Lecture 5

Lecture 5 - 3.4 Drawing Lewis Structures of ...

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Unformatted text preview: 3.4 Drawing Lewis Structures of Molecules and Polyatomic Ions Lewis Structures of Molecules Lewis Structures of Polyatomic Ions Lewis Structure, Stability, Mul;ple Bonds, and Bond Energies Isomers Lewis Structures and Resonance Lewis Structures and Excep;ons to the Octet Rule Lewis Structures and Molecular Geometry; VSEPR Theory Periodic Structural Rela;onships Lewis Structures and Polarity 3.5 Proper;es Based on Electronic Structure and Molecular Geometry Solubility Boiling Points of Liquids and Mel;ng Points of Solids Guidelines for drawing Lewis Structure ①  Use symbols for elements to write the skeletal structure of the compound The least electronega;ve atom placed at the center Hydrogen and halogens occupy terminal posi;ons Carbon oNen forms chains of carbon- carbon covalent bonds ②  Determine the number of valence electrons for each atom Combine to determine the total number of valence electrons Polyatomic ca;ons, subtract one electron for every posi;ve charge Polyatomic anions, add one electron for every nega;ve charge ③  Connect the central atom to each of the surrounding with single bonds. ④  Next, complete octets of all the atoms bonded to the central atom Hydrogen needs only two electrons. Electrons not involved in bonding are represented as lone pairs. ANer the terminal atoms have an octet, provide the central atom with an octet if valence electrons are s;ll available. ⑤  If not enough valence electrons to give the central atom an octet, move lone pair electrons from terminal atoms to form a new bond with the central atom. (Con;nue to shiN the electrons un;l all atoms have an octet) ⑥  Recheck that all atoms have the octet rule sa;sfied and that the total number of valance electrons are used Lewis Structures of Covalent Compounds Carbon dioxide, CO2 1) Draw a skeletal structure of the molecule, C- O- O or O- C- O ? 2) Place the least electronega;ve atom as the central atom, EN of O=3.5 C=2.5 ∴ O- C- O 3) Find the number of valence electrons for each atom and the total 1 C atom × 4 valence electrons = 4 e- 2 O atoms × 6 valence electrons = 12 e- O−C−O 16 e- total 4) Use electron pairs to connect the C to each O with a single bond O―C―O 5) Place electron pairs around the atoms O C O This sa;sfies the rule for the O atoms, but not for C 6) Move 2 e- from each O, placing them between C―O ∴ O C O This is the most probable structure Four electrons are between C and O These electrons are share in covalent bonds Four electrons in this arrangement signify a double bond 7) Recheck the electron distribu;on 8 electron pairs = 16 valence electrons, number counted at start 8 electrons around each atom, octet rule sa;sfied Lewis Structure of Polyatomic Anions Draw the Lewis structure of carbonate ion, CO32- Draw a skeletal structure of the molecule 1.  Carbon is less electronega;ve than oxygen •  This makes carbon the central atom •  Skeletal structure and charge: 2.  The total number of valence electrons is determined by adding one electron for each unit of nega;ve charge 1 C atom x 4 valence electrons = 4 e- 3 O atoms x 6 valence electron = 18 e- + 2 nega;ve charges = 2 e- 24 e- total 3.  Distribute these e- around the skeletal structure 4.  Distribu;ng the electrons around the central carbon atom (4 bonds) and around the surrounding O atoms ajemp;ng to sa;sfy the octet rule results in: 5.  6.  This sa;sfies the octet rule for the 3 oxygen, but not for the carbon Move a lone pair from one of the O atoms to form another bond with C Using the guidelines presented, write Lewis structures for the following: 1.  H2O 2.  NH3 3.  CO 4.  NH4+ 5.  N2 Lewis Structure, Stability, Mul;ple Bonds, and Bond Energies •  Single bond - one pair of electrons are shared between two atoms •  Double bond - two pairs of electrons are shared between two atoms •  Triple bond - three pairs of electrons are shared between two atoms •  Very stable Bond energy - the amount of energy required to break a bond triple bond > double bond > single bond Bond length - the distance separa;ng the nuclei of two adjacent atoms single bond > double bond > triple bond Lewis Structures and Resonance •  Write the Lewis structure of CO32- •  In some cases it is possible to write more than one Lewis structure that sa;sfies the octet rule for a par;cular compound O O C : C :: O O :: O : C O :: :: O : : : O O •  Experimental evidence shows all bonds are the same length, meaning there is not really any double bond in this ion •  None of theses three Lewis structures exist, but the actual structure is an average or hybrid of these three Lewis structures •  Resonance - two or more Lewis structures that contribute to the real structure Lewis Structures and Excep;ons to the Octet Rule 1.  Incomplete octet - less then eight electrons around an atom other than H –  Let’s look at BeH2 1 Be atom × 2 valence electrons = 2 e- 2 H atoms × 1 valence electrons = 2 e- total 4 e- –  2.  Resul;ng Lewis structure: H – Be – H Odd electron - if there is an odd number of valence electrons, it is not possible to give every atom eight electrons •  Let’s look at NO, nitric oxide •  It is impossible to pair all electrons as the compound contains an ODD number of valence electrons 3.  Expanded octet - an element in the 3rd period or below may have 10 and 12 electrons around it •  Expanded octet is the most common excep;on •  Consider the Lewis structure of PF5 •  Phosphorus is a third period element F FP F F : :F: : :: : 1 P atom × 5 valence electrons = 5 e- 5 F atoms × 7 valence electrons = 35 e- 40 e- total F :: F: :F P F: •  Distribu;ng the electrons results in this Lewis structure :F: : Lewis Structures and Molecular Geometry: VSEPR Theory •  Molecular shape plays a large part in determining proper;es and shape •  VSEPR theory - Valance Shell Electron Pair Repulsion theory •  Used to predict the shape of the molecules •  All electrons around the central atom arrange themselves so they can be as far away from each other as possible – to minimize electronic repulsion. VSEPR Theory ①  In the covalent bond, bonding electrons are localized around the nucleus ②  The covalent bond is direc)onal, having a specific orienta;on in space between the bonded atoms ③  Ionic bonds have electrosta;c forces which have no specific orienta;on in space •  BeH2 (stable excep;on to octet rule) –  Only 4 electrons surround the beryllium atom –  These electrons in the bonds to the two atoms have minimal repulsion when located on opposite sides of the structure –  Linear structure having bond angles of 180° •  BF3 –  There are 3 bonded atoms around the central atom –  These atoms have minimal repulsion when placed in a plane, forming a triangle –  Trigonal planar structure with bond angles of 120° •  Consider CH4 –  There are 4 bonded atoms around the central carbon. –  Minimal electron repulsion when electrons are placed at the four corners of a tetrahedron. –  Each H- C- H bond angle is 109.5° •  Tetrahedron is the primary structure of a full octet Consider NH3 •  There are three bonded atoms and one lone pair (four groups). –  A lone pair is more electronega;ve with a greater electron repulsion –  The lone pair takes one of the corners of the tetrahedron without being visible, distor;ng the arrangement of electron pairs •  Ammonia has a trigonal pyramidal structure with 107° angles H2O •  There are two bonded atoms and two lone pair (four groups) –  All 4 electron pairs are approximately tetrahedral to each other –  The lone pairs take two of the corners of the tetrahedron without being visible, distor;ng the arrangement of electron pairs •  Water has a bent or angular structure with 104.5° bond angles Predic;ng Geometric Shape Using Electron Pairs 1.  2.  3.  4.  Basic Procedure to Determine Molecular Shape Write the Lewis structure Count the number of bonded atoms and lone pairs around the central atom If no lone pairs are present, geometry is: •  2 bonded atoms - linear •  3 bonded atoms - trigonal planar •  4 bonded atoms - tetrahedral If there are lone pairs, look at the arrangement of the atoms and name the geometry. Names include: •  Angular •  Trigonal pyramid Periodic Structural Rela;onship Group VIA, oxygen, sulfur, and selenium (Se). Each has six valence electrons, two more electrons to complete its octet, each reacts with hydrogen, forming H2O, H2S, and H2Se. H2O, H2S and H2Se all are angular molecules with similar Lewis structures. This logic applies equally well to the other representa;ve elements. More Complex Molecules Dimethyl ether •  Has 2 different central atoms: •  oxygen •  carbon –  CH3 (methyl group) has tetrahedral geometry (like methane) –  PorIon of the molecule linking the two methyl groups would bond angles similar to water Detrmine te Molecular Geomety PCl3 S O2 P H3 SiH4 Lewis Structures and Polarity •  A molecule is polar if its centers of posi;ve and nega;ve charges do not coincide •  Polar molecules when placed in an electric field will align themselves in the field •  Molecules that are polar behave as a dipole (having two “poles” or ends) •  One end is posi;vely charged the other is nega;vely charged •  Nonpolar molecules will not align themselves in an electric field Determining Polarity To determine if a molecule is polar: •  Write the Lewis structure •  Draw the geometry •  Use the following symbol to denote the polarity of each bond Posi;ve end of the bond, the less electronega;ve atom Nega;ve end of the bond, more electronega;ve atom ajracts the electrons more strongly towards it Determining Polarity Guidelines: •  Molecules that have no lone pair on the central atom, and all terminal atoms are the same are nonpolar. •  Molecules with one lone pair on the central atom are polar. •  Molecules with more than one lone pair on the central atom are usually polar. Practce Detrmining Polarit Determine whether the following bonds and molecules are polar: 1.  Si – Cl 1. O2 2.  H – C 2. HF 3.  C – C 3. CH4 4.  S – Cl 4. H2O •  Intramolecular forces – ajrac;ve forces within molecules – Chemical bonds •  Intermolecular forces – ajrac;ve forces between molecules •  Intermolecular forces determine many physical proper;es –  Intermolecular forces are a direct consequence of the intramolecular forces in the molecules Solubility and Intermolecular Forces Solubility - the maximum amount of solute that dissolves in a given amount of solvent at a specific temperature •  “Like dissolves like” –  Polar molecules are most soluble in polar solvents –  Nonpolar molecules are most soluble in nonpolar solvents •  Does ammonia, NH3, dissolve in water? Yes, both molecules are polar Interac;on of Water and Ammonia •  The δ- end of ammonia, N, is ajracted to the δ+ end of the water molecule, H •  The δ+ end of ammonia, H, is ajracted to the δ- end of the water molecule, O •  The ajrac;ve forces, called hydrogen bonds, pull ammonia into water, distribu;ng the ammonia molecules throughout the water, forming a homogeneous solu;on Interac;on of Water and Oil •  What do you know about oil and water? “They don’t mix” –  Because water is polar and oil is nonpolar •  Water molecules exert their ajrac;ve forces on other water molecules •  Oil remains insoluble and floats on the surface of the water as it is less dense Boiling Points of Liquids and Mel;ng Points of Solids •  Energy is used to overcome the intermolecular ajrac;ve forces in a substance, driving the molecules into a less associated phase •  The greater the intermolecular force, the more energy is required leading to –  Higher mel;ng point of a solid –  Higher boiling point of a liquid Factors Influencing Boiling and Mel;ng Points •  Strength of the ajrac;ve force holding the substance in its current physical state •  Molecular mass •  Larger molecules have higher m.p. and b.p. than smaller molecules as it is more difficult to convert a larger mass to another phase •  Polarity •  Polar molecules have higher m.p. and b.p. than nonpolar molecules of similar molecular mass due to their stronger ajrac;ve force Mel;ng and Boiling Points – Selected Compounds by Bonding Type ...
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