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Bonding

Bonding. Chapter 8. Types of Chemical Bonds. Ionic Bonds – metals/nonmetals Electrons are transferred Ions paired have lower energy (greater stability) than separated ions Electrostatic forces Covalent Bonds – nonmetals Electrons are shared by nuclei Pure covalent – non-polar covalent

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Bonding

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  1. Bonding Chapter 8

  2. Types of Chemical Bonds • Ionic Bonds – metals/nonmetals • Electrons are transferred • Ions paired have lower energy (greater stability) than separated ions • Electrostatic forces • Covalent Bonds – nonmetals • Electrons are shared by nuclei • Pure covalent – non-polar covalent • Electrons are shared evenly (F-F) • Polar Covalent • Electrons are shared unequally • Atoms end up with fractional charges • δ+ or δ-

  3. Covalent Bond Length • Distance at which the system energy is at a minimum • Forces at work • Attractive forces – protons and electrons • Repulsive forces – electron-electron and proton-proton • Energy is given off when two atoms achieve greater stability together than apart • Bond energy

  4. Electronegativity • The ability of atom in a molecule to attract shared electrons to itself • Trend – increases across and up

  5. Electronegativity & Bonds • Greater electronegativity difference between two elements means less covalent character and greater ionic character • Any compound that conducts an electric current when melted is an ionic compound • If the electronegativity difference < 1.67, then the atoms will share electrons.

  6. Bond Polarity & Dipole Moments • Dipolar Molecules • Molecules with a somewhat negative and a somewhat positive end • Dipole moment • Molecules with preferential orientation in an electric field • Slight negative side will be attracted to positive • All diatomic molecules with a polar covalent bond are dipolar

  7. Bond Polarity & Dipole Moments • Some molecules have polar bonds, but no dipole moment • Linear, radial, or tetrahedral symmetry of charge distribution • Charge balances/evens out CO2 CCl4

  8. Bonding & Noble Gas e- Configurations • Ionic bonds – electrons are transferred until each species attains a noble gas configuration • Covalent bonds – electrons are shared in order to complete the valence configurations of both atoms • Predicting Formulas of Ionic Compounds • Based on placement in the periodic table • Na  Na+ • Sizes of ions • Cations are smaller than parent ion • Anions are larger • Isoelectronic ions – size decreases as the nuclear charge increases

  9. Binary Compounds Lattice Energy

  10. Binary Ionic Compounds • Lattice energy – change in energy that takes place when separated gaseous ions are packed together to form an ionic solidM+(g) + X-(g)  MX(s)

  11. Determining ΔHf° • Step 1: Sublimation • Solid  Gas • Step 2: Ionization Energy • Gas  Ion • Step 3: Bond Energy • Eg: Diatomic  Single • Step 4: Electron Affinity • X + e- X- • Step 5: Formation of solid compound (LE) • Sum = ΔHf° Metal Nonmetal

  12. Example – formation of LiF

  13. Binary Ionic Compounds • The formation of ionic compounds is endothermic until the formation of the lattice • The lattice formed by alkali metals and halogens (1:1 ratio) is cubic except for cesium salts

  14. Lattice Energy Calculations • Lattice Energy = k Q1Q2r • k = constant dependent on the solid structure and the electron configurations • Q1 and Q2 = numerical ion charges • r = shortest distance between centers of the cations and the anions • Lattice Energy increases as the ionic charge increases and the distance between anions and cations decreases • Charge has more impact than distance

  15. Partial Ionic Character of Covalent Bonds

  16. Calculating Percent Ionic ( ) • % Ionic Character = measured dipole moment of X-Y x 100% Calculated dipole moment of X+Y- • Ionic compounds generally have > 50% ionic character • % ionic character is difficult to calculate for compounds containing polyatomic ions

  17. Covalent Chemical Bond • Strengths of the Bond Model • Associates the quantities of energy with the formation of bonds between elements • Allows the drawing of structures showing the spatial relationship between atoms in a molecule • Provides a visual tool to understanding chemical structure • Weaknesses of the Bond Model • Bonds are not actual physical structures • Bonds can not adequately explain some phenomena • Resonance structures

  18. Multiple Bonds • Single bonds – 1 pair of shared electrons • Double bonds – 2 pairs of shared electrons • Triple bonds – 3 pairs of shared electrons • As the number of shared electrons increases, bond length shortens • Multiple bonds typically have higher bond energy

  19. Bond Energy & Enthalpy • ΔH = sum of energies required to break old bonds (endothermic) - sum of the energies released in forming new bonds (exothermic) • ΔH = ΣD(bonds broken) – ΣD(bonds formed) • D = bond energy per mole • D always has a positive sign

  20. Localized Electron Bonding Model • Lone electron pairs • Electrons localized on an atom • Bonding electron pairs • Electrons found in the space between atoms • Shared pairs • Localized Electron Model • “A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.” • Derivations • Valence electron arrangement using Lewis structures • Prediction of molecular geometry using VSEPR • Description of the type of atomic orbitals used to share or hold lone pairs of electrons

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