Chemical Bonding Copyright© by Houghton Mifflin Company. All rights reserved.

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<ul><li> Slide 1 </li> <li> Chemical Bonding Copyright by Houghton Mifflin Company. All rights reserved. </li> <li> Slide 2 </li> <li> Bonding Intramolecular Bonding Within molecules Intermolecular Bonding Between molecule Copyright by Houghton Mifflin Company. All rights reserved. 2 </li> <li> Slide 3 </li> <li> Octet Rule Atoms tend to gain, lose, or share electrons in order to acquire a full set of valence electrons A very important rule An Octet consist of eight electrons This refers to the outermost eight electrons Compounds form to satisfy the Octet Rule Ionic or Covalent Bonds 3 </li> <li> Slide 4 </li> <li> Types of IntramolecularBonds Ionic Formed by transfer of electrons Held together by electrostatic attraction Covalent Formed by sharing of electrons Held together by shared electrons Metallic Formed when electron(s) become detached from metal atoms Sea of electrons among + ions metal of atoms Held together by electrostatic attraction 4 </li> <li> Slide 5 </li> <li> 5 </li> <li> Slide 6 </li> <li> Electronegativity values for selected elements. 6 </li> <li> Slide 7 </li> <li> Electronegativity Difference and Bond Type Electronegativity Difference (Approximate) Bond Type Example 0-0.4 Nonpolar Covalent H-H 0.4-1.0 Moderately Polar Covalent H-Cl 1.0-2 Very Polar Covalent H-F 2.0IonicNaCl 7 </li> <li> Slide 8 </li> <li> Ions as packed spheres. 8 </li> <li> Slide 9 </li> <li> Ionic bonding occurs when one atom transfers an electron to another atom Both atoms involved become charged Both atoms involved become charged One is negatively charged One is negatively charged One is positively charged One is positively charged This occurs when a metal reacts with a nonmetal This occurs when a metal reacts with a nonmetal 9 Na +1 + Cl -1 NaCl (table salt) Ionic Bonds </li> <li> Slide 10 </li> <li> 10 </li> <li> Slide 11 </li> <li> Ions and Charges CationsPositive Ions AnionsNegative Ions Group 1, 2, 3 Form ions with charges equal to their group number Group 5, 6, 7 Form ions with charges equal to group number minus eight 11 </li> <li> Slide 12 </li> <li> Ionic Compounds and Formulas The formula of a compound describes what elements are in the compound and in what proportions. The formula of a compound describes what elements are in the compound and in what proportions. Compounds that are held together by ionic bonds are called ionic compounds. Compounds that are held together by ionic bonds are called ionic compounds. 12 </li> <li> Slide 13 </li> <li> 13 Common polyatomic ion names FormulaName NH 4 + Ammonium ion CO 3 2- Carbonate ion PO 4 3- Phosphate ion SO 4 2- Sulfate ion OH - Hydroxide ion NO 3 - Nitrate ion </li> <li> Slide 14 </li> <li> Formation of Ionic Compounds 14 Al 3+ SO 4 2- Al 3+ SO 4 2- Al 2 (SO 4 ) 3 Al 2 (SO 4 ) 3 Ionic compounds forms such that total of ionic charges is zero. For above case: 2 x (+3) = 6 3 x (-2) = -6 Total 0 </li> <li> Slide 15 </li> <li> 15 The Lewis dot structure for Oxygen O Oxygen is in group VIA so it has 6 valence electrons </li> <li> Slide 16 </li> <li> 16 The Lewis dot structure for Chlorine Cl chlorine is in group VIIA so it has 7 valence electrons </li> <li> Slide 17 </li> <li> 17 The Lewis dot structure for calcium Ca calcium is in group IIA so it has 2 valence electrons </li> <li> Slide 18 </li> <li> 18 Making calcium chloride +Ca Cl CaCl 2 </li> <li> Slide 19 </li> <li> Covalent Bonds Covalent Bonds A covalent bond is a chemical bond that is formed when two atoms share a pair of electrons. A covalent bond is a chemical bond that is formed when two atoms share a pair of electrons. H. + H. H:H H. + H. H:H Covalent Compounds and Formulas Covalent Compounds and Formulas In the above example, each hydrogen has a filled valence shell simulating the electron configuration of helium. In the above example, each hydrogen has a filled valence shell simulating the electron configuration of helium. Compounds that are held together by covalent bonds are called covalent compounds. Compounds that are held together by covalent bonds are called covalent compounds. Covalent compounds form from atoms on the right side of the periodic table Covalent compounds form from atoms on the right side of the periodic table 19 </li> <li> Slide 20 </li> <li> Multiple Bonds. In electron dot notations, a pair of electrons can be represented by a pair of dots :. In electron dot notations, a pair of electrons can be represented by a pair of dots :. This can be a bonding pair or a lone pair (non- bonding pair). This can be a bonding pair or a lone pair (non- bonding pair). Bonding pairs can also be represented by lines connecting atoms. Bonding pairs can also be represented by lines connecting atoms. H : H = H - H H : H = H - H When one pair of electrons is shared, it is called a single bond. When one pair of electrons is shared, it is called a single bond. H-H H-H 20 </li> <li> Slide 21 </li> <li> When two pairs of electrons are shared it is called a double bond. When two pairs of electrons are shared it is called a double bond. When three pairs of electrons are shared it is called a triple bond. When three pairs of electrons are shared it is called a triple bond. 21 </li> <li> Slide 22 </li> <li> Bond Length Bond Length varies for covalent bonds between different elements See p. 187 Two important trends As one moves down a group, bond length between atoms increases FF 0.128 nm ClCl 0.198 nm BrBr 0.228 nm Multiple bonds are shorter than single bonds Example: two carbon atoms bonded together Single covalent bond0.154 nm Double covalent bond0.134 nm Triple covalent bond0.120 nm 22 </li> <li> Slide 23 </li> <li> 23 </li> <li> Slide 24 </li> <li> Exceptions to the Octet Rule Exceptions to the octet rule include those for atoms that cannot fit eight electrons, and for those that can fit more than eight electrons, into their outermost orbital. Hydrogen forms bonds in which it is surrounded by only two electrons. Boron has just three valence electrons, so it tends to form bonds in which it is surrounded by six electrons. Main-group elements in Periods 3 and up can form bonds with expanded valence, involving more than eight electrons, e.g. PF 5 and SF 6 </li> <li> Slide 25 </li> <li> Electron-Dot Notation To keep track of valence electrons, it is helpful to use electron-dot notation. Electron-dot notation is an electron-configuration notation in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the elements symbol. The inner-shell electrons are not shown. </li> <li> Slide 26 </li> <li> Lewis Structures Electron-dot notation can also be used to represent molecules. Chapter 6 The pair of dots between the two symbols represents the shared electron pair of the hydrogen-hydrogen covalent bond. For a molecule of fluorine, F 2, the electron-dot notations of two fluorine atoms are combined. </li> <li> Slide 27 </li> <li> Lewis Structures The pair of dots between the two symbols represents the shared pair of a covalent bond. Chapter 6 In addition, each fluorine atom is surrounded by three pairs of electrons that are not shared in bonds. An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom. See Example, p. 185 </li> <li> Slide 28 </li> <li> Multiple Covalent Bonds Double and triple bonds are referred to as multiple bonds, or multiple covalent bonds. (See Table 2, p. 187) In general, double bonds have greater bond energies and are shorter than single bonds. Triple bonds are even stronger and shorter than double bonds. When writing Lewis structures for molecules that contain carbon, nitrogen, or oxygen, remember that multiple bonds between pairs of these atoms are possible. Chapter 6 </li> <li> Slide 29 </li> <li> Molecular Shape Arises because electrons repulse one another Called VSEPR Valence Shell Electron Pair Repulsion VSEPR states that in a small molecule, the pairs of valence electrons are arranged as far apart from each other as possible 29 </li> <li> Slide 30 </li> <li> Summary of Molecular Shapes TypeBond Angle Unshared Pairs Example Linear180Balanced (2 each side) CO 2 Trigonal Planar 120noneBCl 3 Tetrahedral109.5noneCH 4 Pyramidal107oneNH 3 Bent105twoH2OH2O 30 </li> <li> Slide 31 </li> <li> Bent molecular structure of the water molecule. 31 </li> <li> Slide 32 </li> <li> A closer look at a water molecule (bent) 32 </li> <li> Slide 33 </li> <li> Molecular structure of methane (Tetrahedral). 33 </li> <li> Slide 34 </li> <li> Pyramidal 34 </li> <li> Slide 35 </li> <li> Molecular Orbitals When two atoms combine, the molecular orbital model assumes that their atomic orbitals overlap to produce molecular orbitals, or orbitals that apply to the entire molecule. Just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole. A molecular orbital that can be occupied by two electrons of a covalent bond is called a bonding orbital. 35 8.3 </li> <li> Slide 36 </li> <li> Molecular Orbitals Sigma Bonds When two atomic orbitals combine to form a molecular orbital that is symmetrical around the axis connecting two atomic nuclei, a sigma bond is formed. 36 8.3 </li> <li> Slide 37 </li> <li> Molecular Orbitals When two fluorine atoms combine, the p orbitals overlap to produce a bonding molecular orbital. The FF bond is a sigma bond. 37 8.3 </li> <li> Slide 38 </li> <li> Molecular Orbitals Pi Bonds In a pi bond (symbolized by the Greek letter ), the bonding electrons are most likely to be found in sausage-shaped regions above and below the bond axis of the bonded atoms. 38 8.3 </li> <li> Slide 39 </li> <li> Polarity Bonds can be polar or nonpolar Depends on electronegativity difference Molecules can be polar or nonpolar Polar molecules are called dipoles One end of a polar molecule has + charge; other end has - charge Will align in electric field Will be attracted or deflected by magnetic field Polarity of a molecule is determined by shape and the type of bonds between its atoms 39 </li> <li> Slide 40 </li> <li> 40 The chlorine atom attracts the electron cloud more than the hydrogen atom does. Why? The chlorine atom attracts the electron cloud more than the hydrogen atom does. Why? </li> <li> Slide 41 </li> <li> Polar Molecules A hydrogen chloride molecule is a dipole. 41 8.4 </li> <li> Slide 42 </li> <li> Intermolecular Forces Attraction between molecules Holds groups of molecules together Intermolecular forces are known as van der Waals forces Dipole-Dipole attraction Dispersion Forces Hydrogen Bonds 42 </li> <li> Slide 43 </li> <li> van der Waals Forces Dipole-Dipole attraction Electrostatic attraction between polar molecules Molecules line up like magnets Dispersion Forces Polarity arises due to electron imbalance Weak electrostatic forces arise Hydrogen Bonds Strong dipoles arising from covalent bonds of hydrogen atom to a very electronegative atom H electronegativity = 2.1F=4.0 Molecules line up like magnets Results in liquids with high boiling points 43 </li> <li> Slide 44 </li> <li> Probability representations of the electron sharing in HF. 44 </li> <li> Slide 45 </li> <li> Charge distribution in the water molecule. 45 </li> <li> Slide 46 </li> <li> Water molecule behaves as if it had a positive and negative end. 46 </li> <li> Slide 47 </li> <li> Polarity Formaldehyde CH 2 O Forms dipole due to high electronegativity of O Carbon Dioxide CO 2 Has polar bond but they cancel out CO 2 is nonpolar Water H 2 O Water is a bent molecule due to unshared pairs unlike CO 2 bond polarities do not cancel out Water is polar and forms dipole Why is water liquid and Carbon Dioxide a gas? 47 </li> <li> Slide 48 </li> <li> Large Molecules Examples: protein, Subunits linked together in a chain Often bend and twist to form 3D shape Chains, rings, balls 48 </li> <li> Slide 49 </li> <li> The three states of water: 49 </li> <li> Slide 50 </li> <li> 50 </li> <li> Slide 51 </li> <li> Table 12.3 51 </li> <li> Slide 52 </li> <li> Relative sizes of some ions and their parent atoms. 52 </li> <li> Slide 53 </li> <li> The three possible types of bonds. 53 </li> <li> Slide 54 </li> <li> Figure 12.1: The formation of a bond between two atoms. 54 </li> <li> Slide 55 </li> <li> Tetrahedral arrangement of electron pairs. 55 109.5 </li> <li> Slide 56 </li> <li> Tetrahedral arrangement of four electron pairs around oxygen. 56 </li> <li> Slide 57 </li> <li> 57 </li> <li> Slide 58 </li> <li> Table 12.1 58 </li> <li> Slide 59 </li> <li> Figure 12.6: Polar water molecules are strongly attracted to negative ions by their positive ends. 59 </li> <li> Slide 60 </li> <li> Figure 12.6: Polar water molecules are strongly attracted to positive ions by their negative ends. 60 </li> <li> Slide 61 </li> <li> The NH 3 molecule pyramidal structure. 61 </li> <li> Slide 62 </li> <li> Bond Energies and Bond Lengths for Single Bonds Section 2 Covalent Bonding and Molecular Compounds Chapter 6 </li> <li> Slide 63 </li> <li> Hybrid Orbitals When atoms bond, their outer orbitals become distorted Orbitals do not look like those of unbonded atom Hybrid orbitals are formed which are a cross between the bonding orbitals sp sp 2 sp 3 Orbital hybridization contributes to shape of molecule 63 </li> <li> Slide 64 </li> <li> 64 Hybrid Orbitals </li> </ul>

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