2 Illustrating Bonds - Lewis Dot Structures Lewis Dot structures are also known as electron dot diagrams These diagrams illustrate valence electrons and subsequent bonding A line shows each shared electron pair Dots represent unpaired electrons called lone pairs
3 Steps to Draw Lewis Dot Diagrams 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 # of electrons Add lone pair electrons, :, to terminal atoms to satisfy octet rule Extras go to central atom
4 Steps to Draw Lewis Dot Structures If central atom is not octet and extra electrons are left unpaired, form multiple bonds! Carbon bonded to N, O, P, S tend to form double bonds NOTE - Hydrogen is ALWAYS a terminal atom Only makes 1 bond
5 Using Electronegativity Differences to Draw Appropriate Lewis Dot Structures Using the steps discussed on the last two slides, write the electron dot diagrams for each element in the compound Check the electronegativity difference between the elements to determine if electrons are transferred or shared If the electronegativity difference > 1.67, the reaction forms an IONIC BOND Remove the electrons from the metal and add them to the nonmetal
6 Lewis Dot Structure of Ionic Compounds Write the charges of the ions formed and use coefficients to show how many of each ion are needed to balance the overall charge + 2-2Na,[ O ] Ionic sodium oxide, Na 2 O
7 Lewis Dot Structure of Covalent Compounds If the electronegativity difference < 1.67, then the atoms will share electrons Position shared electron pairs between the two atoms, and connect them with a single line to represent a covalent bond Place the extra pairs of electrons around atoms until each has eight
8 Lewis Dot Structure of Covalent Compounds If an atom other than hydrogen or a metal has less than eight electrons, move unshared pairs to form multiple bonds Add extra atoms, if needed, to obtain the octets Atoms with positive oxidation numbers should be bonded to those with negative oxidation numbers If extra electrons still remain, add them to the central atom All oxidation numbers should add up to zero for a compound
9 Single Covalent Bonds H H C H F F H Do atoms (except H or metals) have octets?
10 Lewis Dot Structures with Multiple Bonds Example CO 2 Step 1 Draw any possible structures C-O-O O-C-O You may want to use lines for bonds. Each line represents 2 electrons.
11 Lewis Dot Structures with Multiple Bonds Step 2 Determine the total number of valence electrons. CO 2 1 carbon x 4 electrons = 4 2 oxygen x 6 electrons = 12 Total electrons = 16
12 Lewis Dot Structures with Multiple Bonds Step 3 Try to satisfy the octet rule for each atom - all electrons must be in pairs - make multiple bonds as required Try the C-O-O structure C O O No matter what you try, there is no way satisfy the octet for all of the atoms.
13 Lewis Dot Structures with Multiple Bonds O C O This arrangement needs too many electrons. How about making some double bonds? O=C=O That works! = is a double bond, the same as 4 electrons
14 Exceptions to the Octet Rule Not all compounds obey the octet rule Three types of exceptions: Species with more than eight electrons around an atom Species with fewer than eight electrons around an atom Species with an odd total number of electrons
15 Atoms with Fewer than Eight Electrons Referred to as electron deficient Examples - Beryllium and boron will both form compounds where they have less than 8 electrons around them.... :F.. B F:.. :F:..
16 Atoms with More than Eight Electrons Except for species that contain hydrogen, this is the most common type of exception For elements in the third period and beyond, the d orbitals can become involved in bonding Examples 5 electron pairs around P in PF 5 5 electron pairs around S in SF 4 6 electron pairs around S in SF 6
17 Species with an Odd Total number of Electrons A very few species exist where the total number of valence electrons is an odd number This must mean that there is an unpaired electron which is usually very reactive A radical is a species that has one or more unpaired electrons They are believed to play significant roles in aging and cancer
18 : Species with an Odd Total Number of Electrons Nitrogen monoxide (aka nitric oxide) is an example of a compound with an odd number of electrons It has a total of 11 valence electrons - six from oxygen and 5 from nitrogen The best Lewis structure for NO is:. :N O:
19 The VSEPR Theory Advanced Chemistry Ms. Grobsky
20 Determining Molecular Geometries Lewis dot structures only show the number and types of bonds they do not provide any information about the shape of the molecule The shape of a molecule is determined by its bond angles and bond lengths These are affected by the energetically favorable and energetically unfavorable interactions of electrons and protons In order to predict molecular shape, we use the Valence Shell Electron Pair Repulsion (VSEPR) theory
21 The VSEPR Model 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 To determine the shape of a molecule, we must consider all electron domains in a molecule Defined as the regions where electron pairs may be found Therefore, we must 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 In general, each non-bonding pair, single bond, or multiple bond produces a single electron domain around the central atom in a molecule
22 The Five Electron-Domain Geometries 2 e- domains 3 e- domains 4 e- domains 5 e- domains 6 e- domains
23 Why Do We Have Different Electron Domains? This is because each electron domain affects the amount of electron repulsion around the atom differently 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! Multiple bonds exert a greater repulsive force on adjacent electron pairs than do single bonds as a result of higher electron density All electrons are considered when determining 3-D shape Called the coordination number
24 Factors Affecting Bond Angles and Molecular Geometries Multiple Bonds Non-Bonding (Lone) Pairs
25 Steps in Determining a Molecular Shape Draw the Lewis dot structure Count number of electron domains (bonding pairs and lone pairs) around the central atom For molecules with MORE THAN ONE CENTRAL ATOM, work with one central atom at a time Determine the electron-domain geometry by arranging the electron domains about the central atom so that the repulsions among them are minimized Determine the molecular geometry with the help of the following formula : AX m E n A - central atom X surrounding atom E non-bonding valence electron group m and n - integers
26 The Single Molecular Shape of Linear Electron-Group Arrangement AX2 X A X
27 Example - CO2
28 The 2 Molecular Shapes of Trigonal Planar Electron-Domain Arrangement Trigonal Planar AX 3 Bent AX 2 E Example NO2 X X A X X E A X
29 Example of Trigonal Planar - BCl3
30 The 3 Molecular Shapes of the Tetrahedral Electron-Group Arrangement Tetrahedral AX 4 Trigonal Pyramidal AX 3 E Bent AX 2 E 2 X E E X A X X X A X X X A X E
31 Example of Tetrahedral - CH4
32 Example of Trigonal Pyramidal - NH3
33 Example of Bent H 2 O O H H
34 The 4 Molecular Shapes of the Trigonal Bipyramidal Electron-Group Arrangement Trigonal Bipyramidal AX 5 Examples PCl 5, PF 5, AsF 5, SOF 4 See-Saw T-Shaped Linear AX 4 E Examples SF 4, XeO 2 F 2, IF 4 +, IO 2 F 2 - AX 3 E 2 Examples ClF 3, BrF 3 AX 2 E 3 Examples XeF 2, I 3 -, IF 2 -
35 The 3 Molecular Shapes of the Octahedral Electron-Group Arrangement Octahedral AX 6 Examples SF 6, IOF 5 Square Pyramidal AX 5 E Examples BrF 5, XeOF 4, TeF 5 - Square Planar AX 4 E 2 Examples XeF 4, ICl 4 -
36 Time to Explore the Different Molecule Arrangements! VSEPR Simulation
38 VSEPR Theory and Polar/Nonpolar Molecules Most bonds between atoms of different elements in a molecule are polar This does not mean that the molecule will be polar O = C = O The electronegativity values show that the C-O bond would be polar with electrons being pulled towards the oxygens The vectors represent a dipole moment, µ, which is a separation of the charge in a molecule (slightly positive/slightly negative poles) The dipole moment increases as the magnitude of the charge that is separated increases and as the distance between the charges increases However due to the linear geometry of CO 2, the pull happens in equal and opposite directions so overall, the molecule is NON-POLAR! Electronegativities: Oxygen = 3.5 Carbon = 2.5 Difference 1.0 (Polar Bond)
39 VSEPR Theory and Polar/Nonpolar Molecules As we just saw, in order for a molecule to be polar, the effects of bond polarity must not cancel out One way is to have a geometry that is not symmetrical like the bent geometry in water H O H Electronegativity Difference = 1.3 Here, the effects of the polar bonds do not cancel, so the molecule is polar
40 Polar vs. Nonpolar Molecules A molecule is nonpolar if the central atom is symmetrically substituted by identical atoms CO 2, CH 4, CCl 4 A molecule will be polar if the geometry is not symmetrical H 2 O, NH 3, CH 2 Cl 2 The degree of polarity is a function of the number and type of polar bonds as well as the geometry
41 Geometry and Polar Molecules Again, in order for a molecule to be polar: Must have polar bonds Must have the proper geometry Examples: CH 4 Non-polar CH 3 Cl Polar CH 2 Cl 2 Polar CHCl 3 Polar CCl 4 Non-polar WHY?
42 Why is Polarity So Important? It affects physical properties such as melting point, boiling point and solubility Chemical properties also depend on polarity
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