Advanced Chemistry

Ms. Grobsky

Chemical Bonding

• 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

What is Bonding? Why do Atoms 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

Ionic Bonds

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

• Ions and electrons are located at fixed positions

• Strong interactions between ions

• Large melting points

• Solids at room temperature

Ionic Bonds (Continued)

• Bond between two non-metals atoms

• Valence electrons are shared between nuclei of bonding atoms• When shared equally, bond is called non-polar covalent

• When shared unequally, bond is called polar covalent and dipoles are established

• 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

Covalent Bonds

• Sharing of valence electrons• Electrons in the highest occupied energy shell of the atom

• TOTAL highest energy s and p electrons

• Focus on ns, np, and d electrons of transition elements

How Do Covalent Bonds Form?

• Single bond • One pair of electrons shared

• Double bond• Two pairs of electrons shared

• Triple bond• Three pairs of electrons shared

Single and Multiple Bonds

• 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

• The shorter the bond implies that atoms are held together more tightly when there are multiple bonds

• Multiple bonds are stronger than single bonds

Multiple Bonds and Bond Lengths

• Called the Localized Electron Model• Used to describe covalent bonds

• Assumes that electrons are localized (restricted to certain areas) on an atom or the space between atoms• Lone pair electrons

• Bonding pair electrons

• You will learn about 2 parts of the model:• Lewis Dot structure describe valence electron arrangement

• Geometry is predicted with VSEPR

How Do We Describe the Structure of Covalent Bonds?

• Lewis Dot structures are also known as electron dot diagrams

• These diagrams show only the valence (bonding) electrons• Unpaired (single) electrons will participate in bonding

• Paired electrons will not participate in bonding

• Octet Rule• 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)

• There are EXCEPTIONS to this rule!• 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

Lewis Dot Structures

• Determine total number of valence electrons

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

Steps to Draw Lewis Dot Diagrams for Elements

• Determine total number of valence electrons• Add them up for BOTH compounds!

• 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 and extra electrons are left unpaired, form multiple bonds!• Carbon bonded to N, O, P, S tend to form double bonds

• Hydrogen is ALWAYS a terminal atom• Only makes 1 bond

Steps to Draw Lewis Dot Structures for Compounds

• Ionic Lewis Dot structures are drawn exactly the same way as covalent compounds• ONE EXCEPTION – Ionic compounds only form SINGLE bonds!

• Metal donates all valence electrons to non-metal

Ionic Compounds and Lewis Dot Structures

• Sometimes, an atom is unable to form a stable compound by following the octet rule

• Some atoms can make compounds using paired electrons in their inner shell (d and f-orbitals)

• This causes expanded octets• Create more bonds than expected

• Example: BrF3 and PCl5

Expanded Octets and Lewis Dot Structures

• Some covalently bonded atoms can have a few extra or fewer electrons, resulting in an overall charge• Negative charge (anions) – additional electrons must be added

• Positive charge (cations) – electrons need to be reduced (subtract)

• Examples: NH4+ and SO42-

Polyatomic Ions and Lewis Dot Structures