Chemical Bonding Chemical Bonds A bond is a force that holds groups of two or more atoms together and makes them function as a unit. A bond is a force that holds groups of two or more atoms together and makes them function as a unit. Bond energy is the amount of energy required to break a particular bond. Bond energy is the amount of energy required to break a particular bond. *determines the strength of a bond *determines the strength of a bond Bonding Review What type of elements form ionic bonds? What type of elements form ionic bonds? What type of elements form covalent bonds? What type of elements form covalent bonds? Where are these types of elements found on the periodic table in relationship to each other? Where are these types of elements found on the periodic table in relationship to each other? Chemical Bonding: An Overview Valence electrons play a fundamental role in chemical bonding. When metals and nonmetals combine, valence electrons usually are transferred from the metal to the nonmetal atoms, giving rise to ionic bonds. In combinations involving only nonmetals, one or more pairs of valence electrons are shared between the bonded atoms, producing covalent bonds. In losing, gaining, or sharing electrons to form chemical bonds, atoms tend to acquire the electron configurations of noble gases. Types of Chemical Bonds: Ionic Bonding Ionic bonding occurs when there is an attraction between oppositely charges ions. Electrons are transferred to form oppositely charged ions. Ionic bonding occurs when there is an attraction between oppositely charges ions. Electrons are transferred to form oppositely charged ions. Ionic compounds result when an atom that loses electrons relatively easily (a metal) reacts with an atom that has a high affinity for electrons (a non-metal). Ionic compounds result when an atom that loses electrons relatively easily (a metal) reacts with an atom that has a high affinity for electrons (a non-metal). Example: Sodium chloride, NaCl Example: Sodium chloride, NaCl Na + & Cl - ions react to form solid sodium chloride. Types of Chemical Bonds: Covalent Bonding Covalent Bond: A bond that results from the sharing of electrons between atoms. Non-polar covalent bonding occurs between identical atoms, like H 2, in which the electrons are shared equally. Types of Chemical Bonds: Polar Covalent Bonding Polar covalent bonding occurs when atoms are not different enough to transfer electrons (ionic) but at the same time are not identical and cannot share electrons equally (covalent); instead there is an unequal sharing of electrons. Polar covalent bonding occurs when atoms are not different enough to transfer electrons (ionic) but at the same time are not identical and cannot share electrons equally (covalent); instead there is an unequal sharing of electrons. Example: HF, which has an unequal charge distribution or a partial charge. Electrons are more attracted to, or want to spend more time with, fluorine than hydrogen. This is what we call bond polarity. Example: HF, which has an unequal charge distribution or a partial charge. Electrons are more attracted to, or want to spend more time with, fluorine than hydrogen. This is what we call bond polarity. The partial charge is represented with a delta, , and a positive or negative sign depending on where electrons are positioned. The partial charge is represented with a delta, , and a positive or negative sign depending on where electrons are positioned. Bond Polarity & Dipole Moments A dipole moment indicates the direction of charge distribution between two atoms. Can be represented by a center of positive charge and a center of negative charge. A dipole moment indicates the direction of charge distribution between two atoms. Can be represented by a center of positive charge and a center of negative charge. The dipole character of a molecule is often represented by an arrow, which points toward the negative center and its tail indicated the positive center. The dipole character of a molecule is often represented by an arrow, which points toward the negative center and its tail indicated the positive center. Any diatomic molecule that has a polar bond has a dipole moment. Any diatomic molecule that has a polar bond has a dipole moment. Bond Polarity NaCl HCl Cl 2 Ionic and covalent bonding are the extremes of types of bonding. Ionic > Polar Covalent > Non-Polar Covalent Electronegativity Electronegativity is the relative ability of an atom in a molecule to attract shared electrons to itself. Electronegativity is the relative ability of an atom in a molecule to attract shared electrons to itself. Trend: Electronegativity increases across the period and decreases down the group. Trend: Electronegativity increases across the period and decreases down the group. Fluorine is the most electronegative atom. Fluorine is the most electronegative atom. Noble Gasses have NO electronegativity! Noble Gasses have NO electronegativity! Electronegativity & Polarity The polarity of a bond depends on the difference between the electronegativity values of the atoms forming the bond. The polarity of a bond depends on the difference between the electronegativity values of the atoms forming the bond. Difference of Zero Non-polar Covalent Bond Intermediate Difference (greater than 1.0 to 2.0) Polar Covalent Bond Large Difference (greater than 2.0) Ionic Bond Lewis Symbols (Electron Dot Symbols) In a Lewis symbol, the chemical symbol for the element represents the nucleus and core electrons of the atom. Dots around the symbol represent the valence electrons. In writing Lewis symbols, the first four dots are placed singly on each of the four sides of the chemical symbol. Dots are paired as the next four are added. Lewis symbols are used primarily for those elements that acquire noble-gas configurations when they form bonds. Lewis Structures of Simple Molecules A Lewis structure is a combination of Lewis symbols that represents the formation of covalent bonds between atoms. Lewis structure shows the bonded atoms with the electron configuration of a noble gas; that is, the atoms obey the octet rule. Lewis Structures The shared pairs of electrons in a molecule are called bonding pairs. In common practice, the bonding pair is represented by a dash (). The other electron pairs, which are not shared, are called nonbonding pairs, or lone pairs. Each chlorine atom sees an octet of electrons. Multiple Covalent Bonds The covalent bond in which one pair of electrons is shared is called a single bond. Multiple bonds can also form: In a double bond two pairs of electrons are shared. Note that each atom obeys the octet rule, even with multiple bonds. In a triple bond three pairs of electrons are shared. Writing Lewis Structures: A Method 1.Add together the number of Happy electrons. (H/He = 2e-, B = 6e-, all other atoms = 8e-) 2.Add up the number of valence electrons. 3.(Step #1 Step #2) / 2 = # of Bonds 4.Put the least electronegative atom in the middle (called the central atom) and attach all other atoms as terminal atoms. # bonds = # attachments all single bonds # bonds > # attachments multiple bonds (double, triple) # bonds < # attachments EXCEPTION to the octet rule (attach all attachments with single bonds) 5. Count electrons to make sure all atoms are happy. (H/He = 2e-, B = 6e-, all other atoms = 8e-) If short add the correct number of lone pairs to make each atom Happy **If the compound is an exception to the octet rule, make attachments happy and place extra valence electrons in pairs on the central atom may hold more than 8 electrons. Example Write a plausible Lewis structure for nitrogen trifloride, NF 3. Example Write a plausible Lewis structure for carbonate ion, C Example Write the Lewis structure for boron tribromide, BBr 3. Exceptions to the Octet Rule Exceptions to the Octet Rule Expanded octets (more than 8 electrons) can only occur in elements that have d-orbitals (periods 3-7). Expanded octets (more than 8 electrons) can only occur in elements that have d-orbitals (periods 3-7). 1. Do the math # bonds < # attachments 2. Add all attachments -- single bonds and lone pairs to make happy. 3. Count number of electrons in structure and compare with valence electrons (step 2). 4. If short valence electrons, add them in pairs to the central atom only. Example Write the Lewis structure for iodine pentafluoride, IF 5. Resonance: Delocalized Bonding When a molecule or ion can be represented by two or more plausible Lewis structures that differ only in the distribution of electrons, the true structure is a composite, or hybrid, of them. The different plausible structures are called resonance structures. The actual molecule or ion that is a hybrid of the resonance structures is called a resonance hybrid. Electrons that are part of the resonance hybrid are spread out over several atoms and are referred to as being delocalized. Three pairs of electrons are distributed among two bonds. Example Write three equivalent Lewis structures for the SeO 3 molecule that conform to the octet rule, and describe how the resonance hybrid is related to the three structures. Molecular geometry is simply the shape of a molecule. Molecular geometry is simply the shape of a molecule. Molecular geometry is described by the geometric figure formed when the atomic nuclei are joined by (imaginary) straight lines. Molecular geometry is described by the geometric figure formed when the atomic nuclei are joined by (imaginary) straight lines. Molecular geometry is found using the Lewis structure, but the Lewis structure itself does NOT necessarily represent the molecule s shape. Molecular geometry is found using the Lewis structure, but the Lewis structure itself does NOT necessarily represent the molecule s shape. A water molecule is angular or bent. A carbon dioxide molecule is linear. Molecular Geometry Valence-Shell Electron-Pair Repulsion (VSEPR) is a simple method for determining geometry. Valence-Shell Electron-Pair Repulsion (VSEPR) is a simple method for determining geometry. Basis: pairs of valence electrons in bonded atoms repel one another. Basis: pairs of valence electrons in bonded atoms repel one another. These mutual repulsions push electron pairs as far from one another as possible. These mutual repulsions push electron pairs as far from one another as possible. B A B B B B A When the electron pairs (bonds) are as far apart as they can get, what will be the B-A-B angle? VSEPR Two Attachments Linear Three Attachments Trigonal Planar Four Attachments Tetrahedral Five Attachments Trigonal Bipyramidal Six Attachments Octahedral Example Use the VSEPR method to predict the shape of the nitrate ion, NO 3 -. Example Use the VSEPR method to predict the shape of the SF 4 Molecular Polarity Although most bonds within a molecule are polar, the arrangement of these bonds may cause the overall molecule to be non-polar. Although most bonds within a molecule are polar, the arrangement of these bonds may cause the overall molecule to be non-polar. IDENTICAL dipoles (same elements b/w bonds) can cancel at specific angles or combination of the following. 180 o IDENTICAL dipoles (same elements b/w bonds) can cancel at specific angles or combination of the following. 180 o 120 o 120 o o o HINT: if the molecule contains less than ( Non-Polar Covalent Electronegativity _________________ is the relative ability of an atom in a molecule to attract shared electrons to itself. _________________ is the relative ability of an atom in a molecule to attract shared electrons to itself. Trend: ___________________________________________________________________ Trend: ___________________________________________________________________ ___________ is the most electronegative atom. ___________ is the most electronegative atom. _______________ have NO electronegativity! _______________ have NO electronegativity! Electronegativity & Polarity The polarity of a bond depends on the difference between the electronegativity values of the atoms forming the bond. The polarity of a bond depends on the difference between the electronegativity values of the atoms forming the bond. Difference of Zero Non-polar Covalent Bond Intermediate Difference (greater than zero to 2.0) Polar Covalent Bond Large Difference (greater than 2.0) Ionic Bond