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Chapter 10 Chemical Bonding

Chapter 10 Chemical Bonding. Atoms interact with other atoms to form molecules, this is chemical bonding Bonding theories – are models that predict how atoms bond together to form molecules. Bonding Theories are applied to design molecules that will

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Chapter 10 Chemical Bonding

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  1. Chapter 10 Chemical Bonding Atoms interact with other atoms to form molecules, this is chemical bonding Bonding theories – are models that predict how atoms bond together to form molecules Bonding Theories are applied to design molecules that will interfere with the active site of HIV-protease. This delays or inhibits the onset of AIDS. HIV-protease protease inhibitor

  2. CHAPTER OUTLINE

  3. CHEMICAL BOND • The nature and type of the chemical bond is directly responsible for many physical and chemical properties of a substance: (e.g. melting point, conductivity). • Most matter in nature is found in form of compounds: 2 or more elements held together through a chemical bond. • Elements combine together (bond) to fill their outer energy levels and achieve a stable structure (low energy). • Noble gases are un-reactive since their energy levels are complete.

  4. CHEMICAL BOND • This difference in conductivity between salt and sugar is due to the different types of bonds between their atoms. • Two common types of bonding are present: ionic & covalent. • When the conductivity apparatus is placed in salt solution, the bulb will light. • But when it is placed in sugar solution, the bulb does not light.

  5. Gilbert Newton Lewis (1875 - 1946) was a famous American physical chemist known for the discovery of the covalent bond (see his Lewis dot structures and his 1916 paper "The Atom and the Molecule") Other major contributions were his theory of Lewis acids and bases and Lewis coined the term "photon" for the smallest unit of radiant energy.

  6. The Origin of Lewis Symbols of Atoms Drawings of cubical atoms, the corners of the cube represented possible electron positions Lewis later cited these notes in his classic 1916 paper on chemical bonding, as being the first expression of his ideas.

  7. LEWIS SYMBOLS OF ATOMS • Lewis structures use Lewis symbols to show valence electrons in molecules and ions of compounds. • Lewis symbols for the first 3 periods of representative elements are shown below: • In Lewis symbols, valence electrons for each element are shown as a dot.

  8. Lewis Bonding Theory • atoms bond because it results in a more stable electron configuration • atoms bond together by either transferring or sharing electrons so that all atoms obtain an outer shell with 8 electrons • Octet Rule • there are some exceptions to this rule – the key to remember is to try to get an electron configuration like a noble gas

  9. •• •• Li• Li+1:F: [:F:]-1 • •• Lewis Symbols of Ions • Cations have Lewis symbols without valence electrons • Lost in the cation formation • Anions have Lewis symbols with 8 valence electrons • Electrons gained in the formation of the anion

  10. Ionic Bonds • metal to nonmetal • metal loses electrons to form cation • nonmetal gains electrons to form anion • ionic bond results from + to - attraction • larger charge = stronger attraction • smaller ion = stronger attraction • Lewis Theory allow us to predict the correct formulas of ionic compounds

  11. Metal Nonmetal IONIC BOND • After bonding, each atom achieves a complete shell (noble gas configuration). • Ionic bonds occur between metals and non-metals. • Ionic bonds occur when electrons are transferred between two atoms.

  12. IONIC BOND • Atoms that lose electrons (metals) form positive ions (cations). • Atoms that gain electrons (non-metals) form negative ions (anions). • The smallest particles of ionic compounds are ions (not atoms). Anion Cation

  13. ∙ ∙ ∙ ∙ ∙ Ca Ca Cl Cl Cl ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ ∙ Using Lewis Theory to Predict Chemical Formulas of Ionic Compounds Predict the formula of the compound that forms between calcium and chlorine. Draw the Lewis dot symbols of the elements Transfer all the valance electrons from the metal to the nonmetal, adding more of each atom as you go, until all electrons are lost from the metal atoms and all nonmetal atoms have 8 electrons Ca2+ CaCl2

  14. Covalent Bonds • typical of molecular species • atoms bonded together to form molecules • strong attraction • sharing pairs of electrons to attain octets • molecules generally weakly attracted to each other • observed physical properties of molecular substance due to these attractions

  15. COVALENT BOND • Covalent bonds form when electrons are shared between two atoms. • Covalent bonds form between twonon-metals. • The smallest particles of covalent compounds are molecules. Electrons shared

  16. •• •• • • •• F F •• •• •• •• •• •• •• •• F F F F •• •• Single Covalent Bonds • two atoms share one pair of electrons • 2 electrons • one atom may have more than one single bond •• • • • • H H O •• •• •• H H •• O ••

  17. •• •• • • • • O O •• •• •• •• O •• •• •• •• O O O Double Covalent Bond • two atoms sharing two pairs of electrons • 4 electrons • shorter and stronger than single bond

  18. •• •• • • • • N N • • N N •• •• •• •• •• N N Triple Covalent Bond • two atoms sharing 3 pairs of electrons • 6 electrons • shorter and stronger than single or double bond

  19. POLAR & NON-POLARBONDS Electrons shared equally • Two types of covalent bonds exist: • Non-polar covalent bonds occur between similar atoms. Polar & Nonpolar • In these bonds the electron pair is shared equally between the two protons.

  20. HF POLAR & NON-POLARBONDS • Polar covalent bonds occur between different atoms. • In these bonds the electron pair is shared unequally between the two atoms. • As a result there is a charge separation in the molecule, and partial charges on each atom. + 

  21. Dipole Moments • A dipole is a material with positively and negatively charged ends • Polar bonds or molecules have one end slightly positive, d+; and the other slightly negative, d- • not “full” charges, come from nonsymmetrical electron distribution • Dipole Moment, m, is a measure of the size of the polarity • measured in Debyes, D

  22. ELECTRONEGATIVITY • Electronegativity (E.N.) is the ability of an atom involved in a covalent bond to attract the bonding electrons to itself. • Linus Pauling derived a relative Electronegativity Scale based on Bond Energies. F 4.0 Cs 0.7 Most electronegative Least electronegative

  23. ELECTRONEGATIVITY Electronegativity increases

  24. BOND POLARITY &ELECTRONEGATIVITY Polarity is a measure of the inequality in the sharing of bonding electrons The more different the electronegativity of the elements forming the bond The larger the electronegativity difference(EN) The more polar the bond formed

  25. As difference in electronegativity increases Bond polarity increases POLARITY &ELECTRONEGATIVITY Most polar Least polar

  26. POLARITY &ELECTRONEGATIVITY EN = 0 Non-polar covalent 0 < EN <1.7 Polar covalent 1.7 < EN Ionic

  27. H H Electronegativity 2.1 Electronegativity 2.1 Hydrogen Molecule POLARITY &ELECTRONEGATIVITY The molecule is nonpolar covalent EN = 0

  28. + - H Cl Electronegativity 2.1 Electronegativity 3.0 Hydrogen Chloride Molecule POLARITY &ELECTRONEGATIVITY The molecule is polar covalent EN = 0.9

  29. Na+ Cl- Electronegativity 0.9 Electronegativity 3.0 Sodium Chloride POLARITY &ELECTRONEGATIVITY No molecule exists The bond is ionic EN = 2.1

  30. SUMMARYOF BONDING Ionic Bond (large EN) EN > 1.7 Non-polar (similar electronegativities) EN = 0 Polar (moderate EN) Covalent Bond (small to moderate EN) 0 < EN < 1.7

  31. Bonding & Lone Pair Electrons • Electrons that are shared by atoms are called bonding pairs • Electrons that are not shared by atoms but belong to a particular atom are called lone pairs • also known as nonbonding pairs

  32. LEWIS STRUCTURES • In a Lewis structure, a shared electron pair is indicated by two dots between the atoms, or by a dash connecting them. • Unshared pairs of valence electrons (called lone pairs) are shown as belonging to individual atoms or ions.

  33. LEWIS STRUCTURES • Structures must satisfy octet rule (8 electrons around each atom). • Hydrogen is one of the few exceptions and forms a doublet (2 electrons). • Covalent molecules are best represented with electron-dot or Lewis structures.

  34. LEWISSTRUCTURES • Bonding electrons can be displayed by a dashed line. • Non-bonding electrons must be displayed as dots.

  35. Polyatomic Ions • The polyatomic ions are attracted to opposite ions by ionic bonds • Form crystal lattices • Atoms in the polyatomic ion are held together by covalent bonds

  36. Lewis Formulas of Molecules • shows pattern of valence electron distribution in the molecule • useful for understanding the bonding in many compounds • allows us to predict shapes of molecules • allows us to predict properties of molecules and how they will interact together

  37. LEWISSTRUCTURES • More complex Lewis structures can be drawn by following a stepwise method: 1. Count the number of electrons in the structure. • Draw a skeleton structure. • - most metallic element generally central • - halogens and hydrogen are generally • terminal • - many molecules tend to be symmetrical • - in oxyacids, the acid hydrogens are attached to an oxygen

  38. LEWISSTRUCTURES • More complex Lewis structures can be drawn by following a stepwise method: 3. Connect atoms by bonds (dashes or dots). 4. Distribute electrons to achieve Octet rule. 5. Form multiple bonds if necessary.

  39.     Example 1: Write Lewis structure for H2O Step 1: H2O = 8 electrons 2 (1) + 6 = 8 Step 2: H O H Step 3: Skeleton structure should be symmetrical Hydrogen has doublet 4 electrons used4 electrons remaining Octet rule is satisfied Step 4:

  40. Example 2: Write Lewis structure for CO2 Step 1: CO2 = 16 electrons 4 + 2(6) = 16 Step 2:     O C O   Step 3:       Skeleton structure should be symmetrical Step 4: Octet rule is NOT satisfied 10 electrons used6 electrons remaining 4 electrons used12 electrons remaining Octet rule is satisfied Step 5:

  41. Writing Lewis Structures forPolyatomic Ions • the procedure is the same, the only difference is in counting the valence electrons • for polyatomic cations, take away one electron from the total for each positive charge • for polyatomic anions, add one electron to the total for each negative charge

  42.                 Example 3: Write Lewis structure for CO32- Step 1: CO32- = 24 electrons 4+3(6)+2 = 24 Step 2: 0 electrons remaining 18 electrons remaining 12 electrons remaining 6 electrons remaining O C O O Step 3:   Step 4: Step 5: Octet rule is NOT satisfied Octet rule is satisfied

  43. Example 4: Determine if each of the following Lewis structures are correct or incorrect. If incorrect, rewrite the correct structure. Structure has 14 electrons Only 12 electrons shown 2(5) + 4(1) = 14 2 4 2 Structure is incorrect 2 2 Octets are complete

  44. Exceptions to the Octet Rule • H & Li, lose one electron to form cation • Li now has electron configuration like He • H can also share or gain one electron to have configuration like He • Be shares 2 electrons to form two single bonds • B shares 3 electrons to form three single bonds • expanded octets for elements in Period 3 or below • using empty valence d orbitals • some molecules have odd numbers of electrons • NO

  45. Some molecules, such as SF6 and PCl5 have more than 8 electrons around a central atom in their Lewis structure. SF6 and PCl5 can violate the octet rule through the use of empty d orbitals: both S and P can utilize empty d orbitals to hold pairs of electrons that help bond halogen atoms.

  46. Resonance • we can often draw more than one valid Lewis structure for a molecule or ion • in other words, no one Lewis structure can adequately describe the actual structure of the molecule • the actual molecule will have some characteristics of all the valid Lewis structures we can draw

  47. Resonance • Lewis structures often do not accurately represent the electron distribution in a molecule • Lewis structures imply that O3 has a single (147 pm) and double (121 pm) bond, but actual bond length is between, (128 pm) • Real molecule is a hybrid of all possible Lewis structures • Resonancestabilizes the molecule • maximum stabilization comes when resonance forms contribute equally to the hybrid

  48. Resonance • we can often draw more than one valid Lewis structure for a molecule or ion • Real molecule is a hybrid of all possible Lewis structures represents resonance structures The three oxygens are chemically equivalent, so it makes no difference to the ion which oxygen assumes the double bond.

  49. MOLECULARSHAPES • The three-dimensional shape of the molecules is an important feature in understanding their properties and interactions. • All binary molecules have a linear shape since they only contain two atoms. • More complex molecules can have various shapes (linear, bent, etc.) and need to be predicted based on their Lewis structures. • A very simple model , VSEPR (Valence Shell Electron Pair Repulsion) Theory, has been developed by chemists to predict the shape of large molecules based on their Lewis structures.

  50. MOLECULARSHAPES • Based on VSEPR, the electron pair groups in a molecule will repel one another and seek to minimize their repulsion by arranging themselves around the central atom as far apart as possible. • Electron pair groups can be defined as any one of the following: bonding pairs non-bonding pairs multiple bonds

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