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 following

1. 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”

Dog Analogy Ionic Bonds

big greedy dog stealing the other dogs bone Polar Covalent Bonds

Unevenly matched but willing to share Nonpolar Covalent Bonds

Dogs of equal strength Metallic Bonds

Mellow dogs with plenty of bones to go around

See the Dogs

Identifying Bond Type Ionic – metal and a nonmetal Covalent – 2 nonmetals Metallic – metals

OR Use electronegativity differences Ionic: 1.7 or more Polar Covalent: 0.5-1.6 Nonpolar Covalent: 0.0-0.4

Identifying Bond Types

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

2. CCl43. MgCl24. O2

5. Hg

6. H2O

Ionic BondsIonic Bonds

Transfer of 1 or more electrons Transfer of 1 or more electrons from a metal to a nonmetalfrom a metal to a nonmetalElectronegativity difference is Electronegativity difference is ≥ 1.7≥ 1.7

Example: Sodium Chloride (NaCl)Example: Sodium Chloride (NaCl)Na ClX

Na electron transferred to Cl

Monatomic IonsMonatomic Ions

One atom in an ionOne atom in an ion

Look at the valence electrons to determine Look at the valence electrons to determine the chargesthe charges

Examples: KExamples: K++, O, O2-2-

Polyatomic IonsPolyatomic Ions

More than one atom in the ionMore than one atom in the ionReference Table EReference Table ECharge belongs to the entire ion, not an Charge belongs to the entire ion, not an individual atomindividual atomWithin the polyatomic ion the atoms are Within the polyatomic ion the atoms are held together by covalent bondsheld together by covalent bondsWhen writing it, place ( ) around the entire When writing it, place ( ) around the entire ion, with the charge outsideion, with the charge outside

Examples: (NHExamples: (NH44))++, (H, (H33O)O)++, (CO, (CO33))2-2-

Writing Ionic FormulasWriting Ionic Formulas

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

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

ExamplesExamples

1.1. NaNa1+1+ + Cl + Cl1-1-

2.2. MgMg2+2+ + Cl + Cl1-1-

3.3. CaCa2+2+ + CO + CO332-2-

4.4. AlAl3+3+ + O + O2-2-

Criss Cross MethodCriss Cross Method

1.1. Write the symbol for the cation and anionWrite the symbol for the cation and anion

2.2. Write each ion’s charge as a superscriptWrite each ion’s charge as a superscript

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

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

4.4. Reduce to least common multiple Reduce to least common multiple (empirical formula)(empirical formula)

Ionic FormulasIonic Formulas

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

1.1. MgMg2+2+ + Cl + Cl--

2.2. CaCa2+2+ + CO + CO332-2-

3.3. AlAl3+3+ + O + O2-2-

4.4. Ca + OHCa + OH

Naming Ionic CompoundsNaming Ionic Compounds

Name the cation first, the anion secondName the cation first, the anion secondCation keeps its name, anion changes its Cation keeps its name, anion changes its ending to –ide (Chlorine ending to –ide (Chlorine → Chloride)→ Chloride)Do not change the ending of polyatomic Do not change the ending of polyatomic ionsionsExamples: Examples:

1.1. NaClNaCl

2.2. CaCOCaCO33

3.3. MgFMgF22

Stock System – Stock System – only used for positive ionsonly used for positive ions

Some cations have more than one Some cations have more than one positive oxidation statespositive oxidation states

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

Stock System ExamplesStock System Examples

1.1. Iron (II) ChlorideIron (II) Chloride

2.2. Iron (III) OxideIron (III) Oxide

3.3. Copper (II) OxideCopper (II) Oxide

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

b. What is the formula for copper II sulfate?b. What is the formula for copper II sulfate?

Ionic SaltsIonic Salts

Salts are ionic compounds made up of cations Salts are ionic compounds made up of cations and anionsand anions

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

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

Properties of Ionic SaltsProperties of Ionic Salts

Ionic Bonds are very strongIonic Bonds are very strong

Very high melting and boiling pointsVery high melting and boiling points

HardHard

BrittleBrittle

Properties of Salts (cont’d)Properties of Salts (cont’d)

Do not conduct electricity as solidsDo not conduct electricity as solids

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

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

movemove

Melting and Boiling Points of Melting and Boiling Points of CompoundsCompounds

Compound Compound NameName

FormulaFormula Type of CompoundType of Compound mpmp

((ooC)C)

bpbp

((ooC)C)Magnesium FlourideMagnesium Flouride MgFMgF22 Ionic Ionic 12611261 25122512

Sodium ChlorideSodium Chloride NaClNaCl IonicIonic 801801 16861686

Calcium IodideCalcium Iodide CaICaI22 IonicIonic 784784 13731373

Iodine MonoChlorideIodine MonoChloride IClICl CovalentCovalent 2727 370370

Carbon tetrachlorideCarbon tetrachloride CClCCl44 CovalentCovalent -23-23 350350

Hydrogen FlourideHydrogen Flouride HFHF CovalentCovalent -83-83 293293

Hydrogen SulfideHydrogen Sulfide HH22SS CovalentCovalent -86-86 212212

MethaneMethane CHCH44 CovalentCovalent -182-182 109109

Covalent BondsCovalent BondsSharing of electrons between 2 nonmetalsSharing of electrons between 2 nonmetalsElectronegativity difference is Electronegativity difference is ≤ 1.6≤ 1.6

Non-Polar CovalentNon-Polar Covalent

Electrons are shared equallyElectrons are shared equally Uniform distribution of electronsUniform distribution of electrons Bond is symmetricalBond is symmetrical Electronegativity difference of 0-0.4Electronegativity difference of 0-0.4 All diatomic molecules have non-polar All diatomic molecules have non-polar

covalent bondscovalent bonds

Nonpolar Covalent ExamplesNonpolar Covalent Examples

1.1. Flourine (FFlourine (F22))

a.a. e-neg difference = e-neg difference =

b.b. Dot diagram: Dot diagram:

2.2. Hydrogen (HHydrogen (H22))

a.a. e-neg difference = e-neg difference =

b.b. Dot diagram: Dot diagram:

Polar CovalentPolar Covalent Unequal Sharing of electronsUnequal Sharing of electrons Unequal distribution of electronsUnequal distribution of electrons

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

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

Electronegativity difference of 0.5-1.6Electronegativity difference of 0.5-1.6

Polar Covalent ExamplesPolar Covalent Examples

1.1. HClHCla.a. e-neg difference: e-neg difference:

b.b. Dot diagram: Dot diagram:

2.2. HH22OO

a.a. e-neg difference: e-neg difference:

b.b. Dot diagram: Dot diagram:

DipolesDipoles

Form when the charge in a bond is Form when the charge in a bond is asymmetricalasymmetrical

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

Polar Bonds / DipolesPolar Bonds / Dipoles

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

Example: Example:

H H - Cl- Cl

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

Dipole ExamplesDipole Examples

1.1. Which molecule contains more polar Which molecule contains more polar bonds? bonds?

a. CCla. CCl44

b. CHb. CH44

2. Which has a stronger dipole? 2. Which has a stronger dipole? a.a. HClHCl

b.b. HBrHBr

Properties of Molecular Substances Properties of Molecular Substances (Covalent Compounds)(Covalent Compounds)

SoftSoftLow melting points and boiling pointsLow melting points and boiling points

Many exist as gasesMany exist as gasesPoor conductors of heat and electricity (in Poor conductors of heat and electricity (in

all phases)all phases)

Examples: HExamples: H22O, CClO, CCl44, NH, NH33, C, C66HH1212OO66, O, O22

Molecular Formulas Molecular Formulas (Covalent Compounds)(Covalent Compounds)

Contain covalent bondsContain covalent bondsTells you how many atoms are present in a Tells you how many atoms are present in a

single moleculesingle moleculeNamed similarly to ionic compounds, Named similarly to ionic compounds,

except use prefixes to indicate the number except use prefixes to indicate the number of atoms per moleculeof atoms per molecule

PrefixesPrefixes

Mono- is only used for the second elementMono- is only used for the second element Example: CO = carbon monoxideExample: CO = carbon monoxide

Mono-Mono- 11 Hexa-Hexa- 66

Di-Di- 22 Hepta-Hepta- 77

Tri-Tri- 33 Octa-Octa- 88

Tetra-Tetra- 44 Nona-Nona- 99

Penta-Penta- 55 Deca-Deca- 1010

ExamplesExamples

1.1. CClCCl44

2.2. HH22OO

3.3. NONO

4.4. NN22OO55

5.5. BBrBBr33

Structural FormulasStructural Formulas

Specifies how atoms Specifies how atoms are bonded togetherare bonded together

Dashes represent Dashes represent bondsbonds

2 atoms can share up 2 atoms can share up to 3 pairs of electronsto 3 pairs of electrons

Single BondsSingle Bonds

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

Examples: Examples:

1.1. Ammonia (NHAmmonia (NH33))

2.2. Chlorine (ClChlorine (Cl22))

3.3. Hydrochloric Acid (HCl)Hydrochloric Acid (HCl)

Double Covalent BondsDouble Covalent Bonds

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

2 bonds between 2 atoms2 bonds between 2 atoms

Examples: Examples:

1.1. Carbon Dioxide (COCarbon Dioxide (CO22))

2.2. Oxygen (OOxygen (O22))

Triple Covalent BondTriple Covalent Bond

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

3 bonds between 2 atoms3 bonds between 2 atoms

Examples: Examples:

1.1. Nitrogen (NNitrogen (N22))

2.2. Ethyne (CEthyne (C22HH22))

Bond Length/StrengthBond Length/Strength

Length: Length: Single Single > Double > Triple> Double > Triple

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

As you move down a group bond length As you move down a group bond length increasesincreasesDue to increasing molecular sizeDue to increasing molecular size

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

the most energy to break the most energy to break

Network SolidsNetwork Solids

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

Examples: Diamond Examples: Diamond (C), Graphite (C), (C), Graphite (C), Silicon Carbide (SiC), Silicon Carbide (SiC), and Silicon Dioxide and Silicon Dioxide (SiO(SiO22))

Network SolidsNetwork Solids

Properties:Properties:HardHardHigh melting and boiling pointsHigh melting and boiling pointsDo not conduct heat and electricityDo not conduct heat and electricity

Metallic BondingMetallic Bonding

Sea of ElectronsSea of ElectronsElectrons are free to move through the solid. Electrons are free to move through the solid.

+ + + ++ + + +

+ + + +

Properties of Metallic SolidsProperties of Metallic Solids

Very StrongVery StrongGood conductors of heat and electricity Good conductors of heat and electricity

because electrons are free to move aboutbecause electrons are free to move aboutLusterLusterHigh melting point (except Hg)High melting point (except Hg)Malleable, DuctileMalleable, Ductile

VSEPR TheoryVSEPR Theory

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

RepulsionRepulsion

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

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

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

(Asymmetrical = Polar)(Asymmetrical = Polar) Bond Angle = 180Bond Angle = 180oo

See MoleculesSee Molecules Examples: HExamples: H22, CO, CO22, HCl, HCl

Trigonal Planar

Trigonal PlanarTrigonal Planar

A central atom is bonded to 3 other atoms, A central atom is bonded to 3 other atoms, with no extra electrons on the central atomwith no extra electrons on the central atom

Forms a flat “Y” shape (triangle shape)Forms a flat “Y” shape (triangle shape) If the ends are all the same, NONPOLARIf the ends are all the same, NONPOLAR If the ends are different, POLARIf the ends are different, POLARBond Angle = 120Bond Angle = 120oo

See MoleculesSee Molecules

Examples: BClExamples: BCl33, BH, BH22FF

PyramidialPyramidial

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

3-D, like a pyramid3-D, like a pyramidAlways POLARAlways POLARBond Angle = 107Bond Angle = 107oo

See MoleculesSee MoleculesExample: NHExample: NH33

Tetrahedral

TetrahedralTetrahedral

A central atom bonded A central atom bonded to 4 other atomsto 4 other atoms

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

H

H109.5º

TetrahedralTetrahedral

If all the ends are the same, NONPOLARIf all the ends are the same, NONPOLAR If the ends are different, POLARIf the ends are different, POLARBond Angle = 109.5Bond Angle = 109.5oo

See MoleculesSee MoleculesExamples: Examples:

1. CH1. CH44

2. CH2. CH33ClCl

Bent

BentBent

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

Always POLARAlways POLARBond angle = 105Bond angle = 105oo

See MoleculesSee MoleculesExample: HExample: H22OO

Intermolecular Intermolecular Attractions/ForcesAttractions/Forces

Forces between moleculesForces between molecules Determines boiling point, melting Determines boiling point, melting

point, vapor pressure, surface point, vapor pressure, surface tensiontension The stronger the intermolecular The stronger the intermolecular

attractions, the higher the boiling pointattractions, the higher the boiling point All intermolecular attractions are All intermolecular attractions are

weaker than actual bondsweaker than actual bonds

Dipole-Dipole ForcesDipole-Dipole Forces

Occurs between 2 polar moleculesOccurs between 2 polar molecules The positive end of one molecule is The positive end of one molecule is

attracted to the negative end of attracted to the negative end of another moleculeanother molecule

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

Example: HClExample: HCl

Dipole ExamplesDipole Examples

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

a. HCla. HCl

b. HBrb. HBr

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

a. Ha. H22SS

b. Hb. H22O O

Hydrogen BondsHydrogen Bonds

Special, Strong type of Special, Strong type of Dipole AttractionsDipole Attractions

Attraction of a covalently Attraction of a covalently

bonded bonded HH atom to a atom to a F, F, O, or NO, or N atom on atom on another covalent another covalent compoundcompound

HH

O+ -

+

H HO+-

+

Hydrogen BondsHydrogen Bonds

VERY STRONGVERY STRONG Molecules with H bonds will have Molecules with H bonds will have

high boiling points, melting points, high boiling points, melting points, and surface tensionand surface tension

Example: NHExample: NH33

H-bonds ExamplesH-bonds Examples

1. Which sample has Hydrogen 1. Which sample has Hydrogen Bonds? Bonds?

a. Ha. H22 b. HF c. F b. HF c. F22 d. HCl d. HCl

2. Which is the strongest?2. Which is the strongest?a. Hydrogen Bondsa. Hydrogen Bonds

b. Covalent Bondsb. Covalent Bonds

c. Dipole-Dipole Attractions c. Dipole-Dipole Attractions

Molecule – Ion Molecule – Ion AttractionsAttractions

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

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

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

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

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

Molecule-Ion ExamplesMolecule-Ion Examples

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

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

Van Deer Waals ForcesVan Deer Waals Forces

Very weakVery weak Exist between non-polar moleculesExist between non-polar molecules Caused by momentary dipolesCaused by momentary dipoles Increases as molecular mass Increases as molecular mass

increasesincreases

VDW ExamplesVDW Examples

1. Rank in order from weakest to 1. Rank in order from weakest to strongest: strongest:

Hydrogen BondsHydrogen Bonds Covalent BondsCovalent Bonds Van deer Waals ForcesVan deer Waals Forces Dipole-Dipole AttractionsDipole-Dipole Attractions

2. Which would have the strongest 2. Which would have the strongest intermolecular forces? intermolecular forces?

a. Ha. H22 b. Cl b. Cl2 2 c. F c. F22 d. Br d. Br22