Unit 5 Chemical Bonding

Transcription

1 Unit 5 Chemical Bonding Ionic and Metallic Bonding Ionic Compounds Compounds composed of cations and anions are called ionic compounds. Although they are composed of ions, ionic compounds are electrically neutral. The that hold ions together in ionic compounds are called. Formula Units NaCl is the chemical formula for sodium chloride, but there are more than one sodium ion and chloride ion within the ionic crystal. 1

2 Practice Properties of Ionic Compounds Most ionic compounds are crystalline solids at room temperature. Ionic compounds generally brittle and have high melting points. Metallic Bonds The valence electrons of metal atoms can be modeled as a sea of electrons.» The valence electrons are mobile and can drift freely from one part of the metal to another.» Metallic bonds 2

3 Metallic Bonds A force can change the shape of a metal, but a force can shatter an ionic crystal. Metals also have crystal structures. Metal atoms are arranged in very compact and orderly patterns. Properties of Metals Name three properties of metals. Describe how these properties can be explained by the nature of a metallic bond. 3

4 Alloys Alloys are mixtures composed of two or more elements, at least one of which is a metal. Alloys are important because their properties are often superior to those of their component elements. The most important alloys today are steels. Steels have a wide range of useful properties, such as corrosion resistance, ductility, hardness, and toughness. Unit 5 Chemical Bonding Molecular Compounds and Lewis Dot Structures Molecular Compounds What is a covalent bond? The atoms held together by sharing electrons are joined by a. A is a neutral group of atoms joined together by covalent bonds. Molecules are composed of and. 4

5 Ionic vs Covalent Properties of Molecular Compounds Molecular compounds tend to have relatively melting and boiling points than ionic compounds. This is because molecular compounds do not have closely packed crystal structures. Physical properties are dependent upon. Ethane, a component of natural gas, is also a molecular compound. Lewis Dot Structure Rules Lewis Dot Structures can be used to represent covalent bonds between atoms. Let s examine the rules for drawing these structures: 1. Count the number of valence electrons in each atom in the molecule. 5

6 2. In molecules with multiple bonding atoms, decide which element should be the central element. Rule of Thumb: An atom makes the same number of bonds as the number of valence electrons it needs to complete its valence shell. The atom that makes the most number of bonds should be the central element. - Duet Rule - Octet Rule 3. Draw the central atom first with its valence electrons. Draw in the other atoms in the molecule, sharing one electron from each atom, forming single bonds. Draw each atom s lone pair electrons. 4. Count the total number of electrons drawn. Does it match your count from step # 1? 5. Check to determine if the Octet Rule is satisfied for each atom. If not, you might have to draw double or triple bonds. A bond that involves two shared pairs of electrons is a. A bond formed by sharing three pairs of electrons is a. C 4 Let s Practice N 3 C 3 O CO 2 CO 6

7 Polyatomic Ions A polyatomic ion, such as N 4+, is a covalently bonded group of atoms that has a positive or negative charge and behaves as a unit. Rule of Thumb: When counting the total number of valence electrons in the structure add or subtract electrons as the charge indicates. Draw the rest of the structure as usual. Draw CO 3 2- Draw SO 2 This leads to the following structures. These equivalent structures are called Resonance A resonance structure is a structure that occurs when it is possible to draw two or more valid electron dot structures that have the same number of electron pairs for a molecule or ion. Draw ozone: O 3 The real structure is a hybrid of all possible resonance forms. 7

8 Thiocyanate Ion, SCN - S C S N C S C N N Which is the most important resonance form? Calculate formal charge. Checking yourself Formal Charge Atoms in molecules often bear a charge (+ or -). The predominant resonance structure of a molecule is the one with charges as close to 0 as possible, and with the negative charge on the most electronegative element. Formal charge = [Total number of valence electron] [1/2 the # of bonding electrons] - [# of lone pair electrons] Carbon Dioxide, CO 2 O C O 8

9 Exceptions to the Octet Rule The octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number. There are also molecules in which an atom has fewer, or more, than a complete octet of valence electrons. Draw NO 2 Unit 5 Chemical Bonding Molecular Structure Theory VSEPR Bonding Theory According to VSEPR theory, Valence Shell Electron Pair Repulsion theory. 9

10 1. Linear: Molecular Geometries No. of e- Pairs Around Central Atom Example 2 F Be F 180 Geometry linear 3 F B F F 120 planar trigonal C tetrahedral 2. Trigonal Planar Molecular Geometries No. of e- Pairs Around Central Atom Example 2 F Be F 180 Geometry linear 3 F B F F 120 planar trigonal C tetrahedral 3. Tetrahedral: Molecular Geometries No. of e- Pairs Around Central Atom Example 2 F Be F 180 Geometry linear 3 F B F F 120 planar trigonal C tetrahedral 10

11 Structure Determination by VSEPR 4. Ammonia, N 3 1. Draw electron dot structure 2. Count BP s and LP s = areas of e- denisty 3. The 4 areas of e- density are at the corners of a tetrahedron. N lone pair of electrons in tetrahedral position Structure Determination by VSEPR The lone pair electrons take up more space. They are only attracted to the nitrogen nucleus N lone pair of electrons in tetrahedral position The MOLECULAR GEOMETRY the positions of the atoms - Bond angel is Structure Determination by VSEPR 5. Water, 2 O 1. Draw electron dot structure 2. ow many areas of e- density? 3. ow many lone pairs? O 11

12 Bond Angles Practice Predict the electron pair geometry and molecular geometry of this molecule. Formaldehyde, C 2 O O C Practice Predict the molecular geometries in this molecule around both the carbon and oxygen atoms. Methanol, C 3 O Draw the Lewis Dot Structure first. 12

13 Valence Bond Theory A bond + Cl Cl Overlap of (1s) and Cl (2p) Note that each atom has a single, unpaired electron. ow are hybridized orbitals formed?? In hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals. Linear: 1. Two pairs of bonding electrons. 2. Two atomic orbitals needed to produce two bonding orbitals ybridization Trigonal Planar: 1. Three pairs of bonding electrons. 2. Three atomic orbitals needed to produce three bonding orbitals 13

14 ybridization Tetrahedral: 1. Four pairs of bonding electrons. 2. Four atomic orbitals needed to produce four bonding orbitals Sigma Bonds Types of Bonds 2 F 2 14

15 Sigma Bonds in Methane Pi Bonds Types of Bonds Bond Properties Free rotation around C C single bond No rotation around C=C double bond 15

16 Unit 5 Chemical Bonding Polarity Polar Covalent Bonds A, known also as a polar bond, is a covalent bond between atoms in which the electrons are shared. The more electronegative atom attracts electrons more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge. +δ -δ Cl ow can you tell a bond is polar? When the atoms in a bond pull equally (as occurs when identical atoms are bonded), the bonding electrons are shared equally, and the bond is a. 16

17 Practice Classify each of the following bonds as nonpolar, polar, or ionic. 1. F-F 2. -F 3. Mg-O 4. N- 5. C- 6. -Br Dipole Moment In a, one end of the molecule is slightly negative and the other end is slightly positive. Dipole - Dipole Moment Molecules will have a dipole moment if a) b) All above do not have dipoles. 17

18 Does this molecule have a dipole moment? F B B F F F F Does this molecule have a dipole moment? Intermolecular Forces (Van der Waals Forces) Intermolecular Forces (IMF) are attractive forces BETWEEN molecules like interstates are roads that connect between states. Intermolecular attractions are weaker than either ionic or covalent bonds. These attractions are responsible for determining whether a molecular compound is a gas, a liquid, or a solid at a given temperature. 18

19 Dipole Interactions Dipole interactions occur when polar molecules are attracted to one another due to their slight opposite charges. The opposite charges align. These are the strongest among IMF because they are permanent! Most famous, and strongest is hydrogen bonding» bonds to O, N, or F Other Intermolecular Forces London Dispersion Forces, the weakest of all molecular interactions, are caused by the motion of electrons about bonding nuclei. ALL POLAR AND NONPOLAR MOLECULES EXPERIENCE DISPERSION FORCES. TIS IS TE ONLY IMF BETWEEN NONPOLAR MOLECULES. The strength of dispersion forces generally increases as the number of electrons in a molecule increases. This is a much weaker IMF than dipole interactions. It is not as constant. Dispersion Forces 19

20 Predicting State Using IMF First determine the type of IMF that exists between molecules of the following compounds. a. F 2 b. KBr c. C d. N 3 e. C