1 / 32

Basic Concepts of Chemical Bonding

Basic Concepts of Chemical Bonding. Why are some compounds composed of ions and some composed of molecules? The answer lies in the relationship between electronic structure and chemical bonding, the forces holding atoms together.

brandon-nixon
Download Presentation

Basic Concepts of Chemical Bonding

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Basic Concepts of Chemical Bonding

  2. Why are some compounds composed of ions and some composed of molecules? The answer lies in the relationship between electronic structure and chemical bonding, the forces holding atoms together.

  3. Held together by electrostatic forces between particles of opposite charge generally formed between metals and nonmetals Formed by electron transfer Formed by the sharing of electrons Generally between two nonmetals of similar electronegativity Ionic Bonds Covalent Bonds

  4. Metallic Bonds Formed between atoms of the same metal Freely moving electrons are delocalized around nuclei Lustrous, good conductors of electricity

  5. Ionic Bonds Formed when a metal with low ionization energy combines with a nonmetal of high electron affinity: Na. + :Cl. Na+ + Cl- Principal reason for ionic bond stability is the force of attraction between ions of unlike charge. A measure of that stability is the lattice energy: energy required to completely separate a mole of a solid ionic compound into its gaseous elements. Magnitude of lattice energy defined by: Eel = k Q1Q2 d

  6. Q1 and Q2 are the charges on the atoms measured in Coulombs; d represents the distance between the centers or size of the atoms, and k is a constant, 8.99 x 109 Jm/C2. Lattice energy increases as the charge on the atoms increases and decreases as the atom gets larger. Atoms should want to lose more electrons, Q1 and Q2 would be larger, but there is a balance between: metals: lattice energy and ionization energy nonmetals: lattice energy and electron affinity Transition metals lose valence shell s electrons first and then as many d electrons as need to form an ion of a particular charge. “s before d”

  7. Arrange the following compounds in order of increasing lattice energy: NaF CsI CaO The lattice energies for KF, MgO, and ScN are as follows: 808, 3795, and 7547 kJ/mol respectively. Account for the trend in lattice energy.

  8. Born-Haber Cycle Hess’s Law like series of steps that allows us to calculate Hf of a compound. Hfo = Na(s) + 1/2 Cl2(g)  NaCl(s) -410.9 kJ/mol Step 1: Generate gaseous atoms of Na and Cl Na(s)  Na(g) Hfo = 107.7 1/2 Cl2(g)  Cl(g) 121.7 Both processes are endothermic

  9. Step 2: Remove an electron from Na and add it to Cl Na(g)  Na+(g) + 1e- I1(Na) = 496 kJ/mol Cl(g) + 1e-  Cl-(g) Ea = -349 Step 3: Bring the gaseous ions together to form NaCl(s) Na+(g) + Cl-(g)  NaCl(s) H = ? Add the steps together. Determine the Born-Haber cycle for the formation of CaBr2.

  10. Ionic Size • Crucial to structure and stability of ionic solids • size determines packing and lattice energy • Cations are smaller than parent atoms, why? • Anions are larger than parent atoms, why? Arrange in order of decreasing size: Mg2+, Ca2+, Ca.

  11. Arrange the following in order of decreasing size: S2- K+ Ca2+ Cl- Sr2+ Y3+ Rb+

  12. Covalent Bonding • Forms compounds that are gases, liquids, and soft solids. • Low melting points, vaporize easily • Share a pair of electrons • Electron density concentrated between nuclei. • Lewis electron dot structures are used to represent the bonding arrangement in an ion or molecule.

  13. Lewis Structures Use NAS formula for molecules obeying octet rule. N = number of electrons needed; multiply all non-H atoms by 8; H by 2. A = number of electrons available; add up valence electrons. S = N-A represents the shared electrons. Divide this number by 2 = number of bonds that must be drawn. Draw the Lewis structure for the following: NF3 CH3Cl ClO4- CO2

  14. Bond Polarity and Electronegativity • When a covalent bond is formed, the electron pair may not be equally shared by both atoms; in this case, the bond is said to be polar. • A non-polar covalent bond is one where the atoms are shared equally. • Electronegativity is used to estimate the polarity of a bond; EN is the ability of an atom to attract electrons to itself.

  15. Pauling’s electronegativity scale was developed to quantify EN based on thermochemical data. On Pauling’s scale, F is the most EN element and has an assigned value of 4.0, Cs has the lowest value, 0.7. The trend for EN is: I The difference in EN can be used to determine bond polarity. EN difference 0.0-0.4 non-polar covalent; 0.4-1.7 polar covalent; >1.7 ionic

  16. Which bond is more polar? B - Cl C - Cl P - F P - Cl

  17. Dipole Moments • A polar molecule is one in which the centers of positive and negative charge do not coincide. • Molecules with polar bonds that are not symmetric are said to be polar and have a dipole moment > 0. • A molecule’s polarity helps determine many of its physical characteristics including state.

  18. Dipole Moments While it is not necessary to calculate dipole moments, you should be aware of the following: symbol used:   = Qr measured in Debye (D) higher the  value, more polar the molecule Molecular nomenclature: Molecules are named by exactly describing the number of atoms present using a prefix system. OF2 = oxygen difluoride

  19. Formal Charge -Means to decide between what appear to be equivalent Lewis structures. Formal charge = Number of valence electrons - (1/2 the number of bonding electrons + number of non-bonding electrons) Choose structure with smallest formal charge. NO2- N = 5- ((1/2)(6) + 2) = 0 O = 6- ((1/2)(4) + 4) = 0 O = 6- ((1/2)(2) + 6) = -1

  20. Draw Lewis structures for the following molecules, determine formal charge and select the most likely structure in each case: NNO or NON HCN or HNC NOBr or ONBr

  21. http://cheminf.cmbi.ru.nl/wetche/organic/

  22. Resonance Structures • Two alternative Lewis structures that are equivalent. • All forms are equivalent but only one form of the molecule exists that is observed experimentally. • Lewis structures are limited in describing electron distribution.

  23. Exceptions to the Octet Rule • NAS cannot be used here. • Molecules with an odd number of electrons: ClO2, NO • Molecules in which an atom has less than an octet: BF3 • Molecules in which an atom has more than an octet: PCl5, SF6. Second period elements will NEVER expand their octets as no d electrons are available.

  24. Strengths of Covalent Bonds • Stability of molecule relates to the strengths of the covalent bonds. • Use average bond energies to calculate enthalpy of reaction. • H =  Bonds Broken -  Bonds Formed • Molecules with strong chemical bonds have less tendency to undergo chemical change.

  25. Using the table of Bond enthalpies, calculate the enthalpy of reaction for the combustion of ethane. Bonds Lengths and Bond Strengths generally as the number of bonds increases, bond length decreases and bond gets stronger. Application of bond stability: explosives decompose exothermically produce large amounts of gas decompose rapidly usually produce N2, CO2 or CO

More Related