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Chemical Bonding

Chemical Bonding. Pre-AP Chemistry Modern Chemistry Chapter 6 Created By Laura Peck, 2011. A chemical bond is a mutual attraction between the nuclei and valence electrons of different atoms that binds the atoms together.

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Chemical Bonding

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  1. Chemical Bonding Pre-AP Chemistry Modern Chemistry Chapter 6 Created By Laura Peck, 2011

  2. A chemical bond is a mutual attraction between the nuclei and valence electrons of different atoms that binds the atoms together • Ionic Bonding: bonding that results from the electrical attraction between cations and anions. (far left bonding with far right) • Covalent Bonding: sharing of electron pairs between two atoms.

  3. Ionic or covalent? Nonpolar covalent: a covalent bond in which the bonding electrons are shared equally. (generally only found in pure elemental molecules. H2, O2, etc…..) Polar: Electrons are shared unequally, so there are charges at either end of molecule. Polar-covalent bond: a covalent bond in which there is an unequal attraction to electrons.

  4. Electronegativity differences • If B is slightly more electronegative than A • B will attract the electron pair rather more than A does. • That means that the B end of the bond has more than its fair share of electron density and so becomes slightly negative. • At the same time, the A end (rather short of electrons) becomes slightly positive. In the diagram, "" (read as "delta") means "slightly" - so + means "slightly positive".

  5. Electronegativity + – 0 0 H Cl H H

  6. Calculating EN differences (page 176) • The first step in defining the polarity of a bond is to calculate electronegativity difference ( EN) •  EN = EN large - EN small • E.g. for NaCl,  EN = 2.9 - 1.0 = 1.9 • Next, estimate from fig 7.12 the % ionic character: about 65% (60 - 70%) Q-Give the % ionic character forMgO,CH,HCl MgO = 3.5 - 1.3 = 2.2 … 80% ionic (75-85) CH = 2.5 - 2.1 = 0.4 … 7% ionic (5 - 10) HCl = 2.9 - 2.1 = 0.8 … 20 % ionic (15-25) Note if % ionic is 20%, then % covalent is 80%

  7. Defining polarity • For our purposes we will define polarity in the following fashion: 0-10 % is non-polar, 10-50% is polar (covalent), 50%+ is ionic • This is a crude estimate. In reality, the only non-polar bond between 2 atoms occurs in diatomic molecules (O2:  EN = 3.5 - 3.5 = 0) Q - what is the polarity of the bonds in MgO, CH, HCl? MgO = 80% ionic = ionic CH = 7% ionic = non-polar HCl = 20 % ionic = polar covalent

  8. 6.2 Covalent Bonding and Molecular Compounds • Molecule: a neutral group of atoms that are held together by covalent bonds. • Molecular Compound: a chemical compound whose simplest units are molecules. • Chemical Formula: Shows the relative numbers of atoms of each kind in a chemical compound using Symbols and Subscripts. • Ex. H2O2, CO2, C3H6, CH3COOH • Molecular Formula: shows the types and numbers of atoms combined in a single molecule of a particular molecular compound.

  9. Formation of a Covalent Bond

  10. Characteristics of the Covalent Bond • Bond Energy(enthalpy): energy required to break a chemical bond and form neutral isolated atoms. • Bond Length: average distance between two bonded atoms at their minimum potential energy.

  11. Exceptions to the Octet rule • What are the four elements that only have s orbitals in their outermost shell? • Boron: Tends to form bonds in which it is surrounded by just six electrons. • Expanded Valence: Elements that can be surrounded by more than 8 valence electrons by involving electrons in their d orbitals. • Phosphorus & Sulfer

  12. Lewis Structures • Place pairs of electrons around an imaginary square. Use only the valence electrons.

  13. Practice on the Whiteboards. • Draw the following lewis structures: 1. Si 2. C 3. I 4. K 5. Mg

  14. Showing bonding with Lewis structures.

  15. Bond types • Lewis Structures: formulas in which atomic symbols represent nuclei and inner-shell electrons; dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds. Unshared pairs are lone pairs. • Structural formula: indicates the kind, number, arrangement, and bonds – but not the unshared pairs. • Single Bonds: One electron from each atom is shared. Represented by a single line. These are the weakest/longest bond. • Double Bonds: two electrons from each atom are shared. Represented by an ‘equal sign’. These are stronger/shorter than single bonds. • Triple Bonds: three electrons from each atom are shared. Represented by three stacked lines. These are the strongest/shortest bonds of all. • Quadruple Bonds: Theoretically possible between two carbon atoms. But do to repulsive forces cannot actually exist. (favorite test question!)

  16. Resonance structures • Same chemical formula – different physical structures

  17. Ionic Bonds • In chemical bonds, atoms can either transfer or share their valence electrons. In the extreme case where one or more atoms lose electrons and other atoms gain them in order to produce a noble gas electron configuration, the bond is called an ionic bond. • Typical of ionic bonds are those in the alkali halides such as sodium chloride, NaCl • Ionic bonding can be visualized with the aid of Lewis diagrams

  18. Polyatomic Ions • Polyatomic Ion Rap • One or more elements combine to form a charged molecule. • Positive or negative sign is placed outside the brackets. • Superscript # indicates how many excess or deficient electrons the molecule has. • Mrs. Peck will draw these on the board, copy them down: • NH4+, NO3-, SO42-, PO42-

  19. Metallic Bonding • Electron Sea Model for Metallic Bonding • To account for the bonding in metals, Lorentz proposed a model known as electron gas model or electron sea model. • Metals have vacant orbitals due to their ability to hybridize lower and higher orbitals • For example, lithium {(Li, Z = 3) 1s22s1} has 2p-orbitals vacant; • Sodium {(Na, Z = 11) 1s22s22p6 3s1} has 3p-and 5d-orbitals vacant; • Magnesium {(Mg, Z = 12) 1s22s22p6 3s2} has 3p-and 3d-orbitals vacant • The important features of electron sea model are: • The positively charged nuclei of metal atoms are arranged in a regular fashion in a metallic lattice. • Loosely held valence electrons, surround each nuclei in metallic lattice. The valence electrons enjoy complete freedom in the metallic lattice and are regarded as mobile electrons. • Thus, the simultaneous force of attraction between the mobile electrons and the positive nuclei that binds the metal atoms together, is known as metallic bond

  20. Malleability: ability of a substance to be Hammered or beaten into thin sheets. Ductility: ability of a substance to be drawn, pulled or extruded though a small opening to produce a wire. Enthalpy of Vaporization: the amount of energy as heat required to vaporize the metal is a measure of the strength of the bonds that Hold the metal together.

  21. Brain BreakVSEPR intro

  22. VSEPR Theory(Not in Chem 1 standards – but since this is honors and pre-AP – we’re going to learn it!) • VSEPR: “Valence Shell Electron Pair Repulsion” • Repulsion between valence electrons causes them to be oriented as far apart as possible. • Predicts the molecular shape of a bonded molecule • Electrons around the central atom arrange themselves as far apart from each other as possible • Unshared pairs of electrons (lone pairs) on the central atom repel the most • So only look at what is connected to the central atom

  23. Linear • 2 atoms attached to center atom • 0 unshared pairs (lone pairs) • Bond angle = 180o • Type: AB2 • Ex. : BeF2

  24. Trigonal Planar • 3 atoms attached to center atom • 0 lone pairs • Bond angle = 120o • Type: AB3 • Ex. : AlF3

  25. Tetrahedral • 4 atoms attached to center atom • 0 lone pairs • Bond angle = 109.5o • Type: AB4 • Ex. : CH4

  26. Trigonal Bipyramidal • 5 atoms attached to center atom • 0 lone pairs • Bond angle = • equatorial -> 120o • axial -> 90o • Type: AB5 • Ex. : PF5

  27. Octahedral • 6 atoms attached to center atom • 0 lone pairs • Bond angle = 90o • Type: AB6 • Ex. : SF6

  28. These are for molecules with both paired and unshared (lone) pairs of electrons around the central atom.

  29. Bent • 2 atoms attached to center atom • 2 lone pairs • Bond angle = 104.5o • Type: AB2E2 • Ex. : H2O

  30. Trigonal Pyramidal • 3 atoms attached to center atom • 1 lone pair • Bond angle = 107o • Type: AB3E • Ex. : NH3

  31. Hybridization: mixing of two or more Atomic orbitals of similar energies on the Same atom to produce new hybrid orbitals Of equal energies (video) Hybrid Orbitals: Orbitals of equal Energy produced By the combination Of two or more Orbitals on the Same atom

  32. Second part of Chapter 6 is In the Intermolecular Forces ppt! http://www.sciencecartoonsplus.com/gallery/chemistry/index.php

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