Chapter 11 The Nature of Covalent Bonding

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1 Name Chapter 11 The Nature of Covalent Bonding There are two major types of bonds 1. Ionic 2. Covalent 1. Covalent bonds electrons are shared between atoms. a. polar covalent electrons are not shared equally between the two bonded atoms. the electrons are pulled toward the more electronegative of the elements Example: HF b. nonpolar covalent electrons are shared equally between the two bonded atoms 2. Ionic bonds electrons are transferred between atoms, creating cations and anions. 1

2 9_12 IA IIA H 2.1 IIIA IVA VA VIA VIIA Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Na 0.9 Mg 1.2 VIIIB IIIB IVB VB VIB VIIB IB IIB Al 1.5 Si P 2.1 S 2.5 Cl 3.0 K 0.8 Ca 1.0 Sc 1.3 Ti 1.5 V 1.6 Cr 1.6 Mn 1.5 Fe Co Ni Cu Zn 1.6 Ga 1.6 Ge As 2.0 Se 2.4 Br 2.8 Rb 0.8 Sr 1.0 Y 1.2 Zr 1.4 Nb 1.6 Mo Tc Ru Rh Pd Ag Cd 1.7 In 1.7 Sn Sb Te 2.1 I 2.5 Cs 0.7 Ba 0.9 La Lu Hf 1.3 Ta 1.5 W 1.7 Re Os Ir Pt Au 2.4 Hg Tl Pb Bi Po 2.0 At Fr 0.7 Ra 0.9 Ac No How do you determine which type of bond will form? 1. Ionic bonds are formed when there is an electronegativity difference (ΔEN) greater than Polar covalent bonds form when there is a ΔEN between 0.5 and 2.0. a. If the ΔEN is between 1.7 and 2.0 = an ionic bond will form if a metal is one of the elements, and a polar covalent bond will form if only nonmetals or metalloids are present. 3. Nonpolar covalent bonds form when there is a ΔEN between 0 and What type of bond is formed between the following elements? 1. N and O 2. K and F 3. Mg and Cl 4. P and F 2

3 5. C and H Valence electrons - # of electrons in the highest occupied energy level of an atom. The number of valence electrons is related to the group numbers on the periodic table. Group 1 elements = 1 valence electron. Group 2 elements = 2 valence electrons. Group 13 elements = 3 valence electrons. Group 14 elements = 4 valence electrons. Group 15 elements = 5 valence electrons. Group 16 elements = 6 valence electrons. Group 17 elements = 7 valence electrons. Group 18 elements = 8 valence electrons. Lewis Dot Diagrams Steps in drawing Lewis Dot Diagrams 1. Determine the number of valence electrons in the atom or ion. 2. Write one dot on each side of the element symbol. 3. Double up until you have accounted for every valence electron. Examples: Al Br P -3 Si Ionic bonding 3

4 Determining Valence Electrons for an Ion or a Compound 1. Multiply the number of valence electrons by the number of moles of each element. 2. Add up all the electrons for each of the elements. 3. If there is a charge and it is negative, add that number of electrons to the total. 4. If there is a charge and it is positive, subtract that number of electrons from the total. 5. Total # of electrons should always be an even number! Determining Valence Electrons Examples NH 4 +1 CH 2 ClBr PO 4-3 Electron Dot Structures Valence electrons are the only electrons involved in bonding, and are the only ones written when drawing electron dot structures. In forming compounds, atoms tend to achieve the electron configuration of a noble gas, having 8 valence electrons. (Octet Rule) Exceptions to the octet rule: Less than 8 electrons H needs 2 electrons to be stable Be needs 4 electrons to be stable B needs 6 electrons to be stable Exceptions to the octet rule: more than 8 electrons Expanded valence shell Phosphorous ends in 3s 2 3p 6 (3d sublevel can also be found in energy level 3) Steps to Writing Lewis Dot Structures 1) Determine the central atom: You need a central atom when you have more than 2 atoms in a compound or molecule. The central atom is always the first atom in the formula EXCEPT if hydrogen is first. Hydrogen will never be the central atom. 4

5 2) Count the total valence electrons for the molecule: To do this, find the number of valence electrons for each atom in the molecule, and add them up. 3) Place two electrons between the central atom and all other terminal atoms to create a single bond. 4) Subtract the number of bonding electrons from the total # valence electrons: Or, in other words, subtract the number you found in #3 above from the number you found in #2 above. The answer you get will be equal to the number of bonding electrons in the molecule. 4) Add electrons around the terminal atoms to fulfill the octet rule : Remember, the exceptions to the octet rule (H, Be, and B) 5) Subtract the number of electrons you added from #4 to the total number of valence electrons: 6) If there are remaining electrons, put them on the central atom as lone pairs. Examples H 2 BCl 3 H 2 O PF 5 5

6 NF 3 SO 4-2 Single covalent bond a bond in which two atoms share a pair of electrons. Double covalent bond a bond in which two atoms share two pairs of electrons. Triple covalent bond a bond in which two atoms share three pairs of electrons. Single bonds are longer (length between the atoms) than double and triple bonds. Double bonds are longer than triple bonds. Single bonds are not as strong as double bonds, and can be broken much easier than double bonds. Triple bonds are stronger than double bonds. When do you use double or triple bonds instead of single bonds? If you have used up all of the valence electrons and you still need two more electrons to make the central atom stable, you must have one double bond. If you still need four more electrons to make the central atom stable, you must have either one triple bond or two double bonds. Double and triple bonds exist most commonly between C, N, O, and S atoms. Resonance structures: some molecules have more than one valid Lewis structure. Two or more valid Lewis structures for a molecule that differ only in the arrangement of electrons, not the arrangement of atoms, are called resonance structures. Draw Lewis structures for: N 2 6

7 NOCl HCN SiO 3-2 O 3 Bonding Theory The valence-shell electron pair repulsion (VSEPR) model predicts the shapes of molecules and ions by assuming that the valence shell electron pairs are arranged as far from one another as possible. To predict the relative positions of atoms around a given atom using the VSEPR model, you first note the arrangement of the electron pairs around that central atom. Steps in Predicting Molecular Geometry 1. Draw the Lewis structure 2. Determine how many bonding pairs are around the central atom. Count a multiple bond as one pair. 3. Determine how many lone pairs, if any, are around the central atom. 7

8 Draw the Lewis structure and determine the VSEPR Shape for NH 3 PO 4 3- CHFClBr CO 2 NOCl SiO 3-2 8

9 SF 6 Polar Bonds and Molecules Nonpolar covalent bond equal sharing of electrons between two atoms. Polar covalent bond unequal sharing of electrons between two atoms. In polar covalent bonds the electrons are pulled closer to the atom with the larger electronegativity value. a) The easiest way to determine if a molecule or ion is polar or nonpolar is to look at the central atom. b) If the central atom has lone pairs of electrons, the molecule or ion is polar. c) If the central atom does not have any lone pairs of electrons, the molecule or ion is nonpolar. Polar or Nonpolar? SO 2 PO 4-3 N 2 BrO 2-1 9

10 Polar Bonds and Molecules 1. Attractions Between Molecules 2. Molecules are attracted to one another by a variety of forces. 3. These intermolecular forces are weaker than ionic or covalent bonds. 4. These forces are responsible for whether or not a molecular compound is a solid, liquid, or a gas. Type of Forces 1. van der Waals dispersion forces weakest of all intermolecular forces a) All molecules contain dispersion forces. b) As molar mass and the number of electrons increase, dispersion forces increase. c) Halogens are the most common molecules to have dispersion forces. Fluorine is a gas, Bromine is a liquid and Iodine is a solid. 2. Dipole interactions occur when polar molecules or ions are attracted to one another. This occurs when a partial positive charge and a partial negative charge come close to each other. a) Dipole interactions are very similar to, but much weaker than ionic bonds. 3. Hydrogen bonds force exerted between a hydrogen atom bonded to an F, O, or N atom in one molecule a) Hydrogen bonds can have a great effect on the boiling point of a substance Examples CH 4 H 2 O 10

11 NH 3 SF 2 11

12 12

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