chapter 7 chemical formulas and...

BondingHonors Chemistry 412Chapter 6

• Chemical Bond▫ Mutual attraction between the nuclei and valence

electrons of different atoms that binds them together.

• Types of Bonds▫ Ionic Bonds Force of attraction between oppositely charged ions.

▫ Covalent Bond Force of attraction for electrons, that results in a pair of

electrons being shared by two atoms.

Ionic or Covalent?• The Difference in electronegativity determines the

bond type, page 161, according to the following scale:

• Example Na-Cl▫ Na = 0.9 Cl = 3.0 Difference = 3.0 – 0.9 = 2.1▫ Ionic

1.7

50%

3.3

100%

0.3

5%

Ionic Polar Covalent Non Polar

0

0%

Electronegativity difference

% ionic character

Polar Vs. Nonpolar Covalent BondsNon-Polar – electrons are

equally shared▫ Balanced Distribution of

Electrical Charge

Polar – electrons are not equally shared▫ Unbalanced

Distribution of Electrical Charge

δ + δ -

Identifying Bond Types• Determine if the following are going to form

Ionic, Polar Covalent or Non-polar Covalent Bonds:

C-N

Ca-F

Br-Br

H-Br

Solutions

• δ + C-N δ - (.5 Polar)

• +Ca-F -(3.0 Ionic)

• Br-Br (0 Non polar)

• δ + H-Br δ - (.7 Polar)

Example: H2O1. H: 1 each, O: 6 each, Total: 8 valence electrons2. Draw the skeleton:

H - O - H

3. Satisfy the central octet.4. Satisty all other octets.5. Check your work!

A. All atoms have octets/duetsB. Correct number of valence electrons was used

Duet Rule: Hydrogen is satisfied!

Summary of Bond Types

Why do atoms bond?• Atoms want to gain stability, like the noble gases

• Octet Rule ▫ Atoms tend to lose, gain, or share electrons in

order to attain a full set of 8 valence electrons.▫ Like a noble gas!

• Duet Rule▫ Hydrogen and Helium only▫ A full set consists of 2 valence electrons▫ First principal energy level can only accommodate

2 electrons

Covalent Bonds• Force of attraction between atoms that share

electrons

• Covalent Compound▫ Held together by covalent bonds▫ Usually contains two non-metals▫ The smallest piece is called a molecule

• Properties▫ Low Boiling/Melting points (weaker bonds)▫ Most are gases at room temperature▫ Non-conductors of electricity

Formation of Covalent BondsBond results from the attraction forces between the

nuclei and electrons of the two atoms• Attraction Forces ↑ - Potential Energy ↓• Repulsion Forces ↑ - Potential Energy ↑• If the attraction forces overcome the repulsion forces

the bond will form. (AF > RF)

Bond Length – the avg. distance between two atomsBond Energy – energy needed to break the bond

Overall: Attractive Forces ↑ - BL ↓ - BE ↑

Lewis Dot Structures• We can use Lewis dots to show covalent bonds:

H2: H H

F2: F F

We can use a dash (-) for bonds: H-H; F - F• Each dash stands for 2 electrons!

Single Covalent Bond

Unshared pair or “lone” pair(not involved in bonding)

Multiple Covalent Bonds

• Double Bond▫ 2 pairs of shared electrons: (=)

• Triple Bond▫ 3 pairs of shared electrons: ( )

F ___ ___ ___ ___ ___1s 2s 2p

F ___ ___ ___ ___ ___1s 2s 2p

Bonding Electron Pair

N ___ ___ ___ ___ ___1s 2s 2p

N ___ ___ ___ ___ ___1s 2s 2p

Triple Bond – 3 shared pairs

Steps to Drawing a Lewis Structure:1. Determine the total number of electrons contained in the

compound.

2. Draw the “skeleton” structure of the compoundA. If there are more than one type of atom, the least electronegative is

usually the center atomB. If there are multiples of the least electronegative, they will split the

centerC. Attach all other atoms to the center as symmetrically as possible

3. Satisfy the octet of the central atomA. Use single bonds to attach other atomsB. Fill in with unshared pairs as needed

4. Satisfy the octets of all other atoms following the same procedure

5. Check your work

Try some!• Write the Lewis Structure for the following

formulas:▫ NH3

HH - N - H

▫ SCl2Cl - S - Cl

Practice Single Covalent Bonds• Write the Lewis Structure for the following

formulas:▫ C2H6

▫ C3H7Cl


• What if I need to use too many electrons to get my Lewis Structure to work?▫ Indicates you need to make a multiple bond!▫ Rule of thumb: For every 2 electrons over the limit,

you need to make ONE more bond!

• CH2O:O – C – H

H

Too many electrons! You were only allowed to use 12, but there are 14 here!

Solution is to create a Double Bond!

• HCN:

H – C – N


Used 14 electrons, but you only have 10 available! You’ll need to make multiple bonds!Still too many – need to

make another bond!

Polyatomic Ion Structures• Polyatomic ions are covalently bonded atoms that

form a charge due to the gain or loss of electrons.

• When drawing the structures of polyatomic ions:

▫ (–) ions must have extra electrons in the structure that is equivalent to its charge.

▫ (+) ions must lose the number of electrons in the structure that is equivalent to its charge.

▫ Place the structure in brackets and indicate the charge at the top right hand corner

NH4+1

• Nitrogen: 5 valence electrons• Hydrogen: 1 valence electron each• +1 charge: Take away one valence electron

Total: 8 valence electrons

H

H N H

H

1+

For a polyatomic ion, always indicate the charge by putting the structure into brackets and writing the charge at the top right hand corner!

Exceptions to the Octet Rule• Incomplete Octet▫ Atoms that can be satisfied with less than eight

electrons. Boron can be satisfied with only six electrons.

• Expanded Octet▫ Atoms that can have more than eight electrons. (Not

ALWAYS)▫ Common elements that accommodate expanded octets: Cl, Br, I, S, P and all noble gases.

• IMPORTANT!!!▫ Only one atom may exceed the standard eight valence.▫ Must be the central atom only!!

Coordinate Covalent Bond• Like a single covalent bond, but a single atom is

sharing 2 electrons with another atom▫ Does not increase the # of electrons that the atom being

shared with “owns”

• This is determined by one atom having more electrons drawn than it originally contained.

• Indicated by using an arrow rather than a dash. () ▫ The arrow must point away from the atom that “owns”

the electrons being shared

Resonance• Bonding in molecules or ions that cannot be

correctly represented by a single Lewis Structure (in other words, can be drawn multiple ways)

• Molecules constantly “resonate,” or alternate between structures

• Shown by using a double headed arrow between resonance structures:

• CO2:

• O = C = O O ≡ C - O

Ionic Bonds• The result of the attraction between cations and

anions• Remember:▫ Cation: Positive ion, loses electrons▫ Anion: Negative ion, gains electrons

• In an ionic bond, electrons are exchanged between atoms

Ionic Compounds• Stronger than covalent bonds• Usually a metal with a non-metal• Neutral overall – charges must balance• Formula Unit▫ Lowest whole number ratio of atoms in an ionic

compound• Properties▫ High Melting Points (bonds are strong – need a lot

of energy to break)▫ Brittle▫ Dissolve in water▫ Are good conductors in the liquid phase (melted)

• Strength of Array – varies w/ sizes, charges, & number of ions.

• Arrangement of ionsgives the crystal its strength

3D Array of forces is created (Lattice)

CaF2

Lattice Energy• The energy released when one mole of an ionic

compound is formed from gaseous ions• Can be used to compare the strength of bonds (LE↑,

Bond Strength ↑)• Negative values indicate that energy is released• Examples:

Compound Lattice Energy (kJ/mol)

NaCl -787.5

NaBr -751.4

CaF2 -2634.7

LiCl -861.3

LiF -1032

MgO -3760

KCl -715

Ionic Bond Formation• We can use Lewis dot structures to show the exchange

of valence electrons in atoms as they form ionic bonds:Ca + Cl

Na + S

CaCl2Formula Unit of the Compound

Ca Cl

Cl

Practice:• Show the electron transfer for the following

elements and write the formula unit for the binary ionic compound:

1. Mg + F

2. K + P

3. Ba + Se

Metallic Bonding• Bonding resulting from the attraction of positive ions

and mobile electrons

• Only occurs in Metallic Atoms▫ Small Number of Valence Electrons▫ Low Ionization Energy & Electronegativity Easily give up electrons At best – weakly covalent

▫ Electrons are Delocalized Electrons do not belong to any one ion

Metallic Properties

Electron-Sea Model Explains Properties of Metals• Lustrous, Good Conductors of Heat & Electricity• Malleability and Ductility

• Strength of metallic bonds are reflected in the metal’s enthalpy of vaporization values▫ Amount of energy absorbed as heat when a specified

amount of substance vaporizes at constant pressure▫ Higher energy = stronger bonds

VSEPR• Valence Shell Electron Pair Repulsion▫ States that all atoms/electrons attached to a

central atom repel one another, and will situate themselves as far away from each other as possible

▫ Substituent – anything bonded to a central atom Can be another atom OR an unshared pair of

electrons

VSEPR

Orientation (aka Basic arrangement of substituents around the

central atom Depends ONLY upon how many substituents

there are

Shape What the molecule will actually look like in three

dimensions Depends upon how many AND what type of

substituents there are Unshared pairs of electrons are invisible!

Using VSEPR to predict shapes• For generic formulas:▫ A = Central Atom▫ B = Atom bonded to the central atom▫ E = Unshared pair of electrons on the central atom

**We are only concerned with unshared pairs that belong to the central atom!!**

Molecules with One Substituent

• Generic Formula:▫ AB

• Orientation:▫ Linear

• Shape:▫ Linear

**Note** You can add up to 5 unshared pairs of electrons to this, and it will NOT affect the shape of the

molecule!

Molecules with Two Substituents

• Generic Formula:▫ AB2

• Orientation:▫ Linear


Molecules with Three Substituents• Generic Formula:▫ AB3

• Orientation:▫ Trigonal Planar

• Shape:▫ Trigonal Planar

• Generic Formula:▫ AB2E

• Orientation:▫ Trigonal Planar

• Shape:▫ Angular/Bent

Molecules with Four Substituents• Generic Formula:▫ AB4

• Orientation:▫ Tetrahedral

• Shape:▫ Tetrahedral



• Shape:▫ Trigonal Pyramidal

• Generic Formula:▫ AB2E2


• Shape:▫ Angular/Bent

Molecules with Five Substituents

• Generic Formula:▫ AB5

• Orientation:▫ Trigonal Bipyramidal

• Shape:▫ Trigonal Bipyramidal



• Shape:▫ See-Saw

Molecules with Five Substituents• Generic Formula:▫ AB3E2


• Shape:▫ T shaped




Molecules with Six Substituents• Generic Formula:▫ AB6

• Orientation:▫ Octahedral

• Shape:▫ Octahedral



• Shape:▫ Square Pyramid



• Shape:▫ Square Planar

Using VSEPR to predict shape

• In order to use VSEPR to predict the shape of a molecule, you MUST have a valid Lewis Structure for the molecule!▫ Use it to identify the

number and types of substituents

▫ Identify the orientation and shape of the molecule

• SiO2:

O = Si = O

• AB2• 2 substituents▫ Orientation: Linear

• All substituents are atoms, no unshared pairs▫ Shape: Linear

Using VSEPR to predict shape• H2O

H – O – H

• AB2E2

▫ Orientation: Tetrahedral

▫ Shape: Bent

Using VSEPR to predict shape• ClO31-

• AB3E• O: Tetrahedral• S: Trigonal Pyramid

• SF2

• AB2E2• O: Tetrahedral• S: Angular/Bent

Using VSEPR to predict shape

• BF3

• AB3• O: Trigonal Planar• S: Trigonal Planar

• XeF4

• AB4E2• O: Octahedral• S: Square Planar

Molecular Polarity

• Polar Molecules

▫ Aka Dipoles▫ Entire molecules that

have an uneven distribution of charge

▫ Determined by looking at the Lewis structure, VSEPR, and bond polarity

• Example: H2O

O

H H

• O: Tetrahedral• S: Bent• E difference: 1.4

• Overall: Polar!

Molecular Polarity

• Polar Bonds do NOT definitely mean the molecule is polar:

• CCl4:

• O: Tetrahedral• S: Tetrahedral• E difference: 0.5

• The arrows cancel one another out, so this molecule is NON-POLAR!

• No one area has more negative chargeCl

ClCl

C

Cl

Hybridization• VSEPR told us about the shapes of molecules, but

not about the relationship between the shape and the orbitals that are occupied by the bonding electrons

• Hybridization – The mixing of 2 or more atomic orbitals of similar energies on the same atoms to produce an equivalent number of new orbitals with equal energy

Hybridization• Think about CH4:

• Carbon is the central atom, and its normal valence electrons look like this:

• Notice that now all of the electrons are unpaired and able to bond with other atoms!

• New orbitals are named after the original orbitals mixed together to create them:

2p

2sThese orbitals mix together to create 4 new, identical orbitals:

sp3

Hybridization

C H

H

H

H sp3 sp3

sp3

sp3

s

ss s

Possible Hybridization Combinations

• 1 s +1 p = 2 sp orbitals

• 1 s +2 p’s = 3 sp2 orbitals

• 1s +3 p’s = 4 sp3 orbitals

• 1 s +3 p’s +1 d = 5 sp3d orbitals

• 1 s +3 p’s +2 d’s = 6 sp3d2 orbitals

Types of Hybridization• Since hybridization is linked to orientation, here is an

easy guide:

• 2 substituents – Linear – sp

• 3 substituents – Trigonal Planar – sp2

• 4 substituents – Tetrahedral – sp3

• 5 substituents – Trigonal bipyramidal – sp3d

• 6 substituents – Octahedral – sp3d2

Practice• Determine the hybridization and Polarity

of the central atom for the following molecules:1) CO2

2) SO2

3) SF6

4) SiO2

1) O = C = O , (sp hybrid) on C

2 2) S ,(sp hybrid) on S

O O

••

F F

3) F S F , (sp3d2)

F F4) O = Si = O , (sp hybrid) on Si

Molecular Bonding• Molecular Orbitals▫ Resulting orbital formed by the overlap of 2 atomic

orbitals.• 2 Types of Molecular Bonds▫ Sigma Bond Molecular orbital that is formed along the bonding axis.

▫ Pi Bond Molecular orbital that is formed above or below the

bonding axis. Usually found as part of multiple bonds

Sigma Bond

Sigma with Pi bond

Sigma with 2 Pi bond

Sigma bond (σ) – electron density between the 2 atomsPi bond (π) – electron density above and below plane of nuclei of the bonding atoms 10.5

Sigma (σ) and Pi Bonds (π)

Single bond 1 sigma bond

Double bond 1 sigma bond and 1 pi bond

Triple bond 1 sigma bond and 2 pi bonds

How many σ and π bonds are in the acetic acid(vinegar) molecule CH3COOH?

C

H

H

CH

O

O Hσ bonds = 6 + 1 = 7π bonds = 1

10.5

Intermolecular Forces of Attraction• The forces of

attraction between molecules▫ Weaker than an ionic

or covalent bond▫ What pulls molecules

together in the liquid and solid phases

▫ Not all types of F of A are equivalent

• 3 types:▫ Dipole-Dipole

▫ Hydrogen Bonding

▫ London Dispersion

Intermolecular Forces of Attraction• Dipole-Dipole Attractions▫ Attraction between two polar

molecules▫ Partial + of one molecule

attracted to the partial – of another molecule

δ-

δ+

Intermolecular Forces of Attraction• Hydrogen Bonds▫ Special type of dipole-

dipole attraction▫ Attraction between H

in a covalently bonded molecule and unshared pairs of electrons in other molecules

▫ Responsible for many of the special behaviors of water!

Takes much more energy to break H bonds than normal forces of attraction, so water has a very high boiling point

H bonds are responsible for water freezing into a hexagonal pattern, creating an empty space which causes frozen water to have less density and float!

Intermolecular Forces of Attraction• London Dispersion Forces▫ Occur due to the movement of the electrons Electrons can bunch up on one side of the atom,

creating a temporary (induced) dipole, which can then influence other molecules.

▫ Weakest of all Forces of Attraction▫ Typically increase as the number of electrons in the

molecule increase If you look at the halogens, F and Cl are gases at room

temperature, Br is a liquid, and I is a solid!

chapter 7 chemical formulas and...

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