ratio of cations to anions is always such that an ionic compound has no overall charge

• Many of the ions are bonded together to form a crystal

Properties of Ionic Salts

• Ionic Bonds are very strong

• Very high melting and boiling points– All solids at STP

• Hard

• Brittle

Melting and Boiling Points of Compounds

Compound Name

Formula Type of Compound mp

(oC)

bp

(oC)Magnesium Flouride MgF2 Ionic 1261 2512

Sodium Chloride NaCl Ionic 801 1686

Calcium Iodide CaI2 Ionic 784 1373

Iodine MonoChloride ICl Covalent 27 370

Carbon tetrachloride CCl4 Covalent -23 350

Hydrogen Flouride HF Covalent -83 293

Hydrogen Sulfide H2S Covalent -86 212

Methane CH4 Covalent -182 109

Properties of Salts (cont’d)

• Do not conduct electricity as solids

• Do conduct electricity when the salt melts or is dissolved in water (liquid phase or aqueous)– In order to conduct electricity a substance

must have free moving charged particles– In the solid phase the ions are not free to

move

Covalent Bonds• Sharing of electrons between 2 nonmetals

Non-Polar Covalent

• Electrons are shared equally• Equal distribution of electrons

• All diatomic molecules have non-polar covalent bonds

Nonpolar Covalent Examples

1. Flourine (F2)

2. Hydrogen (H2)

Polar Covalent

• Unequal Sharing of electrons• Unequal distribution of electrons

– Partial positive and partial negative charges– The side with the higher electronegativity will

have a greater share of the electron(s) resulting in a partial negative charge

• The greater the electronegativity difference, the more polar the bond is

Polar Covalent Examples

1. HCl

2. H2O

Dipoles

• Form when the charge in a bond is asymmetrical

– Present in polar bonds– Partial positive and partial negative charges

Polar Bonds / Dipoles

• Isn’t a whole charge just a partial charge • means a partially positive • means a partially negative

Example:

H - Cl

+---→• The Cl pulls harder on the electrons (more eneg)• The electrons spend more time near the Cl

Dipole Examples

1. Which molecule contains more polar bonds?

a. CCl4b. CH4

2. Which has a stronger dipole? a. HCl

b. HBr

Properties of Molecular Substances (Covalent Compounds)

• Soft

• Low melting points and boiling points– Many exist as gases or liquids at STP

• Poor conductors of heat and electricity (in all phases)

Examples: H2O, CCl4, NH3, C6H12O6, O2

Molecular Formulas (Covalent Compounds)

• Contain covalent bonds

• Tells you how many atoms are present in a single molecule

• Named similarly to ionic compounds, except use prefixes to indicate the number of atoms per molecule

Prefixes

• Mono- is only used for the second element– Example: CO = carbon monoxide

Mono- 1 Hexa- 6

Di- 2 Hepta- 7

Tri- 3 Octa- 8

Tetra- 4 Nona- 9

Penta- 5 Deca- 10

Examples

1. CCl42. H2O

3. NO

4. N2O5

5. BBr3

Structural Formulas

• Specifies how atoms are bonded together

• Dashes represent bonds• 2 atoms can share up to

3 pairs of electrons

Single Bonds

• 2 atoms share 1 pair of electrons (2 electrons)

Examples:

1. Ammonia (NH3)

2. Chlorine (Cl2)

3. Hydrochloric Acid (HCl)

Double Covalent Bonds

• 2 atoms share 2 pairs of electrons (4 electrons)

• 2 bonds between 2 atoms

Examples:

1. Carbon Dioxide (CO2)

2. Oxygen (O2)

Triple Covalent Bond

• 2 atoms share 3 pairs of electrons (6 electrons)

• 3 bonds between 2 atoms

Examples:

1. Nitrogen (N2)

2. Ethyne (C2H2)

Bond Length/Strength

• Length: – Single > Double > Triple

• The more electrons in a bond, the greater the attraction, therefore shorter

– As you move down a group bond length increases

• Due to increasing molecular size

• Strength:– Triple is the strongest, most stable, requires

the most energy to break

Network Solids

• Covalently bonded atoms are linked into a giant network (macromolecules)

• Examples: Diamond (C), Graphite (C), Silicon Carbide (SiC), and Silicon Dioxide (SiO2)

Network Solids

• Properties:– Hard– High melting and boiling points– Do not conduct heat and electricity

Metallic Bonding

• Sea of Electrons– Electrons are free to move through the solid.

+ + + ++ + + +

+ + + +

Properties of Metallic Solids

• Very Strong

• Good conductors of heat and electricity because electrons are free to move about

• Luster

• High melting point (except Hg)– All solids at STP (except Hg)

• Malleable, Ductile

VSEPR Theory

• In a small molecule, the electron pairs are as far away from each other as possible– VSEPR = Valence Shell Electron Pair

Repulsion

Linear

• Drawn on a straight line• All molecules of only 2 atoms are linear• Many 3 atom molecules are linear, if there are no

unshared electron pairs on the central atom• If both ends are the same, the molecule is

nonpolar (Symmetrical = Nonpolar)• If the ends are different, the molecule will be

polar (Asymmetrical = Polar)• Bond Angle = 180o

See Molecules• Examples: H2, CO2, HCl

Tetrahedral

Tetrahedral

• A central atom bonded to 4 other atoms

• 3-D shape allows the electron pairs to get as far away from each other as possible CH H

H

H109.5º

Tetrahedral

• If all the ends are the same, NONPOLAR

• If the ends are different, POLAR

• Bond Angle = 109.5o

See Molecules

• Examples:

1. CH4

2. CH3Cl

Pyramidial

• A central atom is bonded to 3 other atoms and the central atom has an unshared electron pair

• 3-D, like a pyramid

• Always POLAR

• Bond Angle = 107o

See Molecules

• Example: NH3

Bent

Bent

• A central atom is bonded to 2 other atoms and the central atom has 2 unshared electron pairs

• Always POLAR

• Bond angle = 105o

See Molecules

• Example: H2O

Intermolecular Attractions/Forces

• Forces between molecules• Determines boiling point, melting point, vapor

pressure, surface tension– The stronger the intermolecular attractions, the higher

the boiling point

• All intermolecular attractions are weaker than actual bonds

• Polar molecules will have stronger IMFs than nonpolar molecules– The greater the polarity the stronger the IMF

Dipole-Dipole Forces

• Occurs between 2 polar molecules

• The positive end of one molecule is attracted to the negative end of another molecule

• The greater the electronegativity difference is, the more polar the bond will be and the stronger the dipole will be

• Example: HCl

Dipole Examples

1. Which would have the strongest intermolecular forces? Explain Why.

a. HCl

b. HBr

2. Which would have the weakest intermolecular forces? Explain Why.

a. H2S

b. H2O

Hydrogen Bonds

• Special, Strong type of Dipole Attractions

• Attraction of a covalently

bonded H atom to a F, O, or N atom on another covalent compound

HHO

+ -

+

H HO+-

+

Hydrogen Bonds

• VERY STRONG

• Molecules with H bonds will have high boiling points, melting points, and surface tension– See Water Heating

• Example: NH3

H-bonds Examples

1. Which sample has Hydrogen Bonds?

a. H2 b. HF c. F2 d. HCl

2. Rank in order from strongest (1) to weakest (3). a. Hydrogen Bonds

b. Covalent Bonds

c. Dipole-Dipole Attractions

Molecule – Ion Attractions

• Attraction between a polar compound and an ion (ionic salt)

• Polar substances (such as water) attract ions from ionic compounds in solution

• This allows ionic substances to dissolve in polar solvents (water)– The anion is attracted to the positive end of the polar

solvent– The cation is attracted to the negative end of the polar

solvent– The ion dissociates (falls apart)Example: NaCl(aq)

Molecule-Ion Examples

1. Molecule-Ion attractions are present in which sample? a. HCl(l) c. KCl(l)b. HCl(aq) d. KCl(aq)

2. When sodium chloride dissolves in water the chloride ion is attracted toa. The positive part of the water, the O atomb. The negative part of the water, the O atomc. The positive part of the water, the H atomd. The negative part of the water, the H atom

Van Deer Waals Forces

• Very weak

• Exist between non-polar molecules

• Caused by momentary dipoles

• Increases as molecular mass increases

VDW Examples

1. Rank in order from weakest to strongest: – Hydrogen Bonds– Covalent Bonds– Van deer Waals Forces– Dipole-Dipole Attractions

2. Which would have the strongest intermolecular forces?

a. H2 b. Cl2 c. F2 d. Br2