Chemical Bonding AP Chemistry

Ms. Grobsky

Lesson Objectives

• Describe the basic form of chemical bonding – ionic and covalent – and the differences between them

• Predict from the formula whether a compound has ionic or covalent bonding

• Describe the bases of the localized electron (LE) model and Lewis theory of bonding

• Determine the number of valence electrons for any atom or ion and write its Lewis symbol

• Draw Lewis structures for molecular compounds • Understand and apply the octet rule

▫ Recognize exceptions to the octet rule

What is Bonding?

Why do Atoms Bond? • Bonding is the interplay between interactions between atoms

▫ Energetically favored

Electrons on one atom interacting with protons of another atom

▫ Energetically unfavorable

Electrons on one atom interacting with electrons of another atom

Protons on one atom interacting with protons of another atom

• A bond will form if the system can LOWER its total energy in the process

Illustration of Bonding

Types of Bonds

• Ionic bond ▫ Bond between a metal cation and non-metal

anion

Formula determined by ionic charges

▫ Electron(s) transferred from cation to anion

▫ Electrostatic in nature

Interactions between charged objects (ions)

Bond energy given by Coulomb’s Law

Coulomb’s Law

𝐸 = 2.31 𝑥10−19𝐽 ∙ 𝑛𝑚𝑄1𝑄𝑠𝑟

E = joules

r = distance between ion centers in nanometers

Q1 and Q2 = numerical ion charges

▫ Negative E indicates an attractive force (ion pair has lower energy than separated ions)

▫ Positive E indicates repulsive energy when two like-charged ions are brought together

▫ Small, highly charged ions form strong/favorable ionic compounds

Higher the charges, the greater the attractive energy

Smaller the distance, the greater the attractive energy

Ionic Bonds (Continued)

• Ionic compounds form huge, repeating 3-D crystalline lattices

▫ Ions and electrons are located at fixed positions

• Ionic bond strength is reflected in lattice energy

▫ Modified form of Coulomb’s Law

Ionic Bonds

• Defined as the energy released when gaseous ions react to form one MOLE of a solid ionic compound

• Most negative lattice energy occurs between large charges and small ionic radii ▫ Highly favorable

• Least negative lattice energy occurs between small charges and large ionic radii

▫ Less favorable

• NaCl < NaF < MgS < MgO

𝐿𝑎𝑡𝑡𝑖𝑐𝑒 𝐸𝑛𝑒𝑟𝑔𝑦 = 𝑘(𝑄`𝑄2𝑟)

k = proportionality constant that depends on structure of solid

• Strong interactions between ions have a profound effect on melting points and solubilities

▫ Large melting points ▫ Solids at room temperature

Covalent Bonds

• Bond between two non-metals atoms

▫ Valence electrons are shared between nuclei of bonding atoms

Sharing based on electronegativity of each atom in bond

• Bonds can be single, double, or triple as shown by Lewis structures

• Physical properties vary wildly

Electronegativity (En)

• The ability of an atom IN A MOLECULE (meaning it’s participating in a bond) to attract shared electrons to itself

▫ F is most electronegative Highest Zeff and smallest radius so that the nucleus is closed to the

“action” ▫ Fr is least electronegative Lowest Zeff and largest radius so that the nucleus is farthest from the

“action”

• Can use the difference in electronegativities to determine type of bond formed

▫ Ionic – electronegativity difference greater than 1.67 ▫ Covalent – electronegativity difference less than 1.67 ▫ Non-polar covalent – electronegativity differences less than 0.4

How Do These Bonds Form?

• Valence electrons ▫ Outermost electrons ▫ TOTAL highest energy s and p electrons Focus on ns, np, and d electrons of transition elements

▫ Most elements obey octet rule Each atom in a covalent bond has a TOTAL of 8 valence electrons

around it Most important requirement for the formation of a stable compound is that

atoms achieve a noble gas configuration (octet)

▫ EXCEPTIONS H – 2 electrons total Be – 4 electrons total B – 6 electrons total n = 3 and above – expanded octets from d orbitals NO, NO2, and ClO2 contain an odd number of valence electrons and

thus, cannot obey octet rule

Single and Multiple Bonds

Their Properties • Bond order is the number of bonding electron pairs shared by two atoms in a molecule • Single bond (bond order = 1)

▫ One pair of electrons shared Called a sigma bond

• Double bond (bond order = 2) ▫ Two pairs of electrons shared

One sigma and one pi bond

• Triple bond (bond order = 3) ▫ Three pairs of electrons shared

One sigma and two pi bonds

• Obviously, combinations of sigma and pi bonds are stronger than sigma alone ▫ Pi bonds are weaker than sigma but NEVER exist alone

• HUGE CONCEPUTAL NOTE ▫ Multiple bonds increase electron density between two nuclei

Decreases nuclear repulsions while enhancing the nucleus to electron density attractions

▫ Nuclei move closer together Bond lengths from shortest to longest are as follows:

Triple bond < Double bond < Single bond

Localized Electron Model

• Bonding theory used to describe covalent bonds • Assumes that electrons are localized on an atom or

the space between atoms ▫ Lone pair electrons ▫ Bonding pair electrons • Has 3 parts:

▫ Lewis Dot structure describe valence electron arrangement

▫ Geometry is predicted with VSEPR ▫ Description of the type of atomic orbitals “blended” by

the atoms to share electrons or hold lone pairs Hybrids – next chapter!

How to Illustrate Covalent Bonding -

Lewis Dot Structures • Illustrate valence electrons and subsequent bonding using Lewis Dot

structures

▫ Dots represent valence electrons

▫ Usually only done for main group elements

• Tips for drawing Lewis Dot structures:

▫ Determine total number of valence electrons

Add for anions, subtract for cations

▫ Predict # of bonds by counting the number of unpaired electrons in Lewis structure

▫ Least electronegative atom is the center atom

Remember the trend!

▫ Draw a single bond , -, (2 electrons) to each atom

▫ Subtract from total

▫ Add lone pair electrons, :, to terminal atoms to satisfy octet rule

Extras go to central atom

▫ If central atom is not octet, use terminal electrons to make double bond

Carbon bonded to N, O, P, S tend to form double bonds

▫ Hydrogen is ALWAYS a terminal atom

Only makes 1 bond

Practice!

• Give the Lewis structure for each of the following:

• HF

• N2

• NH3

• CH4

• CF4

• NO+

Resonance Structures

• Defined as a compound that has multiple equivalent structures

• A compound with resonance is best described as the average of all the equivalent structures

▫ Intermediate bond lengths

• Clarification of common misconceptions:

▫ Structures with resonance do not “flip” between equivalent structures

▫ Structures differ by placement of electrons, not atoms

▫ Octet rules still apply

▫ Not all compounds have resonance

Formal charges

• Defined as the charge assigned to an atom in a molecule assuming that electrons in a bond are shared equally between atoms

▫ Note – oxidation state assumes NO sharing

• Important for determining which resonance structure is most significant contributor

▫ All atoms have formal charge of 0 (most important) ▫ Charges are consistent with electronegativity Atoms that are higher in electronegativity have

negative formal charges

Determining Formal Charges

Atom FC= # valence electrons − [# of lone electrons − # of bonding groups ]

Practice!

• See homework

VSEPR Theory

Determining Molecular Geometries

• In order to predict molecular shape, we use the Valence Shell Electron Pair Repulsion (VSEPR) theory

• This theory proposes that the geometric arrangement of groups of atoms about a central atom in a covalent compound is determined solely by the repulsions between electron pairs present in the valence shell of the central atom ▫ The molecule adopts whichever 3-D geometry minimizes the repulsion between valence electrons

Determining Molecular Geometries

• To determine the shape of a molecule, we distinguish between: ▫ Lone pairs (non-bonding pairs) ▫ Bonding pairs (those found between two atoms)

Multiple bonds are considered as ONE bonding pair even though in reality, they have multiple pairs of electrons

• All electrons are considered when determining 3-D shape

AXmEn

A - central atom X – surrounding atom

E – non-bonding valence electron group m and n - integers

Electron Group Repulsions and the Five

Basic Molecular Shapes

Factors Affecting Electron Repulsion • Two factors that affect the amount of electron repulsion around an

atom: ▫ Multiple bonds

Exert a greater repulsive force on adjacent electron pairs than do single bonds

Result of higher electron density

Distorts basic geometry!

▫ Non-bonding (lone) pairs Lone pairs repel bonding pairs more strongly than bonding pairs repel

each other

The Effect of Non-Bonding Electrons

on Bond Angles • Remember, electron pairs of

bonding atoms are shared by two atoms, whereas the nonbonding electron pairs (lone pairs) are attracted to a single nucleus

▫ As a result, lone pairs can be thought of as having a somewhat larger electron cloud near the parent atom

• This “crowds” the bonding pairs and the geometry is distorted!

▫ Bond angles change!

Factors Affecting Bond Angles

Double Bonds Non-Bonding (Lone) Pairs

Time to Explore the Different Molecule

Arrangements!

VSEPR Simulation

The Single Molecular Shape of Linear

Electron-Group Arrangement • AX2

• Examples

▫ CS2, HCN, BeF2

A X X

The 2 Molecular Shapes of Trigonal

Planar Electron-Group Arrangement Trigonal Planar Bent

• AX3

• Examples

▫ SO3, BF3, NO3-, CO3

2-

• AX2E

• Examples

▫ SO2, O3, PbCl2, SnBr2

A

X

X

X X

X A

E

The 3 Molecular Shapes of the Tetrahedral

Electron-Group Arrangement Tetrahedral

• AX4

• Examples

▫ CH4, SiCl4, SO4

2-, ClO4-

• AX3E

• Examples

▫ NH3, PF3, ClO3, H3O+

Trigonal Pyramidal Bent

• AX2E2

• Examples

▫ H2O, OF2, SCl2

A

X

X X

X A

X X

X

E

A

E

E X

X

The 4 Molecular Shapes of the Trigonal

Bipyramidal Electron-Group Arrangement

• AX5

• Examples

▫ PCl5, PF5, AsF5, SOF4

Trigonal Bipyramidal

See-Saw T-Shaped Linear

• AX4E

• Examples ▫ SF4, XeO2F2,

IF4+, IO2F2-

• AX3E2

• Examples

▫ ClF3, BrF3

• AX2E3

• Examples

▫ XeF2, I3-, IF2-

The 3 Molecular Shapes of the Octahedral

Electron-Group Arrangement Octahedral

• AX6

• Examples

▫ SF6, IOF5

• AX5E

• Examples

▫ BrF5, XeOF4, TeF5-

Square Pyramidal Square Planar

• AX4E2

• Examples

▫ XeF4, ICl4-

Steps in Determining a Molecular

Shape • Write the Lewis structure • Determine the electron-group arrangement, ideal

bond angles, and VSEPR class • Place the surrounding atoms and lone pairs in

appropriate positions around the central atom and predict any deviations from the ideal bond angles

• Name the molecular shape • For molecules with MORE THAN ONE CENTRAL

ATOM, find the electron-group arrangement and corresponding shape around EACH central atom ▫ One central atom at a time

• The ability of an atom IN A MOLECULE (meaning it’s participating in a bond) to attract shared electrons to itself

• Can use the difference in electronegativities to determine type of bond formed

▫ Ionic – electronegativity difference greater than 1.67

▫ Covalent – electronegativity difference less than 1.67

▫ Non-polar covalent – electronegativity differences less than 0.4

Electronegativity (En)

• Electronegativities determine polarity since it measures a nucleus’ attraction or “pull” on the bonded electron pair

▫ When two nuclei are the same, sharing is equal Non-polar

▫ When 2 nuclei are different, the electrons are not shared equally

Polar ▫ When electrons are shared unequally to a greater extent,

IONIC • Bonds can be polar while the entire molecule is not

▫ Determined by geometry! More on this later!

• Dipole moment ▫ Separation of the charge in a molecule (slightly

positive/slightly negative poles) ▫ IF octet rule is obeyed AND all the surrounding bonds are the

same (even if they’re very polar), then the molecule is NONPOLAR

Example: CCl4

Molecular Polarity

VSEPR and Polarity

• Knowing the geometry of a molecule allows one to predict whether it is polar or nonpolar ▫ A bond between unlike atoms is usually polar with a

positive end and a negative end • The symmetry of the molecule determines polarity

▫ A diatomic molecule containing two different atoms is polar HF, CO

▫ A diatomic molecule containing the same two atoms is nonpolar N2, O2

▫ A polyatomic molecule may be nonpolar even if it contains polar bonds because, in such cases, the polar bonds are counteracting each other CO2, CH4 = nonpolar

VSEPR Symmetry and Molecular

Polarity

Bond Energies • Again, the greater the number of electron pairs between a pair of atoms, the

shorter the bond ▫ This implies that atoms are held together more tightly when there are multiple

bonds ▫ There IS a relation between bond order and the energy required to separate them

• In order for bonds to be broke, energy must be added to the system (endothermic reaction)

• In order for bonds to be formed, energy must be released from the system (exothermic reaction)

▫ Enthalpy change for a reaction is the sum of the energies required to break old bonds plus the sum of the energies released in the formation of new bonds

∆𝐻 = 𝐷 𝑏𝑜𝑛𝑑𝑠 𝑏𝑟𝑜𝑘𝑒𝑛 − 𝐷 (𝑏𝑜𝑛𝑑𝑠 𝑓𝑜𝑟𝑚𝑒𝑑)

D represents the bond energy per mole of bonds (ALWAYS POSITIVE)