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Chapter 6

Chapter 6. Chemical Bonding. Sect. 6-1: Introduction to Chemical Bonding. Chemical bond – electrical attraction between nuclei and valence electrons of different atoms that binds them together Being bonded lowers potential energy and creates more stable arrangement of matter.

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Chapter 6

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  1. Chapter 6 Chemical Bonding

  2. Sect. 6-1: Introduction to Chemical Bonding • Chemical bond – electrical attraction between nuclei and valence electrons of different atoms that binds them together • Being bonded lowers potential energy and creates more stable arrangement of matter

  3. Types of Chemical Bonding • Ionic bonding – electrical attraction between cations and anions • Electrons are transferred from one atom to another forming ions • Covalent bonding - sharing of electrons between two atoms

  4. Difference in Electronegativities of the 2 atoms determines what type of bond it will have (chart on pg. 151) • If difference is greater than 1.7, it is considered ionic, less than 1.7, covalent • Atoms from same element always covalently bond

  5. Nonpolar-covalent bond – electrons are shared equally between atoms (difference of about 0-0.3) • Polar-covalent bond – unequal sharing of electrons (difference of 0.3-1.7) • One atom is partially negative (δ-)and the other partially positive (δ+)

  6. Sect. 6-2: Covalent Bonding and Molecular Compounds • Molecule – group of atoms held together by a covalent bond • Chemical formula – indicates relative number and type of atoms in a compound • Molecular formula – chemical formula for a molecular compound • Diatomic molecule – only 2 atoms of same element

  7. For a covalent bond to form, repulsive forces between electrons or protons must balance out the attractive force between protons and electrons • Bond length – average distance between 2 bonded atoms • Bond energy – amount of energy required to break a chemical bond (equal to amount released when bond was formed) • Pg. 168

  8. Octet rule – atoms gain, lose, or share electrons in order to have 8 valence electrons • Exception – boron likes to only have 6 • Exception – some elements are surrounded by more than 8 when paired with highly electronegative elements such as F, O, & Cl (called expanded valence)

  9. Electron-dot Notation – shows valence electrons as dots around the element symbol, which represents nucleus and inside electrons • Lewis-structures – uses pairs of dots for unshared (lone pair) electrons and dashes for shared electron pairs • Structural formula – similar to Lewis-structure, but does not show lone pair electrons

  10. Steps to drawing Lewis-structures • Determine # and type of atoms in compound • Write electron dot notation for each type of atom in compound • Determine total number of valence electrons in compound

  11. Arrange atoms to form skeleton structure. Carbon, or least electronegative element goes in center. • Connect atoms by shared electron pairs. • Add unshared pairs of electrons until each atom has 8 electrons (except hydrogen which only wants 2) • Count # electrons used and be sure it matches # available, if not adjust to form multiple bonds.

  12. Multiple Covalent Bonds • Double bond – sharing of 2 pairs of electrons • Triple bond – sharing of 3 pairs of electrons • The more electron pairs shared, the shorter the bond length and higher the bond energy

  13. Resonance – bonding that allows for more than one correct Lewis structure for a compound • Represented by double-headed arrow between the 2 possible structures

  14. Sect. 6-3: Ionic Bonding and Ionic Compounds • Ionic compound – composed of cations & anions, paired so that charges balance • Formula unit – simplest ratio for which an ionic compound’s formula can be written • Use dot structure to show how electrons transfer to form ionic compounds

  15. Ionic compounds are arranged in a 3-D structure called a crystal lattice • Lattice energy – energy released when one mole on an ionic compound is formed • Negative values show that energy is released

  16. Comparison of Covalent/Ionic • Covalent – low melting/boiling because weaker bonds; soft solids, liquid, or gas; do not conduct electricity • Ionic – high melting/boiling because stronger bonds; very brittle & hard; conduct electricity in liquid state or dissolved in water, but not in solid state

  17. Polyatomic ion – a group of covalently bonded atoms with a charge • Will bond with ions of opposite charge to form ionic compounds • Lewis structures are drawn with brackets around it and the charge written outside the brackets

  18. Sect. 6-4: Metallic bonding • Metallic bonding – chemical bond that results from the attraction between a sea of electrons and metal atoms • Properties • Excellent conductor of electricity & heat • Lustrous - shiny • Malleable – hammered in to thin sheets • Ductile - drawn into wires • Heat of vaporization – used to measure strength of bonds in metals

  19. Sect. 6-5: Molecular Geometry • VSEPR Theory (Valence Shell Electron Pair Repulsion Theory) – repulsion between valence electron pairs causes them to spread as far apart as possible • Central atom with 2 atoms around it and no lone electron pairs will spread apart with 180˚ bond angles and a straight configuration • See pg. 186 for full list

  20. Unshared electron pairs act like an atom, but bond angles are slightly different • Double and triple bonds act like single bonds

  21. Hybridization – mixing of 2 or more orbitals of similar energies on the same atom to produce new orbitals of equal energy • A 2s and three 2p orbitals can combine to form four sp3 orbitals that have an energy higher than the 2s, but less than the 2p

  22. Intermolecular forces – force of attraction between molecules • Dipole-dipole • Hydrogen bonding • London dispersion forces

  23. Dipole – created by equal, but opposite charges separated by a small distance • Direction is from positive to negative, represented by an arrow pointing toward negative and tail crossed at positive end • Dipole-dipole – forces of attraction between polar molecules • Induced dipole

  24. Hydrogen bonding – very strong dipole between hydrogen bonded to a highly electronegative atom (N, O, F) attracted to an unshared pair of electrons of an electronegative nearby atom

  25. London Dispersion Forces – weak intermolecular forces caused by instantaneous dipoles • The only type of intermolecular forces acting between noble gases and non-polar molecules • Forces increase with increasing atomic mass

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