Presentation on theme: "1 Vanessa N. Prasad-Permaul Valencia Community College CHM 1045."— Presentation transcript:
1 Vanessa N. Prasad-Permaul Valencia Community College CHM 1045
2 Ionic Bonds Ionic Bonds: a chemical bond formed by the electrostatic attraction between positive and negative ions Bond forms when one or more electrons are transferred from the valence shell of one atom to the valence shell of the other. Cation: atom losing electrons Anion: atom gaining electrons 2
Lewis Electron-Dot Symbols: electrons in the valence shell of an atom or ion are represented by dots placed around the elemental symbol. Na + Cl Na + + [ Cl ] - Valence electrons are those electrons with the highest principal quantum number (n). 4 Ionic Bonds
5 Lewis Electron-Dot Symbols for Atoms in the Periodic Table of Elements
6 Ionic Bonds EXAMPLE 9.1: Use Lewis Electron-Dot symbols to represent the transfer of electrons form magnesium to fluorine atoms to form ions with noble gas configuration. Mg + Fl MgFl 2 F + Mg + F [ F ] - + Mg 2+ + [ F ] -
7 Ionic Bonds EXERCISE 9.1: Use Lewis Electron-Dot symbols to represent the transfer of electrons form magnesium to oxygen atoms to form ions with noble gas configuration.
8 Electron Configurations of Ions Common monatomic ions found in compounds of the main-group elements fall into 3 categories: 1.Cations of Groups IA to IIIA having noble gas or pseudo-noble-gas configurations ; the ion charges equal group numbers. 2.Cations of Groups IIIA to VA having the ns 2 electrons configurations; the ion charges equal the group numbers minus two. Tl +, Sn 2+, Pb 2+, Bi 3+. 3.Anions of Groups VA to VIIA having noble gas configurations; the ion charges equal the group number minus 8.
9 EXAMPLE 9.2: Write the electron configuration and Lewis symbol for N 3-. N = [He]2s 2 2p 3 + 3e - [He]2s 2 2p 6 N = N N 3- = [ N ] 3- Electron Configurations of Ions
10 EXERCISE 9.2: Write the electron configuration and Lewis symbol for Ca 2+ and S 2-. Electron Configurations of Ions
11 EXERCISE 9.3: Write the electron configuration and Lewis symbol for Pb and Pb 2+. Electron Configurations of Ions
12 Electron Configurations of Ions EXAMPLE 9.3: Write the electron configuration and Lewis symbol for Fe 2+ and Fe 3+. Fe = [Ar]3d 6 4s 2 = Fe Fe 2+ = [Ar]3d 6 = [ Fe ] 2+ Fe 3+ = [Ar]3d 5 = [ Fe ] 3+
13 Electron Configurations of Ions EXERCISE 9.4: Write the electron configuration and Lewis symbol for Mn and Mn 2+.
14 Ionic Radii Ionic Radius: A measure of the size of the spherical region around the nucleus of an ion within the electrons are most likely to be found. If the atom loses an electron, the cation will be smaller. The electron-electron repulsion is initially less. orbitals can shrink to increase the attraction of the electrons for the nucleus.
15 Ionic Radii EXERCISE 9.5: Which has the larger radius. S or S 2-. Explain.
16 Ionic Radii EXERCISE 9.6: Using only the Periodic Table, arrange the following ions in order of increasing ionic radius : Sr 2, Mg 2+, Ca 2+.
17 Ionic Radii The Comparison of Atomic and Ionic Radii DECREASES ACROSS A PERIOD INCREASESDOWN AGROUPINCREASESDOWN AGROUP ATOMIC RADIUS
18 Ionic Radii Isoelectronic: Different species having the same number and configuration of electrons. Na + < Mg 2+ < Al 3+ Decrease in atomic radius As the charge increases, the orbitals contract due to the greater attractive forces of the nucleus the ionic radius decreases with increasing atomic number. In general, across a period the cations decrease in radius. As the anions are reached, there is an abrupt increase in radius and then the radius decreases again.
19 Ionic Radii EXAMPLE 9.4: Arrange the following ions in order of decreasing ionic radius: F -, Mg 2+, O 2-. F - = 1s 2 2s 2 2p 6 Mg 2+ = 1s 2 2s 2 2p 6 O 2- = 1s 2 2s 2 2p 6 All are isoelectronic, so as the nuclear charge increases, the ionic radius decreases. O 2-, F -, Mg 2+
20 Ionic Radii EXERCISE 9.7: Arrange the following ions in order of increasing ionic radius: Cl -, Ca 2+, P 3-.
Covalent Bonds Covalent Bonds are formed by sharing at least one pair of electrons. The attraction (nucleus/electrons) outweighs the repulsions (electron/electron & nucleus/nucleus) 21
Covalent Bonds 22 Every covalent bond has a characteristic length that leads to maximum stability: Bond Length
Strength of Covalent Bonds 23 Energy required to break a covalent bond in an isolated gaseous molecule is called the bond dissociation energy. Same amount of energy released when the bond forms.
Electron-Dot Structures The electron-dot structures provide a simple, but useful, way of representing chemical reactions. Ionic: Covalent: 24
25 Covalent Bonds Coordinate covalent bond: a bond formed when both electrons of the bond are donated by one atom A + B A B H + H + + NH3 H N H H
26 Group 1A tends to lose their ns 1 valence shell electron to adopt a noble gas electron configuration. Group 2A lose both ns 2 Group 3A lose all three ns 2 np 1 Group 7A Gains one electron to attain noble gas configuration Group 8A inert, rarely lose or gain electrons Covalent Bonds OCTET RULE
Single Bonds: Double Bonds: Triple Bonds: 27 Covalent Bonds
29 Covalent Bonds ELECTRONEGATIVITY: a measure of the ability of an atom in a molecule to draw bonding electrons to itself. Bond polarity is due to electronegativity differences between atoms. Pauling Electronegativity: is expressed on a scale where F = 4.0
31 Covalent Bonds EXAMPLE 9.5: Use electronegativity values to arrange the following bonds in order of increasing polarity: P H, H O, C Cl. P H = 0.0 H O = 1.4 C Cl = 0.5 P H, C Cl, H O
32 Covalent Bonds EXERCISE 9.8: Use electronegativity values to arrange the following bonds in order of increasing polarity: C O, C S, H Br.
Drawing Lewis-Dot Structures Rule 1: Count the total valence electrons. Rule 2: Draw the structure using single bonds. Rule 3: Distribute the remaining electron pairs around the peripheral atoms. Rule 4: Put remaining pairs on central atom. Rule 5: Share lone pairs between bonded atoms to create multiple bonds. 33
Lewis-Dot Structures NH 2 F Amino Fluoride: In this molecule, nitrogen is the central atom. Rule 1: Number of electrons = 5 + (2 x 1) + 7 = 14 = 7 pairs 35 Rule 2Rule 3Rule 4
36 EXAMPLE 9.6: Sulfur dichloride SCl 2, is a red fuming liquid used in the manufacture insecticides. Write the Lewis formula for the molecule. S = 6 Cl = 7 each for a total of 20 electrons :Cl : S : Cl: :Cl S Cl: Lewis-Dot Structures ¨¨ ¨¨ ¨ ¨ ¨¨ ¨¨ ¨ ¨
37 EXERCISE 9.9: Dichlorodifluoromethane CCl 2 F 2, is a gas used as a refrigerant and aerosol propellant. Write the Lewis formula for the molecule. Lewis-Dot Structures
38 EXAMPLE 9.7: Carbonyl chloride or phosgene, COCl 2, is a highly toxic gas used as a starting material for the preparation of polyurethane plastics. What is the electron dot structure of this compound? C = 4, O = 6, Cl =7 each for a total of 24 electrons : Cl : C : Cl : :O: : Cl C Cl : :O: Lewis-Dot Structures ¨¨ ¨¨ ¨ ¨ ¨ ¨ ¨ ¨
39 EXERCISE 9.10: Write the electron-dot structure of carbon dioxide. Lewis-Dot Structures
40 EXAMPLE 9.8: Obtain the electron-dot formula for the BF 4 - ion. B = 3, F = 7 each (7 x 4) = 28 for a total 31 electrons. It is an ion with one more electron so a total of 32 electrons. : F : : F : B : F : : F : Lewis-Dot Structures ¨ ¨ ¨¨ ¨ ¨ ¨ ¨ : F : : F B F : : F : ¨ ¨ ¨¨ ¨ ¨
41 EXERCISE 9.11: Write the electron-dot structure of: A. the hydronium ion, H 3 O + B. The chlorite ion, ClO 2 - Lewis-Dot Structures
Resonance Structures When multiple structures can be drawn, the actual structure is an average of all possibilities. The average is called a resonance hybrid. A straight double-headed arrow indicates resonance. 42 OOO OOO
43 EXAMPLE 9.9: Describe the electron structure of the carbonate ion CO 3 2-, in terms of electron-dot formulas. C = 4, O = 3 x 6 = 18 for a total of 22 electrons, but it has gained two electrons so there is a total of 24 electrons. 2- 2- 2- : O : : O : : O : C C C : O : : O : : O : O : : O : O : Lewis-Dot Structures ¨¨¨¨¨¨ ¨¨
44 EXERCISE 9.12: Describe the bonding in NO 3 - using resonance formulas. Lewis-Dot Structures
Formal Charge Formal Charge: Determines the best resonance structure. We determine formal charge and estimate the more accurate representation. Formal Charge = valence e - - # of e - in a bond - (# of lone-pair e - ) 2 45
47 EXAMPLE 9.11: Write the Lewis formula that best describes the charge distribution in the sulfuric acid molecule, H 2 SO 4, according to the rules of formal charge. :O: H O S O HH :O: Formal Charge ¨ ¨ ¨ ¨ ¨ ¨ +2 ¨ O ¨ SOH ¨ ¨ O O ¨ ¨ ¨ ¨
48 Exercise 9.15:Write the Lewis formula that best describes the phosphoric acid molecule, H 3 PO 4.
Resonance Structures How is the double bond formed in O 3 ? The correct answer is that both are correct, but neither is correct by itself. 49
Example 1: Which of the following is correct? 1. Energy is absorbed to form a bond 2. Energy is released when a bond is formed 50
Example 2: Drawing Lewis-Dot Structures Draw electron-dot structures for: C 3 H 8 H 2 O 2 CO 2 N 2 H 4 CH 5 NC 2 H 4 C 2 H 2 Cl 2 CO H 3 S + HCO 3 – 51
Example 3: Formal Charge Calculate the formal charge and determine the most favorable of the following electron dot structures: SO 2 NO 3 – NCO – N 2 OO 3 CO 3 2– 52
Example 4: What is the overall formal charge of the following structure? 1. -2 2. -3 3. -1 4. 0 53
Example 5: Ionic Radii of Ions Compare ionic radii Fe & Fe 3+ Cl & Cl - 54