Covalent Bonds Ch. Why do atoms bond? Atoms want noble gas configuration ( ) For bonds there is a transfer of electrons to get an octet of electrons

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1 Covalent Bonds Ch. Why do atoms bond? Atoms want noble gas configuration ( ) For bonds there is a transfer of electrons to get an octet of electrons For covalent bonds there is a of electrons to get an octet What is a covalent bond? Covalent bond - is the chemical bond that results from sharing of electrons Occurs with elements to each other on the periodic table Between a nonmetal and a nonmetal Molecule is two or more atoms are bonded Examples of Molecules F 2 H 2 O NH 3 ( ) CH 4 (methane) Notice there are no, only non-metals Diatomic molecules Some atoms do not exist as a atom Atoms that exist as two: H 2, O 2, N 2, Cl 2, Br 2, I 2, F 2 HONClBrIF Magnificent -don t forget H Single Covalent Bonds Each atom shares one of electrons Each atom shares electrons bond Weakest bond of the three Double Covalent Bonds In a double bond, each atom shares pairs of electrons Each atom shares 4 electrons Medium length bond Medium bond Triple Covalent Bond In a triple covalent bond, each atom shares three pairs of electrons Each atom shares 6 electrons The shortest bond The strongest bond Carbon, nitrogen, oxygen, and sulfur can form double and covalent bonds Covalent Molecule Properties Covalent molecular solids tend to be solids, liquids, or gases at room temperature melting and boiling points conductors of heat and electricity Non-electrolytes do not conduct electricity in water Strength of Covalent Bond Several factors control bond strength Number of shared electrons-the electrons shared, greater the bond of the atom Bond the greater the bond length, the weaker the bond 1

2 Naming Binary Molecular Compounds 3 Rules to Name 1. Name the first element using the name 2. Second element in the formula- use the root word and in ide ex: Oxygen Oxide Sulfur Sulfide Hydrogen 3. Add a prefix to both words to indicate the of atoms 1. mono- 6. hexa- 2. di- 7. hepta- 3. tri- 8. octa- 4. tetra- 9. nona- 5. penta- 10. deca- 2 Exceptions to the rule 1. When the formula contains one atom of the element, omit (leave out) mono ex: CO 2 Carbon dioxide, not Dioxide 2. Drop the final letter in the prefix if the element begins with a - For prefixes 1 & 4-9 ex: CO Carbon monoxide, not Carbon Example 1 S 4 N 2 Example 2 SO 3 Example 3 P 4 S 5 Example 4 CO Example 5 NH 3 common name: Example 6 CH 4 common name: Example 7 As 2 O 3 Example 8 N 2 O 5 Example 9 P 4 O 10 2

3 5 rules for Lewis Structures Lewis Structures: Electron dot structures for molecules 1. Find the total # of electrons for the molecule 2. Find the center atom (the element with the # of atoms) 3. Draw bonds. Connect the other atoms to the center atom. Then subtract electrons from the total # of valence electrons for each bond drawn. 4. Distribute electrons around each atom to give a total of electrons except H, Al, B, & Be 5. If there are not enough to give 8 around each atom, create & triple bonds. Example 1: CF 4 Example 2: NH 3 Example 3: H 2 S Ions: With Ions we add or take away. Put ions in and the charge on the outside Example 4: PO 4 3- Double Bonds Example 5: CO 2 Incomplete Octet Example 6: BCl 3 Boron does not have 8e- around it. It is but it is okay; it is one of our exceptions. 3

4 Expanded Octet Ex 7: PCl 5 Triple Bond Example 8: CO Try doubling or tripling the bonds if you do not have an. Lewis Resonance Structures: Electron dot structures for molecules A Lewis Resonance Structure is a condition in which there is than 1 valid Lewis Structure. Example 1: NO 3 - Try these examples: remember they are not so they do not need to be in square brackets Example 2: SO 3 Example 3: O 3 What is wrong with my Lewis Structure? It will not be because of the lone pair electronegative repulsion with the double bond. 4

5 NT: VSEPR- V S E P R Lewis Structures are, whereas VSEPR Molecules are VSEPR predicts the or of the molecule pairs of electrons influence the shape by pushing other atoms as far apart from each other as possible Molecular Shape Molecule Lewis Structure Number of (shared) bonding pairs of electrons H 2 Number of lone pairs of electrons Total Number of electron pairs Molecular Shapes (look at lone & bond pairs) Linear Ball and Stick Model Bond Angle CO 2 Linear BH 3 Trigonal planar SiF 4 Tetrahedral PH 3 Trigonal Pyrimidal SCl 2 Bent 5

6 Electronegativity Electronegativity and Polarity Relative ability of an atom to electrons in a chemical bond has the highest e.n. value. The trend is increasing going across a period from left to right and as you go down a group Types of chemical bonds Bonds are never completely ionic or covalent Three types of bonds Ionic- of electrons Non-polar covalent- sharing of electrons Polar covalent- unequal of electrons Ionic Bond There is a large difference in electronegativity Bond between a and a non-metal Example: LiCl Bond between a cation and an anion Example: NH 4 NO 3 Non-polar covalent Polar covalent Sharing of electrons Usually occurs when two atoms are bonded together. Examples: H 2, O 2, N 2, Cl 2, Br 2, I 2, F 2 Polar covalent continued... Polar covalent bond occurs when there is an sharing of electrons The more electronegative atom attracts electrons more strongly and gains a slightly charge. The less electronegative atom has a slightly charge. Each is called a dipole (two poles) Electrons shared equally Partial charges symbolized by (delta) + and - The electronegative atom is located at the partially negative end Example: + - H Cl or H-Cl 6

7 Polar covalent example Polar molecule or not? The of a molecule usually tells if a molecule is polar or not If the VSEPR shape is it is usually non-polar If the molecule is asymmetric it is Polar molecule or not? cont. Name Polar example VSEPR Model Non-polar example VSEPR Model Linear HCl CO 2 Trigonal planar CH 2 O AlH 3 Tetrahedral CH 3 OH CH 4 Trigonal pyramidal (Always polar) Bent (Always polar) NH 3 N/A N/A H 2 O N/A N/A Intermolecular Forces Three types of forces (first two forces are called van der Waals forces) Dispersion force or induced dipole attraction between non-polar molecules; weak force (caused by the motion of electrons) (Ex: ) Dipole-dipole the force between two polar molecules; stronger force (Ex: ) Hydrogen bond the intermolecular force forms between hydrogen end of one dipole and fluorine, oxygen, or nitrogen end of another dipole; force. (Ex: ) Solubility of polar molecules are due to intermolecular forces Like dissolves like Polar substances will dissolve molecules - substances will dissolve non-polar molecules 7

8 Molecular Shape The Shape of Molecules The three dimensional shape or configuration of a molecule is an important characteristic. This shape is dependent on the preferred spatial orientation of covalent bonds to atoms having two or more bonding partners. Three dimensional configurations are best viewed with the aid of models. In order to represent such configurations on a two-dimensional surface (paper, blackboard or screen), we often use perspective drawings in which the direction of a bond is specified by the line connecting the bonded atoms. In most cases the focus of configuration is a carbon atom so the lines specifying bond directions will originate there. As defined in the diagram on the right, a simple straight line represents a bond lying approximately in the surface plane. The two bonds to substituents A in the structure on the left are of this kind. A wedge shaped bond is directed in front of this plane (thick end toward the viewer), as shown by the bond to substituent B; and a hatched bond is directed in back of the plane (away from the viewer), as shown by the bond to substituent D. Some texts and other sources may use a dashed bond in the same manner as we have defined the hatched bond, but this can be confusing because the dashed bond is often used to represent a partial bond (i.e. a covalent bond that is partially formed or partially broken). The following examples make use of this notation, and also illustrate the importance of including non-bonding valence shell electron pairs (colored blue) when viewing such configurations. Methane Ammonia Water Bonding configurations are readily predicted by valence-shell electron-pair repulsion theory, commonly referred to as VSEPR in most introductory chemistry texts. This simple model is based on the fact that electrons repel each other, and that it is reasonable to expect that the bonds and non-bonding valence electron pairs associated with a given atom will prefer to be as far apart as possible. 8