Chemical Bond Mutual attraction between the nuclei and valence electrons of different atoms that binds them together. Bonding onors Chemistry 412 Chapter 6 Types of Bonds Ionic Bonds Force of attraction between oppositely charged ions. Covalent Bond Force of attraction for electrons, that results in a pair of electrons being shared by two atoms. Ionic or Covalent? Polar Vs. Nonpolar Covalent Bonds The Difference in electronegativity determines the bond type, page 161, according to the following scale: Example Na-Cl Na = 0.9 Cl = 3.0 Difference = 3.0 0.9 = 2.1 Ionic Non-Polar electrons are equally shared Balanced Distribution of Electrical Charge Polar electrons are not equally shared Unbalanced Distribution of Electrical Charge Electronegativity difference % ionic character 3.3 100% Ionic Polar Covalent Non Polar 1.7 50% 0.3 5% 0 0% + - Identifying Bond Types Determine if the following are going to form Ionic, Polar Covalent or Non-polar Covalent Bonds: C-N Ca-F Br-Br -Br Solutions + C-N - (.5 Polar) + Ca-F - (3.0 Ionic) Br-Br (0 Non polar) + -Br - (.7 Polar) 1
Example: 2 O 1. : 1 each, O: 6 each, Total: 8 valence electrons 2. Draw the skeleton: Summary of Bond Types -O - Duet Rule: ydrogen is satisfied! 3. Satisfy the central octet. 4. Satisty all other octets. 5. Check your work! A. All atoms have octets/duets B. Correct number of valence electrons was used Why do atoms bond? Atoms want to gain stability, like the noble gases Octet Rule Atoms tend to lose, gain, or share electrons in order to attain a full set of 8 valence electrons. Like a noble gas! Duet Rule ydrogen and elium only A full set consists of 2 valence electrons First principal energy level can only accommodate 2 electrons Covalent Bonds Force of attraction between atoms that share electrons Covalent Compound eld together by covalent bonds Usually contains two non-metals The smallest piece is called a molecule Properties Low Boiling/Melting points (weaker bonds) Most are gases at room temperature Non-conductors of electricity Formation of Covalent Bonds Bond results from the attraction forces between the nuclei and electrons of the two atoms Attraction Forces - Potential Energy Repulsion Forces - Potential Energy If the attraction forces overcome the repulsion forces the bond will form. (AF > RF) Bond Length the avg. distance between two atoms Bond Energy energy needed to break the bond Overall: Attractive Forces -BL - BE 2
Lewis Dot Structures We can use Lewis dots to show covalent bonds: Multiple Covalent Bonds 2 : F 2 : F F Single Covalent Bond Unshared pair or lone pair (not involved in bonding) Double Bond 2 pairs of shared electrons: (=) Triple Bond 3 pairs of shared electrons: ( ) We can use a dash (-) for bonds: -; F - F Each dash stands for 2 electrons! F 1s 2s 2p F 1s 2s 2p N 1s 2s 2p N 1s 2s 2p Bonding Electron Pair Triple Bond 3 shared pairs Steps to Drawing a Lewis Structure: 1. Determine the total number of electrons contained in the compound. 2. Draw the skeleton structure of the compound A. If there are more than one type of atom, the least electronegative is usually the center atom B. If there are multiples of the least electronegative, they will split the center C. Attach all other atoms to the center as symmetrically as possible 3. Satisfy the octet of the central atom A. Use single bonds to attach other atoms B. Fill in with unshared pairs as needed 4. Satisfy the octets of all other atoms following the same procedure 5. Check your work Try some! Write the Lewis Structure for the following formulas: N 3 -N - Practice Single Covalent Bonds Write the Lewis Structure for the following formulas: C 2 6 C 3 7 Cl SCl 2 Cl -S -Cl 3
Multiple Covalent Bonds What if I need to use too many electrons to get my Lewis Structure to work? Indicates you need to make a multiple bond! Rule of thumb: For every 2 electrons over the limit, you need to make ONE more bond! C 2 O: O C Too many electrons! You were only allowed to use 12, but there are 14 here! Solution is to create a Double Bond! Multiple Covalent Bonds CN: C N Still too many need to make another bond! Used 14 electrons, but you only have 10 available! You ll need to make multiple bonds! Polyatomic Ion Structures N 4 +1 Polyatomic ions are covalently bonded atoms that form a charge due to the gain or loss of electrons. When drawing the structures of polyatomic ions: ( ) ions must have extra electrons in the structure that is equivalent to its charge. (+) ions must lose the number of electrons in the structure that is equivalent to its charge. Place the structure in brackets and indicate the charge at the top right hand corner Nitrogen: 5 valence electrons ydrogen: 1 valence electron each +1 charge: Take away one valence electron Total: 8 valence electrons 1+ N For a polyatomic ion, always indicate the charge by putting the structure into brackets and writing the charge at the top right hand corner! Exceptions to the Octet Rule Incomplete Octet Atoms that can be satisfied with less than eight electrons. Boron can be satisfied with only six electrons. Expanded Octet Atoms that can have more than eight electrons. (Not ALWAYS) Common elements that accommodate expanded octets: Cl, Br, I, S, P and all noble gases. IMPORTANT!!! Only one atom may exceed the standard eight valence. Must be the central atom only!! Coordinate Covalent Bond Like a single covalent bond, but a single atom is sharing 2 electrons with another atom Does not increase the # of electrons that the atom being shared with owns This is determined by one atom having more electrons drawn than it originally contained. Indicated by using an arrow rather than a dash. ( ) The arrow must point away from the atom that owns the electrons being shared 4
Resonance Bonding in molecules or ions that cannot be correctly represented by a single Lewis Structure (in other words, can be drawn multiple ways) Molecules constantly resonate, or alternate between structures Shown by using a double headed arrow between resonance structures: CO 2 : Ionic Bonds The result of the attraction between cations and anions Remember: Cation: Positive ion, loses electrons Anion: Negative ion, gains electrons In an ionic bond, electrons are exchanged between atoms O = C = O O C - O Ionic Compounds Stronger than covalent bonds Usually a metal with a non-metal Neutral overall charges must balance Formula Unit Lowest whole number ratio of atoms in an ionic compound Properties igh Melting Points (bonds are strong need a lot of energy to break) Brittle Dissolve in water Are good conductors in the liquid phase (melted) 3D Array of forces is created (Lattice) Strength of Array varies w/ sizes, charges, & number of ions. Arrangement of ions gives the crystal its strength CaF 2 Lattice Energy The energy released when one mole of an ionic compound is formed from gaseous ions Can be used to compare the strength of bonds (LE, Bond Strength ) Negative values indicate that energy is released Examples: Compound Lattice Energy (kj/mol) NaCl -787.5 NaBr -751.4 CaF2-2634.7 LiCl -861.3 LiF -1032 MgO -3760 KCl -715 Ionic Bond Formation We can use Lewis dot structures to show the exchange of valence electrons in atoms as they form ionic bonds: Ca + Cl Na + S Ca Cl Cl CaCl 2 Formula Unit of the Compound 5
Practice: Show the electron transfer for the following elements and write the formula unit for the binary ionic compound: 1. Mg + F 2. K + P 3. Ba + Se Metallic Bonding Bonding resulting from the attraction of positive ions and mobile electrons Only occurs in Metallic Atoms Small Number of Valence Electrons Low Ionization Energy & Electronegativity Easily give up electrons At best weakly covalent Electrons are Delocalized Electrons do not belong to any one ion Metallic Properties Electron-Sea Model Explains Properties of Metals Lustrous, Good Conductors of eat & Electricity Malleability and Ductility Strength of metallic bonds are reflected in the metal s enthalpy of vaporization values Amount of energy absorbed as heat when a specified amount of substance vaporizes at constant pressure igher energy = stronger bonds VSEPR Valence Shell Electron Pair Repulsion States that all atoms/electrons attached to a central atom repel one another, and will situate themselves as far away from each other as possible Substituent anything bonded to a central atom Can be another atom OR an unshared pair of electrons VSEPR Orientation (aka Basic arrangement of substituents around the central atom Depends ONLY upon how many substituents there are Shape What the molecule will actually look like in three dimensions Depends upon how many AND what type of substituents there are Unshared pairs of electrons are invisible! 6
Using VSEPR to predict shapes For generic formulas: A = Central Atom B = Atom bonded to the central atom E = Unshared pair of electrons on the central atom **We are only concerned with unshared pairs that belong to the central atom!!** Molecules with One Substituent AB Linear Linear **Note** You can add up to 5 unshared pairs of electrons to this, and it will NOT affect the shape of the molecule! Molecules with Two Substituents AB2 Linear Linear Molecules with Three Substituents AB3 Trigonal Planar Trigonal Planar AB2E Trigonal Planar Angular/Bent Molecules with Four Substituents AB4 Tetrahedral Tetrahedral AB3E Tetrahedral Trigonal Pyramidal AB2E2 Tetrahedral Angular/Bent Molecules with Five Substituents AB4E Trigonal Bipyramidal See-Saw AB5 Trigonal Bipyramidal Trigonal Bipyramidal 7
Molecules with Five Substituents AB3E2 Trigonal Bipyramidal T shaped AB2E3 Trigonal Bipyramidal Linear Molecules with Six Substituents AB 6 Octahedral Octahedral AB5E Octahedral Square Pyramid AB4E2 Octahedral Square Planar Using VSEPR to predict shape In order to use VSEPR to predict the shape of a molecule, you MUST have a valid Lewis Structure for the molecule! Use it to identify the number and types of substituents Identify the orientation and shape of the molecule SiO2: O = Si = O AB2 2 substituents Orientation: Linear All substituents are atoms, no unshared pairs Shape: Linear Using VSEPR to predict shape 2O O AB2E2 Orientation: Tetrahedral Shape: Bent Using VSEPR to predict shape Using VSEPR to predict shape ClO3 1- SF2 BF3 XeF4 AB3E O: Tetrahedral S: Trigonal Pyramid AB2E2 O: Tetrahedral S: Angular/Bent AB3 O: Trigonal Planar S: Trigonal Planar AB4E2 O: Octahedral S: Square Planar 8
Molecular Polarity Molecular Polarity Polar Molecules Aka Dipoles Entire molecules that have an uneven distribution of charge Determined by looking at the Lewis structure, VSEPR, and bond polarity Example: 2O O O: Tetrahedral S: Bent E difference: 1.4 Overall: Polar! Polar Bonds do NOT definitely mean the molecule is polar: CCl4: Cl Cl C Cl Cl O: Tetrahedral S: Tetrahedral E difference: 0.5 The arrows cancel one another out, so this molecule is NON- POLAR! No one area has more negative charge ybridization VSEPR told us about the shapes of molecules, but not about the relationship between the shape and the orbitals that are occupied by the bonding electrons ybridization The mixing of 2 or more atomic orbitals of similar energies on the same atoms to produce an equivalent number of new orbitals with equal energy ybridization Think about C4: Carbon is the central atom, and its normal valence electrons look like this: 2p 2s These orbitals mix together to create 4 new, identical orbitals: Notice that now all of the electrons are unpaired and able to bond with other atoms! New orbitals are named after the original orbitals mixed together to create them: sp 3 ybridization s sp 3 s sp3 C sp 3 s sp 3 s 10.4 9
Possible ybridization Combinations 1 s +1 p = 2 sp orbitals 1 s +2 p s = 3 sp 2 orbitals 1s +3 p s = 4 sp 3 orbitals 1 s +3 p s +1 d = 5 sp 3 d orbitals 1 s +3 p s +2 d s = 6 sp 3 d 2 orbitals Types of ybridization Since hybridization is linked to orientation, here is an easy guide: 2 substituents Linear sp 3 substituents Trigonal Planar sp 2 4 substituents Tetrahedral sp 3 5 substituents Trigonal bipyramidal sp 3 d 6 substituents Octahedral sp 3 d 2 Practice Determine the hybridization and Polarity of the central atom for the following molecules: 1) CO 2 1) O = C = O, (sp hybrid) on C 2) SO 2 2 3) SF 6 2) S,(sp hybrid) on S 4) SiO 2 O O F F 3) F S F, (sp 3 d 2 ) F F 4) O = Si = O, (sp hybrid) on Si Molecular Bonding Molecular Orbitals Resulting orbital formed by the overlap of 2 atomic orbitals. 2 Types of Molecular Bonds Sigma Bond Molecular orbital that is formed along the bonding axis. Pi Bond Molecular orbital that is formed above or below the bonding axis. Usually found as part of multiple bonds Sigma Bond Sigma with Pi bond 10
Sigma with 2 Pi bond Pi bond ( ) electron density above and below plane of nuclei Sigma bond ( ) electron density between the 2 atoms of the bonding atoms 10.5 Sigma ( ) and Pi Bonds ( ) Single bond Double bond Triple bond 1 sigma bond 1 sigma bond and 1 pi bond 1 sigma bond and 2 pi bonds ow many and bonds are in the acetic acid (vinegar) molecule C 3 COO? O C C O bonds = 6 + 1 = 7 bonds = 1 10.5 10.5 Intermolecular Forces of Attraction The forces of attraction between molecules Weaker than an ionic or covalent bond What pulls molecules together in the liquid and solid phases Not all types of F of A are equivalent 3 types: Dipole-Dipole ydrogen Bonding London Dispersion Intermolecular Forces of Attraction Dipole-Dipole Attractions Attraction between two polar molecules Partial + of one molecule attracted to the partial of another molecule δ - δ + 11
Intermolecular Forces of Attraction ydrogen Bonds Special type of dipoledipole attraction Attraction between in a covalently bonded molecule and unshared pairs of electrons in other molecules Responsible for many of the special behaviors of water! Takes much more energy to break bonds than normal forces of attraction, so water has a very high boiling point bonds are responsible for water freezing into a hexagonal pattern, creating an empty space which causes frozen water to have less density and float! Intermolecular Forces of Attraction London Dispersion Forces Occur due to the movement of the electrons Electrons can bunch up on one side of the atom, creating a temporary (induced) dipole, which can then influence other molecules. Weakest of all Forces of Attraction Typically increase as the number of electrons in the molecule increase If you look at the halogens, F and Cl are gases at room temperature, Br is a liquid, and I is a solid! 12