Upload
duongminh
View
214
Download
0
Embed Size (px)
Citation preview
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
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
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 ]
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
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!
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
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)