Week 3 bonding
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  • 1. Chapter 1 Atoms, Molecules & Chemical bonding Introduction Quantum Theory Atomic Orbitals Electronic Configuration Prepared by: Mrs Faraziehan Senusi PA-A11-7C Molecular Orbitals Bonding and Intermolecular Compounds David P. White Prentice Hall © 2003
  • 2. Lesson Plan At the end of this topic, the students will be able:     To describe atomic orbitals. To write electronic configurations To explain the bonding between different atoms To explain the interactions between molecules David P. White Prentice Hall © 2003
  • 3. Chemical Bonding • Chemical bonding refers to the attractive forces that hold atoms together in compounds. • Ionic bond results from the transfer of electrons from a metal to a nonmetal. It results from electrostatic interactions among ions. Compounds containing predominantly ionic bonding are called ionic compounds. • Covalent bond results from sharing electrons between the atoms. Usually found between nonmetals. Those that are held together mainly by covalent bonds are called covalent compounds. • Metallic bond: attractive force holding pure metals together. David P. White Prentice Hall © 2003
  • 4. David P. White Prentice Hall © 2003
  • 5. David P. White Prentice Hall © 2003
  • 6. Types of Bonds • We can classify bonds based on the kinds of atoms that are bonded together Types of Atoms metals to nonmetals nonmetals to nonmetals metals to metals Type of Bond Ionic Covalent Metallic Barbara A. Gage PGCC CHM 1010 Bond Characteristic electrons transferred electrons shared electrons pooled
  • 7. Types of Bonding Barbara A. Gage PGCC CHM 1010
  • 8. The Octet Rule • All noble gases except He has an s2p6 configuration. • Atoms with one or two valence electrons more than a closed shell are highly reactive because the extra electrons are easily removed to form positive ions. • Atoms with one or two valence electrons fewer than a closed shell are also highly reactive because of a tendency either to gain the missing electrons and form negative ions, or to share electrons and form covalent bonds. • Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs). David P. White Prentice Hall © 2003
  • 9. Ionic Bonding • Between atoms of metals and nonmetals with very different electronegativity • Bond formed by transfer of electrons • Produce charged ions all states. Conductors and have high melting point. • Examples; NaCl, CaCl2, K2O David P. White Prentice Hall © 2003
  • 10. Ionic Bonding Consider the reaction between sodium and chlorine: 2Na(s) + Cl2(g) 2NaCl(s) Hºf = -410.9 kJ David P. White Prentice Hall © 2003
  • 11. Ionic Bonding
  • 12. Ionic Bonding • The reaction is violently exothermic. • We infer that the NaCl is more stable than its constituent elements. Why? • Na has lost an electron to become Na+ and chlorine has gained the electron to become Cl . Note: Na+ has an Ne electron configuration and Cl has an Ar configuration. • That is, both Na+ and Cl have an octet of electrons surrounding the central ion. David P. White Prentice Hall © 2003
  • 13. Ionic Bonding • NaCl forms a very regular structure in which each Na+ ion is surrounded by 6 Cl ions. • Similarly, each Cl ion is surrounded by six Na+ ions. • There is a regular arrangement of Na+ and Cl in 3D. • Note that the ions are packed as closely as possible. • Note that it is not easy to find a molecular formula to describe the ionic lattice. David P. White Prentice Hall © 2003
  • 14. Ionic Bonding David P. White Prentice Hall © 2003
  • 15. Ionic Bonding
  • 16. Ionic Bonding Energetics of Ionic Bond Formation • Lattice energy: the energy required to completely separate an ionic solid into its gaseous ions. • Lattice energy depends on the charges on the ions and the sizes of the ions: Q1Q2 El d is a constant (8.99 x 10 9 J·m/C2), Q1 and Q2 are the charges on the ions, and d is the distance between ions. David P. White Prentice Hall © 2003
  • 17. Ionic Bonding Energetics of Ionic Bond Formation • Lattice energy increases as • The charges on the ions increase • The distance between the ions decreases. David P. White Prentice Hall © 2003
  • 18. The energy change associated with the loss of one mole of electrons by one mole of Na atoms to form one mole of Na+ ions= 496 kJ/mol The energy change for the gain of one mole of electrons by one mole of Cl atoms to form one mole of Cl– ions is given by the electron affinity of Cl= –349kJ/mol The strong attractive force between ions of opposite charge draws the ions together and the energy associated with this attraction is the crystal lattice energy of NaCl, = – 789 kJ/mol. David P. White Prentice Hall © 2003
  • 19. COVALENT BOND bond formed by the sharing of electrons
  • 20. Covalent Bond • Between nonmetallic elements of more or less similar electronegativity. • Formed by sharing electron pairs • Stable non-ionizing particles, they are not conductors at any state • Examples; O2, CO2, C2H6, H2O, SiC
  • 21. Bond Polarity and Electronegativity • In a covalent bond, electrons are shared. • Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons. • There are some covalent bonds in which the electrons are located closer to one atom than the other. • Unequal sharing of electrons results in polar bonds.
  • 22. NONPOLAR COVALENT BONDS when electrons are shared equally H2 or Cl2
  • 23. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons. Oxygen Atom Oxygen Atom Oxygen Molecule (O2)
  • 24. POLAR COVALENT BONDS when electrons are shared but shared unequally H2O
  • 25. - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.
  • 26. Bond Polarity and Electronegativity Electronegativity • Electronegativity: The ability of one atoms in a molecule to attract electrons to itself. • Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F). • Electronegativity increase from left to right across a period and decreases with increasing atomic number within each group.
  • 27. Bond Polarity and Electronegativity Electronegativity
  • 28. Electronegativity David P. White Prentice Hall © 2003
  • 29. Bond Polarity and Electronegativity • Difference in electronegativity is a gauge of bond polarity:  electronegativity differences around 0 result in nonpolar covalent bonds (equal or almost equal sharing of electrons);  electronegativity differences around 2 result in polar covalent bonds (unequal sharing of electrons);  electronegativity differences around 3 result in ionic bonds (transfer of electrons).
  • 30. Bond Polarity and Electronegativity • There is no sharp distinction between bonding types. • The positive end (or pole) in a polar bond is represented + and the negative pole -.
  • 31. Consider HF: • The difference in electronegativity leads to a polar bond. • There is more electron density on F than on H. • Since there are two different “ends” of the molecule, we call HF a dipole.
  • 32. Bond Polarity and Electronegativity Dipole Moments Prentice Hall © 2003 Chapter 8
  • 33. Bond Polarity and Electronegativity Dipole Moments Prentice Hall © 2003 Chapter 6 μ=qxd % = μLiH/ μ
  • 34. Strengths of Covalent Bonds • The energy required to break a covalent bond is called the bond dissociation enthalpy, D • When more than one bond is broken: CH4(g) C(g) + 4H(g) H = 1660 kJ the bond enthalpy is a fraction of H for the atomization reaction: D(C-H) = ¼ H = ¼(1660 kJ) = 415 kJ. • Bond enthalpies can either be positive or negative.
  • 35. Strengths of Covalent Bonds Bond Enthalpies and the Enthalpies of Reactions • We can use bond enthalpies to calculate the enthalpy for a chemical reaction. • We recognize that in any chemical reaction bonds need to be broken and then new bonds get formed. • The enthalpy of the reaction is given by the sum of bond enthalpies for bonds broken minus the sum of bond enthalpies for bonds formed.
  • 36. Strengths of Covalent Bonds Bond Enthalpies and the Enthalpies of Reactions • Mathematically, if Hrxn is the enthalpy for a reaction, then H rxn D bonds broken D bonds formed • We illustrate the concept with the reaction between methane, CH4, and chlorine: CH4(g) + Cl2(g) CH3Cl(g) + HCl(g) Hrxn = ?
  • 37. Strengths of Covalent Bonds CH3Cl(g) + HCl(g) CH4(g) + Cl2(g)
  • 38. Strengths of Covalent Bonds • Bond Enthalpies and the Enthalpies of Reactions In this reaction one C-H bond and one Cl-Cl bond gets broken while one C-Cl bond and one H-Cl bond gets formed. H rxn D C-H D Cl - Cl D C - Cl D H - Cl 104 kJ • • The overall reaction is exothermic which means than the bonds formed are stronger than the bonds broken. The above result is consistent with Hess’s law. Hess' Law states that the heat evolved or absorbed in a chemical process is the same whether the process takes place in one or in several steps.
  • 39. Strengths of Covalent Bonds Bond Enthalpy and Bond Length • multiple bonds are shorter than single bonds. • multiple bonds are stronger than single bonds. • As the number of bonds between atoms increases, the atoms are held closer and more tightly together.
  • 40. METALLIC BOND bond found in metals; holds metal atoms together very strongly
  • 41. Metallic Bonding • Formed between atoms of metallic elements (metal & metal) • Could be the same metal bonded together (i.e copper wire) • Or could be a mixture of different metals held together (alloy) • Alloy: metal made by combining two or more metallic elements, especially to give greater strength or to prevent corrosion • Good electricity conductor at all states, lustrous, very high melting points, malleable & ductile
  • 42. A malleable substance can be rolled or pounded into sheets. A ductile substance can be drawn into wires. Prentice Hall © 2003 Chapter 6
  • 43. Metallic Bonding Electron-Sea Model of Metallic Bonding • We use a delocalized model for electrons in a metal. – – – – The metal nuclei are seen to exist in a sea of electrons. No electrons are localized between any two metal atoms. Therefore, the electrons can flow freely through the metal. Without any definite bonds, the metals are easy to deform (and are malleable and ductile). • Problems with the electron sea model: – As the number of electrons increase, the strength of bonding should increase and the melting point should increase.
  • 44. Ionic Bond, A Sea of Electrons
  • 45. Metallic Bonding Electron-Sea Model of Metallic Bonding – group 6B metals have the highest melting points (center of the transition metals). Prentice Hall © 2003 Chapter 23
  • 46. Ketelaar Triangle Difference in electronegativity A plot of average electronegativity against electronegativity difference can be used to classify the bond type for binary compounds. Ionic bonding ~ large difference in electronegativity ~ high electronegativity of one element and low for the other one element Covalent bonding ~ small difference in electronegativity ~ high electronegativity of element between nonmetals For example: MgO ΔX=3.44 - 1.31=2.13 Xmean=2.38 MgO in ionic region Metallic bonding ~ small difference in electronegativity ~ low electronegativity of element Average electronegativity