Honors Unit 7 Notes Bonding & Molecular Geometry (Chapter 7 in Masterton & Hurley textbook)

Transcription

1 Name: Honors Unit 7 Notes Bonding & Molecular Geometry (Chapter 7 in Masterton & Hurley textbook) Objectives: 1. Students will form an understanding of the conceptual difference between covalent and ionic bonds and the resulting structures based on these two major types of bonds. 2. Students will be able to draw Lewis structures for neutral atoms, ions, ionic bonds, and covalent compounds. 3. Students will apply their understanding of the concept of electronegativity and be able to use that to predict the polarity of individual bonds. 4. Students will be able to successfully apply VSEPR Theory to predict molecular shapes, bond angles and molecular polarity including structures with and without nonbonding pairs of electrons. 5. Students will be able to predict whether or not molecules are dipoles (polar). 6. Students will become familiar with the concept of hybridization and will be able to predict the hybridization of central atoms in simple molecules. Types of Chemical Bonds_ o Definition of a chemical bond-- o Electrons that are available for bonding are called electrons. o Two types of chemical bonds are and. IONIC BOND: COVLENT BOND: o Octet Rule: Most atoms prefer to have electrons in their valence shell. ***Exception: want a duet.

2 Lewis Structures- 2 Lewis dot diagram depicts an atom as Ex.) Lewis Dot Diagram for Carbon How to Write a Lewis Dot Structure for an Element: Electrons are placed one at a time in a clockwise manner around the symbol in the north, east, south and west positions,. Same group # = Ex.) s we can see from the chlorine example, there are six electrons that are paired up and one that is unpaired. When it comes to bonding, atoms tend to. bond that forms when one atom gives an unpaired electron to another atom is called an. bond that forms when atoms share unpaired electrons between each other is called a. ***Note: In the final structure, placement of the dots around the element is not crucial. Maximum # of valence electrons =

3 Lewis Structures for Ionic Compounds 3 1. The ionic bond forms from attraction of for. 2. Structure of ionic compounds = crystal lattice a. Ionic compounds have which show the ratio of ions in the crystal lattice. 3. To write a Lewis structure for an ionic compound: a. Uses or dots, brackets and a superscript to designate ion charge. Ex.) Li 1+ Be 2+ B 3+ C 4+ N 3- O 2- F 1- Ex.) MgO Ex.) Li2O Lewis Structures for Covalent Compounds Covalent compounds are substances made up of atoms which are held together by covalent bonds (in which ). o Covalent compounds are also called.

4 4 Lewis structures for covalent compounds show the bonds and how the atoms will connect. o Shared electrons (bonding pair electrons) = Bonding pair generally shown as a between atoms. o Non-shared electrons (non-bonding electrons) = Ex.) H2O Types of Covalent Bonds o Single Bond are shared in a bond ( from each atom) o Double Bond of e- are shared ( total, from each atom) o Triple Bond of e- are shared ( total, from each atom) Rules for Drawing Lewis Structures 1. dd up the total number of valence e - for each atom in the molecule. a. Each (-) sign counts as 1 electron added to the total, each (+) sign subtracts one electron from the total. 2. Write the symbol for the central atom, then use one pair of electrons to form bonds between the central atom and the remaining atoms. 3. Count the number of electrons remaining and distribute to the surrounding atoms according to the octet rule (or the duet rule for hydrogen). 4. If there are not enough pairs, make sure the most electronegative elements are satisfied. Then, start shifting pairs into double bonds and triple bonds to satisfy the octet rule. 5. If there are extra electrons, stick them on the central atom. *****But Remember...The Structure MUST Have: the right number of atoms for each element, the right number of electrons, the right overall charge, and 8 electrons around each atom (ideally) Hints: (1) H is NEVER a central atom. (2) Halogens are USULLY NOT central atoms. (3) IF you have ONLY ONE of a certain element, it is USULLY the central atom.

5 Examples: Drawing Covalent Lewis Structures F2 HF 5 H2O NH3 OCl - NH4 + PO4 3- SO3 2- O2 N2 CH4 CH3OH

6 Exceptions to the Octet Rule 6 1. Reduced Octets electron deficient molecules (Be and B) Be: B: 2. Expanded Octets (P, S, Cl, s, Se, Br, Kr, e) How to recognize: o The central atom in Period 3 or greater is surrounded by more than 4 atoms. Examples: PF5 o You draw the Lewis diagram & the results don t make sense central atom has > 8 e- Examples: ef4 3. Resonance Structures when a single Lewis structure does not adequately represent a substance; the true structure is o Resonance Structures are created by, NOT. o Resonance Examples: SO2 NO3 -

7 4. Radicals 7 o Example: NO o Radicals are ; they want to have paired electrons to be more stable! Molecular Geometry- Molecular geometry describes the of atoms in a molecule. o There is a fundamental geometry that corresponds to the total number of electron pairs around the central atom: 2, 3, 4, 5 and 6 VSEPR Theory: Dictating the Shape of Molecules VSEPR Theory = o Based on idea that electron pairs want to be as far apart as possible The molecule adopts the shape that minimizes the e - pair Based on molecular shape of Lewis structure o We define the 3D space of LL electron pairs (bonding and non-bonding). by the positions in o The only considers the positions of the bonded electrons.

8 To determine the electron pair geometry: (To determine the molecular geometry) Look at the shape of the molecule based on the bonded atoms (not the lone pair electrons on the central atom). Basic Electron Pair Geometries SHPE 1. linear SUM OF BONDED TOMS & LONE PIR ELECTRONS 2. trigonal planar 3. tetrahedral 4. trigonal bipyramidal 5. octahedral Molecular Geometry Notation : BPs: : LPs: E:

9 Possible Molecular Geometry Options (you must memorize these!!) Total Bonds to Central tom ( + E) Bonded toms () Lone Pairs on C.. (E) E Notation Molecular Geometry Name Shape Bond ngle(s) 9 Example 2 (Linear) 2 0 BeCl2 3 (Trigonal Planar) 3 0 BCl3 2 1 NO2-4 (Tetrahedral) 4 0 CCl4 3 1 NH3 5 (Trigonal Bipyramidal) H2O scl5 4 1 SeCl4 3 2 BrCl3 2 3 ef2 6 (Octahedral) 6 0 TeBr6 5 1 BrF5 4 2 ef4

10 Tips for Determining the Molecular Geometry of a Molecule: o If there is more than one central atom, you must determine the geometry for each central atom separately. 10 o Molecules with only two atoms are always linear. Examples: What shape (molecular geometry) would the compounds have based on VSEPR theory? O2 SeCl2 2- PBr5 CFCl3 H2S C2H2

11 11 Polarity_ Predicting Bond Polarity o Polar bond = Example: o Nonpolar bond = Example: Calculate the electronegativity difference between the Pauling electronegativity values for the 2 elements involved in the bond. The difference will dictate the type of bond. Type of Bond IONIC POLR Covalent NON-POLR Types of toms Electronegativity Difference Non-polar covalent bond Polar covalent bond (more e/n element has greater pull) 1.7 and up Ionic bond (electrons are transferred between atoms) Examples: What type of bond is shown below? Cl--Na O--F

12 Predicting Molecule Polarity 12 Polar molecule (dipole) = o Molecules are polar if the pull in any one direction is not balanced out by an equal and opposite pull in the opposite direction o Positive and negative regions shown by (δ + or δ - ) o Example: CH3Cl o Polar molecules have a ; they will align with an electric field Rules for Determining Molecular Polarity (see flowchart!) molecule is polar if: o It has only 2 atoms in it and both are different. o It has 3 or more atoms and has lone pairs on the central atom (it is classified as an E) **Exception: when lone pairs are symmetrical to axis of bonded atoms (4E2 or 2E3) o It has 3 or more atoms in an n classification and all of the s are not the same atom. Examples: Determine the polarity of the following molecules based on their molecular geometry. BF3 BF2Cl CO2 H2O H3O + ef4 PF5

13 Hybridization of tomic Orbitals 1. The solutions of the Schrodinger equation led to these atomic orbitals: However, overlap of these orbitals does not give a satisfactory explanation. In order to explain bonding, these orbitals are to form new sets of orbitals this method is called. sp sp 2 sp 3 sp 3 d sp 3 d 2 2 sp hybrid orbitals from mixing of a s and a p orbital 3 sp 2 hybrid orbitals from mixing of a s and 2 p orbitals 4 sp 3 hybrid orbitals from mixing of a s and 3 p orbitals 5 sp 3 d hybrid orbitals from mixing of a s and 3 p and a d orbital 6 sp 3 d 2 hybrid orbitals from mixing of a s and 3 p and 2 d orbitals ***Notice that there are five hybrid orbital types (same number as electron pair geometries). Superscripts on s, p, d added together = Examples: Predict the molecular geometry and the hybridization of the central atom in the following: OF2 NH4 + CO2 COCl2 ef4