Presentation on theme: "Chapter 8 Covalent Bonding. The Covalent Bond Atoms will share electrons in order to form a stable octet. l Covalent bond : the chemical bond that results."— Presentation transcript:
Bonding pair: a pair of electrons shared by two atoms Lone pair: an unshared pair of electrons on an atom
l When a single pair of electrons is shared, a single covalent bond forms. l This can be represented with a Lewis structure: uses electron dot diagrams to show how electrons are arranged in molecules.
l Group 17 elements will form one covalent bond. l Group 16 elements will form two covalent bonds. l Group 15 elements will form three covalent bonds. l Group 14 elements will form four covalent bonds.
l Multiple bonds are made up of sigma bonds and pi bonds: formed when parallel orbitals share electrons. l A double covalent bonds has one sigma and one pi bond. l A triple covalent bond has one sigma and two pi bonds.
Strength of Covalent Bonds l The strength of covalent bonds is determined by the bond length: distance between the bond nuclei l Bond length is determined by: l The size of the atoms involved—larger atoms have longer bond lengths l How many pairs of electrons are shared—the more pairs of electrons shared, the shorter the bond length is.
l The strength of a covalent bond is indicated by its bond dissociation energy: the amount of energy required to break the bond. l When a bond forms, energy is released; when a bond breaks, energy must be added. l Each covalent bond has a specific value for its bond dissociation energy. These values are determined experimentally.
l A direct relationship exists between bond energy and bond length: l The total energy change of a chemical reaction is determined from the energy of the bonds broken and formed in the reaction. l Endothermic reactions: have a net absorption of energy l Exothermic reactions: have a net release of energy
How do we show the arrangement of atoms in a molecule? l By using Lewis structures:
Five steps to draw Lewis structures: l Count the total number of valence electrons in all atoms involved. l Write a skeleton equation, subtracting the # of electrons used to show bonds. l Fill the octet of the terminal atoms. Subtract electrons. l Place any remaining electrons around the central atom to satisfy its octet. If the central atom cannot be satisfied, make a multiple bond using a lone pair from the terminal atoms. l Check your work
Hints for skeleton equation: l Hydrogen is always an end atom l The atom with the least attraction for shared electrons is the central atom. This element is usually the one closer to the left on the periodic table. l Place lone pairs around each outer atom to satisfy their octets. Subtract the number of electrons used.
Drawing Lewis structures for polyatomic ions is very similar to drawing Lewis structures for covalent compounds EXCEPT in finding the number of electrons available for bonding l Count the total number of valence electrons in all atoms involved. l If the polyatomic ion is negatively charged, ADD the charge to the number of valence electrons. l If the ion is positively charged, SUBRACT the charge from the number of valence electrons. l Follow the rest of the steps to drawing Lewis structures.
Resonance Structures l When a molecule or polyatomic ion has both a double bond and a single bond, it is possible to have more than one correct Lewis structure:
l Resonance: a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion. The structures are called resonance structures. l A molecule that undergoes resonance behaves as if it has only one structure.
Exceptions to the Octet Rule Always on Central Atom Three Ways Molecules Might Violate the Octet Rule 1. Odd Number of Valence Electrons –Some molecules have an odd number of valence electrons and cannot form an octet around each atom
Sub Octet… Less than 8 l Some compounds form with fewer than 8 electrons present around an atom. l Boron does this. –Coordinate covalent bond: when one atom donates an entire pair of electrons to be shared with atoms or ions that need two more electrons.
Expanded Octet l Some atoms can have expanded octets: have more than eight electrons in their valence shell.
l How? –The d orbital starts to hold electrons. –This occurs in atoms in Period 3 or higher. –When you draw Lewis structures for these compounds, extra lone pairs are added to the central atom OR the central atom will form more than four bonds.
Molecular Shape l When we talk about how a molecule “looks” we talk about three characteristics: 1.Shape 2.Bond Angles 3.Hybridization
VSEPR Model l The shape of a molecule determines whether or not two molecules can get close enough to react. l We describe shape using the VSEPR model.
This model is based on the fact that electrons pairs will stay as far away from each other as possible Valence Shell Electron Pair Repulsion
l Atoms will assume certain bond angles: the angle formed by any two terminal atoms and the central atom l Lone pairs take up more space than bonded pairs do.
Model Building l Draw the Lewis Structure. l Count the Bonded Pairs of Electrons around the central atom. l Count the unbonded pairs of electrons. l Look up the amounts on the VSEPR chart. l Classify the molecular geometry (shape). l Build the molecule.
Hybridization l During bonding, atomic orbitals can undergo hybridization: a process in which atomic orbitals are mixed to form new, identical hybrid orbitals. l Hybridization occurs on the central atom. l Carbon is the most common element that undergoes hybridization. l Carbon has the electron configuration 1s 2 2s 2 2p 2 l S and p orbitals are shaped very differently. If 2 e - s are in the s orbital and 2 e - s are in the p orbitals, why are all 4 bonds the same in the molecule CH 4 ? l One s orbital and three p orbitals combine to form 4 sp3 orbitals. Each one of these orbitals has one valence electron.
l Determine the molecular geometry, bond angle, and type of hybridization for the following: l BF 3 l NH 4 + l OCl 2 l BeF 2 l CF 4
l Determining Bond Type l Electronegativity and Polarity l Classifying Molecules as polar or nonpolar
Electronegativity and Polarity l Electron affinity is the tendency of an atom to accept an electron. l The scale of electronegativities allows chemists to predict electron affinities of atom in compounds. l Upper right high electronegativity. l Lower left low electronegativity
Electronegativity Difference l Determines the type of bond formed between two atoms. l Identical elements have equal electronegativities, no difference… the bond is said to be nonpolar covalent (a pure covalent bond). l Unequal sharing of electrons results in a polar covalent bond. l Very large differences in electronegativity result in electron transfer ( ionic bonding ).
What does Polar mean? l Unequal sharing of electrons l One atom pulls the electrons to itself more of the time. l Partial charges occur at the ends of the bond. l More electronegative atom gets the partially negative charge and l The lower electronegative atom gets the partial positive charge. l Called a Dipole.
Classify the bond l Ionic Bond: electronegativity difference of greater than 1.7 l Polar Covalent Bond: e.d. of 0.3-1.7 l Nonpolar Covalent Bond: e.d. less than 0.3
Molecular Polarity l Molecules are either polar or nonpolar depending on the bonds in the molecule. l We must look at the shape (geometry) of a molecule to determine polarity. l Symmetric molecules are nonpolar.
Solubility of Polar Molecules l Bond type and shape of the molecule determine solubility. l Polar substances and ionic substances will dissolve in polar solvents. l Nonpolar substances will only dissolve in nonpolar substances.
Molecules Polar l Opposite ends l + end - end l Asymmetrical l Stronger bonds l Unequal distribution of electron pairs l Higher boiling pts l Higher melting points l Lower evaporation rates NonPolar l Ends are the same l No charged ends l Symmetrical l Weaker Intermolecular attractions l Temporary dipoles between molecules l Bonding is weak l Low melting and boiling points high evaporation rates