Advanced Chemistry
Ms. Grobsky
What is Bonding?
Why do Atoms Bond?
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Bonding is the interplay between interactions between atoms
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Energetically favored
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Electrons on one atom interacting with protons of another atom
Energetically unfavorable
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Electrons on one atom interacting with electrons of another atom
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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
Ionic Bonds
• Bond between a metal cation and non-metal anion
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Formula determined by ionic charges
• Electron(s) transferred from cation to anion
• Electrostatic in nature
Ionic Bonds (Continued)
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Ionic compounds form huge, repeating 3-D crystalline lattices
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Ions and electrons are located at fixed positions
Strong interactions between ions
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Large melting points
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Solids at room temperature
Covalent Bonds
• Bond between two non-metals atoms
• Valence electrons are shared between nuclei of bonding atoms
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When shared equally, bond is called non-polar covalent
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When shared unequally, bond is called polar covalent and dipoles are established
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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
How Do Covalent Bonds Form?
• Sharing of valence electrons
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Electrons in the highest occupied energy shell of the atom
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TOTAL highest energy s and p electrons
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Focus on ns, np, and d electrons of transition elements
Single and Multiple Bonds
• Single bond
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One pair of electrons shared
• Double bond
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Two pairs of electrons shared
• Triple bond
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Three pairs of electrons shared
Multiple Bonds and Bond Lengths
• Multiple bonds increase electron density between two nuclei
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Decreases nuclear repulsions while enhancing the nucleus to electron density
attractions
• Nuclei move closer together
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Bond lengths from shortest to longest are as follows:
Triple bond < Double bond < Single bond
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The shorter the bond implies that atoms are held together more tightly when there
are multiple bonds
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Multiple bonds are stronger than single bonds
How Do We Describe the Structure of
Covalent Bonds?
• Called the Localized Electron Model
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Used to describe covalent bonds
• Assumes that electrons are localized (restricted to certain areas) on an
atom or the space between atoms
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Lone pair electrons
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Bonding pair electrons
• You will learn about 2 parts of the model:
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Lewis Dot structure describe valence electron arrangement
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Geometry is predicted with VSEPR
Lewis Dot Structures
• Lewis Dot structures are also known as electron dot diagrams
• These diagrams show only the valence (bonding) electrons
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Unpaired (single) electrons will participate in bonding
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Paired electrons will not participate in bonding
• Octet Rule
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Most elements obey octet rule
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Each atom in a covalent bond has a TOTAL of 8 valence electrons around it
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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!
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H – 2 electrons total
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Be – 4 electrons total
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B – 6 electrons total
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n = 3 and above – expanded octets from d orbitals
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NO, NO2, and ClO2 contain an odd number of valence electrons and thus, cannot
obey octet rule
Steps to Draw Lewis Dot Diagrams for
Elements
• Determine total number of valence electrons
• Predict # of bonds by counting the number of unpaired electrons in Lewis structure
Steps to Draw Lewis Dot Structures for
Compounds
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Determine total number of valence electrons
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Add them up for BOTH compounds!
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Add for anions, subtract for cations
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Predict # of bonds by counting the number of unpaired electrons in Lewis structure
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Least electronegative atom is the center atom
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Remember the trend!
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Draw a single bond , -, (2 electrons) to each atom
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Subtract from total
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Add lone pair electrons, :, to terminal atoms to satisfy octet rule
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Extras go to central atom
If central atom is not octet and extra electrons are left unpaired, form multiple bonds!
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Carbon bonded to N, O, P, S tend to form double bonds
Hydrogen is ALWAYS a terminal atom
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Only makes 1 bond
Ionic Compounds and Lewis Dot
Structures
• Ionic Lewis Dot structures are drawn exactly the same way as covalent
compounds
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ONE EXCEPTION – Ionic compounds only form SINGLE bonds!
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Metal donates all valence electrons to non-metal
Expanded Octets 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
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Create more bonds than expected
• Example: BrF3 and PCl5
Polyatomic Ions and Lewis Dot
Structures
• Some covalently bonded atoms can have a few extra or fewer electrons,
resulting in an overall charge
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Negative charge (anions) – additional electrons must be added
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Positive charge (cations) – electrons need to be reduced (subtract)
• Examples: NH4+ and SO42-