Chemical Bonding. 8.1 Types of Bonds. 8.1 Types of Bonds. : A force that holds atoms together in a molecule or compound

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1 : Chemical Bonding Types of Bonds : A force that holds atoms together in a molecule or compound Two types of chemical bonds Ionic Bonds Covalent Bonds Types of Bonds 8-3 1

2 8.1 Types of Bonds A bond created by electrostatic attraction between oppositely charged ions Occurs between a metal and a nonmetal to form ionic compounds Electrons transferred between the cation and the anion Extremely strong bonds Types of Bonds : A bond created by the sharing of electrons between atoms to form molecules Occurs between two nonmetals (no ions) Electrons not transferred in this case Electrons typically shared in pairs Weaker bonds than ionic bonds 8-5 : Identifying Types of Bonding Identify the type of bonding in each of the following substances: 1. NaF 2. ClO 2 3. FeSO 4 4. SO 2 5. Ca(ClO 2 )

3 8.1 Types of Bonds There are 2 types of covalent bonds: : Equal sharing of electrons Occurs only when all of the atoms in a molecule are the same element Typically longer bonds Weaker bonds : Unequal sharing of electrons Occurs when different elements are covalently bonded to one another Due to different elements having different electronegativities Typically shorter bonds Stronger bonds due to their increased ionic character Types of Bonds Types of Bonds In polar covalent bonds, one atom attracts the bonding electrons to itself more strongly than the other does. : the ability of an atom to attract bonding electrons to itself No difference in electronegativity between atoms in a covalen bond results in a nonpolar covalent bond. The greater the difference in electronegativity between bonding atoms, the greater the ionic character and the more polar the bond that joins the atoms. Higher electronegative element pulls bonding electrons more strongly and obtains a partial negative charge (δ ). The less electronegative element loses electron density in the bond and therefore obtains a partial positive charge (δ+) 8-9 3

4 Electronegativity Values Figure Types of Bonds The difference in electronegativity between metals and nonmetals is so large, that the electrons are transferred, not shared, creating ionic bonds. The greater the electronegativity difference in bonded nonmetals, the more polar the bond. Si-F > N-F> O-F >F-F What partial charges go on each atom? 8-11 : Polar Bonds Which of the following molecules have polar bonds? If a bond is polar, which atom has a partial negative charge? 1. SO 2 2. N 2 3. H 2 S 4. CCl 4 5. O

5 8.3 Covalent Bonding When two nonmetals form a bond, the bond is covalent. They are both close to the noble-gas electron configuration, so sharing will allow both to obtain it. In a covalent bond, each shared electron interacts simultaneously with two nuclei. Covalent bonds are formed so that each atom can be isoelectric with a noble gas. Noble gases have 8 valence electrons, an Covalent Bonding : Tendency of an atom to achieve an electron configuration having 8 valence electrons Same as the electron configuration of a noble gas Covalently bonded atoms achieve 8 valence electrons by sharing electrons The 8 electrons exist in 4 pairs H atoms bond with other atoms to obtain a total of 2 electrons like He (duet rule.) How does hydrogen obtain a noble-gas electron configuration? Why is hydrogen a diatomic molecule? Covalent Bonding Do the atoms in each of these molecules have an octet? Why do the halogens exist as diatomic molecules? How many valence electrons does an oxygen atom have? How many does it need to obtain an octet? Can a single covalent bond allow each oxygen atom to satisfy the octet rule? Diatomic oxygen has a double bond (2 pairs of shared electrons)

6 8.3 Covalent Bonding How many valence electrons does a nitrogen atom have? How many does it need to obtain an octet? N 2 has a triple bond (three pairs of shared electrons) : How many bonds do each of the following atoms tend to form? a) H b) Cl c) O d) N e) C Covalent Bonding: Lewis Structures 1. Count the available valence electrons by adding up the valence electrons of each of the atoms. 2. Use 6N+2 to determine if there are multiple bonds, where N = the number of atoms, not including hydrogen. If 6N+2=ve, all single bonds If 6N+2=ve + 2, 1 double bond If 6N+2=ve + 4, 2 double bonds OR 1 triple bond 3. Arrange the atoms Put atoms that want the most atoms in the middle. Remember that H can only form 1 bond. 4. Add bonds 1 line for each pair of electrons Single bold 1 line, Double bond 2 lines, Triple bond 3 lines 5. Use remaining electrons as lone pairs to complete octets Hydrogen will only have a duet. 6. Double check that all non-h atoms have an octet, and that the total number of electrons used matches the number of valence 8-17 electrons Copyavailable. right Mc Graw-Hill Educ ation. Permis sion required for reproduction or dis play. : Writing Lewis Structures Write Lewis structures for each of the following: a. Formaldehyde molecule, CH 2 O b. Hydrocyanic acid c. Carbon dioxide d. Sulfur dioxide e. Nitrite ion f. Nitric acid 1. Count ve 2. Calculate 6N+2 3. Arrange the atoms 4. Add bonds 5. Add lone pairs 6. Double check

7 8.3 Covalent Bonding: Lewis Structures : When a molecule can have more than one Lewis structure while only moving electrons. Sometimes a single Lewis structure does not give an accurate picture of the bonding. For example, there are two Lewis structures that represent the ozone, O 3, molecule. The actual molecule is a blend of these two: Both O O bonds are the same length with partial double bond character. : Write resonance structures for N 2 O. (Nitrogen is in the middle) 8-19 : Resonance Structures Which of the following have resonance? For those that do, how many resonance structures do they have? 1.NO 3 2.HNO 3 3.CH 4 4.H 2 CO Covalent Bonding Molecules with an odd number of electrons example: NO Incomplete octets example: BH 3 or BF 3 Expanded octets example: SF 4 or SF

8 8.4 Bonding in Carbon Compounds : Has four valence electrons Has the ability to form four bonds Has the ability to bond to itself Forms very strong bonds when bonded to itself Carbon molecules are ubiquitous in nature. Hydrocarbons: Compounds that mainly composed of carbon and hydrogen Bonding in Carbon Compounds Alkanes: Have only single bonds Alkenes: Have at least 1 double bond Alkynes: Have at least 1 triple bond : Linear chain hydrocarbons Bonding in Carbon Compounds : hydrocarbons with carbon atoms arranged in a six-atom ring with alternating single and double bonds Delocalized structures

9 8.4 Bonding in Carbon Compounds Bonding in Carbon Compounds : A group of atoms in a hydrocarbon chain Gives the hydrocarbon its characteristic properties The group has a heteroatom, an atom other than C and H Typically O, S, and N : contains a hydroxyl group (-OH) replaces a hydrogen atom in the formula for a hydrocarbon Bonding in Carbon Compounds

10 8.5 Shapes of Molecules The shape of a molecule influences many of its properties including taste and smell. Sweet substances are often of a shape similar to that of glucose. It s H and OH groups fit into a taste receptor site on the tongue. : a theory used to predict molecular shapes based on the fact that negative charges repel one another. Valence electron pairs (bonding or nonbonding) repel one another and take up positions that maximize their distance and angles between them. To determine the shape of a molecule, count the number of electron groups. Electron groups are: - Nonbonding electron pair - Single bond - Double bond - Triple bond Shapes of Molecules 8-29 : Molecular Shapes Determine the parent structure and bond angles for each of the following. 1. NH 3 2. CO 2 3. SO 2 4. BH 3 5. H 2 O

11 8.5 Shapes of Molecules The actual shape of a molecule includes only the atoms, and leaves out the nonbonding lone pairs of electrons. To envision the shape from a parent structure, make the nonbonding electrons invisible (but they are still there). What is the shape of a water molecule? Why do CO 2 and SO 2 have different shapes? 8-31 Steps for Determining VSEPR Shapes 1. Draw a Lewis formula. 2. Count the electron groups around the central atom. 3. Determine parent shape. 4. Count the number of lone pairs 5. Eliminate the lone pairs from the parent shape and determine the VSEPR shape of the central atom

12 8.5 Shapes of Molecules Trigonal planar and trigonal pyramidal molecular shapes are different from one another. One has a lone pair, causing it s shape to be different : VSEPR Structures What is the shape of the nitrite ion? What is the O N O bond angle? Shapes of Molecules Glycine: H 2 NCH 2 CO 2 H 1. Draw the Lewis structure. 2. Determine the parent structure, remembering to consider all the electron domains. 3. Determine the bond angles at each central atom (N, C, C, and O)

13 : What are the bond angles? 8-37 : What are the bond angles? Shapes of Molecules Bond polarity results from atoms in the bond pulling on the bonding electrons unevenly (electronegativity) Molecular polarity results from bond polarity tending to pull the electrons to one side of the molecule. A nonpolar molecule is one that has all nonpolar bonds or one that has polar bonds that cancel out because: - A central atom has no unshared electrons - The atoms around the central atom all have the same electronegativity Sulfur trioxide has polar bonds, but its symmetrical shapes with no nonbonding lone pairs of electrons on the central atom, the bond polarity cancels out and makes it a nonpolar molecule

14 8.5 Shapes of Molecules Hydrocarbons without functional groups are nonpolar Shapes of Molecules In asymmetric molecules like these, the atoms surrounding the central atom are not identical and have different bond polarities. These molecules are polar. Figure : Molecular Polarity Which of these molecules are polar? Predict whether the following are polar or nonpolar molecules. 1. NH 3 2. BH 3 3. CS 2 4. H 2 S 5. O 3 6. CH 2 Cl

15 8.5 Shapes of Molecules Molecular polarity is important in many properties of molecules, such as solubility (Chp 11.) A rule of thumb for solubility is like dissolves like, which refers to similarity in polarity. Ionic salts and polar liquids dissolve better in polar liquids than in nonpolar liquids. Nonpolar liquids dissolve better in other nonpolar liquids than in polar liquids