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CONCEPT: ATOMIC PROPERTIES AND CHEMICAL BONDS Before we examine the types of chemical bonding, we should ask why atoms bond at all. Generally, the reason is that ionic bonding the potential energy between positive and negative ions. Generally, the reason covalent bonds form is to follow the rule, in which the element is then surrounded by 8 valence electrons. There are three models of chemical bonding: In bonding, metals connect to non-metals. transfers an electron to the, creating ions with opposite charges that are attracted to each other. Li F Li F Li F In bonding, non-metals connect to non-metals. In it the nonmetals electron pairs between their nuclei. Cl Cl In bonding, metal atoms pool their valence electrons to form an electron sea that holds the metal-ion together Page 2
CONCEPT: CHEMICAL BONDS (PRACTICE) EXAMPLE: Describe each of the following as either a(n): atomic element, molecular element, molecular compound or ionic compound. atomic element molecular element molecular compound ionic compound a. Iodine b. NH3 c. Graphite d. Na3P e. Ag2(SO4)2 Page 3
CHEMISTRY - CLUTCH CONCEPT: THE IONIC-BONDING MODEL The central idea of ionic bonding is that the metal transfers an electron(s) to a nonmetal. The metal then becomes a(n) (positive ion). and the nonmetal becomes a(n) (negative ion). Their opposite charges cause them to combine into a crystalline solid. PRACTICE: Determine the molecular formula of the compound formed from each of the following ions. a. K+ & P3- b. Sn4+ & O2- c. Al3+ & CO32- Page 4
CONCEPT: ENERGY CONSIDERATIONS IN IONIC BONDING is the enthalpy change that occurs when 1 mol of ionic solid separates into gaseous ions. It tells us the strength of ionic interactions and has an influence in melting point, hardness, solubility and other properties. Li + (g) + F (g) LiF (s) H = 1050 kj/mol In order to calculate the energy of an ionic bond we use the following equation; Ionic Bond Energy = Radius = EXAMPLE: For each pair, choose the compound with the lower lattice energy. a. BaO or MgO b. LiCl or CaS PRACTICE 1: Choose the compound with the lower lattice energy. a. AlN or KBr PRACTICE 2: Choose the compound with the higher lattice energy. a. CsF or LiCl Page 5
CONCEPT: LATTICE ENERGY APPLICATION Lattice Energy represents the energy released when 1 mole of an ionic crystal is formed from its gaseous ions. Mg 2+ (g) + O 2 (g) MgO (s) ΔH = 3800 kj mole Cation Charge Anion Charge Lattice Energy (Electrostatic Energy) = Cation Radius + Anion Radius Increases Generally, it increases going from left to right of a period and increases going up any group because of a(n) in atomic size. ion charges and radii help to increase lattice energy. The larger the lattice energy then the stronger the ionic bond between the ions. Mg Lattice Energy O I n c r e a s e s Results in a higher boiling point and melting point for the ionic compound. EXAMPLE 1: The solubilities of CaCrO 4 and PbCrO 4 in water at 25 C are approximately 0.111 g/l H 2O and 0.0905 g/l H 2O respectively. Based on this information, which compound do you think has the smaller lattice energy? EXAMPLE 2: Which of the following bond will have the highest ionic character? A. BeBr2 B. MgBr2 C. SrBr2 D. BaBr2 Page 6
CONCEPT: BORN-HABER CYCLE The Born-Haber cycle is used a method to calculate the or of a compound. It looks mainly at the formation of an ionic compound from gaseous ions. The metal being from Groups or and the nonmetallic element being a or. M (s) + 1 2 X 2 o ΔH f MX (s) 1 = 1 3 2 = M (g) 2 X (g) 4 3 = M + (g) + X (g) HX (s) 5 4 = ΔH o f = 1 + 2 + 3 + 4 + 5 5 = Page 7
PRACTICE: BORN-HABER CYCLE EXAMPLE: Using the Born-Haber Cycle, demonstrate the formation of cesium chloride, CsCl, and calculate its heat of formation. ΔH Sublimation = 79 kj mol IE 1 = 376 kj mol ΔH Dissociation =122 kj mol EA = 349 kj mol U = 661 kj mol Page 8
CONCEPT: DIPOLE ARROWS Before drawing covalent compounds we first need to understand the idea of polarity and its connection to electronegativity. Polarity arises whenever two elements are connected to each other and there is a significant difference in their electronegativities. Generally, electronegativity going from left to right of a period and going down a group. To show this difference in electronegativity we use a dipole arrow. The dipole arrow points towards the electronegative element. The Effect of Electronegativity Difference on Bond Classification Electronegativity Difference (ΔEN) Bond Classification Example Zero (0.0) Small (0.1 0.4) Intermediate (0.4 1.7) Large (Greater than 1.7) Pure Covalent Nonpolar Covalent Polar Covalent Ionic Page 9
PRACTICE: DIPOLE ARROWS EXAMPLE: Based on each of the given bonds determine the direction of the dipole arrow and the polarity that may arise. a. H Cl b. S O c. Br B Br PRACTICE 1: Based on the given bond determine the direction of the dipole arrow and the polarity that may arise. a. H C PRACTICE 2: Based on the given bond determine the direction of the dipole arrow and the polarity that may arise. a. N F PRACTICE 3: Based on the given bond determine the direction of the dipole arrow and the polarity that may arise. a. H N H Page 10
CONCEPT: CHEMICAL BOND IDENTIFICATION PRACTICE: Answer each of the following questions dealing with the following compounds. KBr NH3 F2 CaO NaClO a. Which of the following compound(s) contains a polar covalent bond? b. Which of the following compound(s) contains a pure covalent bond? c. Which of the following compound(s) contains a polar ionic bond? d. Which of the following compound(s) contains both a polar ionic bond and a polar covalent bond? Page 11
CONCEPT: ELECTRON-DOT SYMBOLS Before we look at the first two bonding models, we have to figure out how to depict the valence electrons of bonding atoms. In the electron-dot symbol, the element symbol represents the nucleus and inner electrons, and the surrounding dots represent the electrons. EXAMPLE: Draw the electron-dot symbol for each of the following elements. 1A 2A 3A 4A 5A 6A 7A 8A Li Be B C N O F Ne It s easy to write the Lewis symbol for any Main-Group element: 1) Remember that Group Number equals Valence Electron Number. 2) Place one dot at a time on the four sides (top, right, bottom, left) of the element symbol. 3) Keep adding dots, pairing them up until you have reach the number of total valence electrons for that element. PRACTICE 1: Draw the electron-dot symbol for the following ion. Mg 2+ PRACTICE 2: Draw the electron-dot symbol for the following ion. N 3- PRACTICE 3: Draw the electron-dot symbol for the following ion. Cr 1+ Page 12
CONCEPT: CHEMICAL BONDING I Rules for Drawing 1. Least electronegative element goes into the center. Important Facts to Know: (a) Electronegativity increases across any Period going from left to right and up any Group going from bottom to top. (b) Hydrogen and Fluorine go in the center and they only make BOND. 2. Number of valence electrons equals group number. 3. Carbon must make bonds, except in rare occasions when it makes bonds. If the carbon atom were positive or negative then it would make bonds 4. Nitrogen likes to make bonds. 5. Oxygen likes to make bonds. 6. Halogens (Group 7A), when not in the center, make bond. 7. Expanded Valence Shell Theory: Nonmetals starting from Period to can have more than 8 valence electrons around them when in the center. Page 13
CONCEPT: INCOMPLETE OCTETS Nonmetals form covalent bonds to generally follow the rule, in which the element is surrounded by 8 valence electrons. Sometimes elements form compounds in which they have 8 valence electrons. These elements are said to have an incomplete octet or to be. EXAMPLE: Draw the following molecular compound. BH 3 PRACTICE: Draw the following molecular compound. BeCl 2 Page 14
CONCEPT: EXPANDED OCTETS Expanded Valence Shell Theory: Nonmetals starting from Period to can have more than 8 valence electrons around them when in the center. EXAMPLE: Draw each of the following molecular compounds. IF3 KrF5 + PRACTICE 1: Draw the following molecular compound. SBr4 PRACTICE 2: Draw the following molecular compound. I3 Page 15
CONCEPT: POLYATOMIC IONS Shortcut: If you have,,,, or connected to oxygen then the negative charge tells us how many oxygens are single bonded. The remaining oxygens are bonded to the central element. EXAMPLE: Draw each of the following molecular compounds. SO 4 2- PO 4 3- H2SO 4 PRACTICE 1: Draw the following molecular compound. SeO4 2- PRACTICE 2: Draw the following molecular compound. XeO6 4- Page 16
CONCEPT: FORMAL CHARGE Structures and polyatomic ions that break the octet rule often have Lewis Structures. The purpose of using the formal charge formula is to determine which Lewis structure is the best answer. Formal Charge = a) Use formal charge formula to check to see if you drew your compound correctly. b) Formal charges must be either,,. c) If you add up all the formula charges in your compound that will equal the overall charge of the compound. EXAMPLE: Calculate the formal charge for each of the following element designated for each of the following. a. The carbon atom in b. The sulfur atom in PRACTICE: Calculate the formal charge for each of the following element designated in the following compound. a. Both oxygen atoms in: A B! Page 17
CONCEPT: RESONANCE STRUCTURES Resonance structures are used to represent bonding in a molecule or ion when a single Lewis structure cannot correctly describe the Lewis structure. EXAMPLE: Determine all the possible Lewis structures possible for NO2. Determine its resonance hybrid. EXAMPLE: Determine the remaining resonance structures possible for the following compound, CO3 2-. O O C O Page 18