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Chemical Bonding

Chemical Bonding. Ms. Manning. Back to Compounds…. 2 Types: Covalent Compounds Formed when non-metals bond with other non-metals Ionic Compounds Formed when metals bond with non-metals. Classification. Properties of Metallic compounds. Relatively dense solids (exception Hg)

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Chemical Bonding

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  1. Chemical Bonding Ms. Manning

  2. Back to Compounds….. • 2 Types: • Covalent Compounds • Formed when non-metals bond with other non-metals • Ionic Compounds • Formed when metals bond with non-metals

  3. Classification

  4. Properties of Metallic compounds • Relatively dense solids (exception Hg) • Good conductors of heat and electricity • Lustrous when clean/ freshly cut • Strong, malleable (can be shaped) and ductile (can be drawn into a wire) • Sonorous: Ringing sound when hit • Relatively high melting and boiling points • Usually form positive ions

  5. Properties of Non-Metals • Non-lustrous • Can exist in any state - generally gases at room temperature • Brittle, non-ductile • Poor conductors of heat and electricity • Usually exist as molecules in their elemental form • Low densities, melting and boiling points. • Combine with other nonmetals to form covalent bonds • Generally form negative ions, e.g.  Cl-, SO42-, and N3-

  6. Properties of Metalloids • Generally look metallic but are brittle (not malleable or ductile) • Neither good conductors or insulators; instead they are semiconductors.

  7. Chemical Bonding • Chemical Bond = the force of attraction holding atoms or ions together • This is how compounds are made!

  8. Classifying Compounds • Ionic Compound = a pure substance formed from a metal and a nonmetal • NaCl • CaO • Molecular Compound = a pure substance formed from two or more different nonmetals • SO2 • CO2

  9. Ionic versus Molecular Compound • Electrical Conductivity = the ability of a material to allow electricity to flow through it • Ionic Compounds  conduct electricity • Molecular Compounds  DO NOT

  10. Electrolyte • Electrolyte = a substance that forms a solution that conducts electricity • Ionic compounds  form electrolytic solutions • Molecular compounds  form non-electrolytic solutions

  11. Ionic Bonding • Ions = atoms that have gained or lost electrons • Ionic Bond = the electrostatic attraction between positive and negative ions in a compound • Metals lose electrons • Non-metals gain electrons • Both form octets = MORE STABLE

  12. Ionic Bonding – Bohr Diagrams

  13. Lewis Dot Diagrams – Ionic Bonding KBr MgCl2

  14. Naming Ionic Bonding • Ionic Compounds • Metal + Non-metal • Metal name  same as on the atom name • Non-Metal  suffix “-ide” • Example: • NaCl = Sodium Chloride • LiF = Lithium Flouride • MgO = Magnesium Oxide

  15. Non-metal Suffixes • Nitrogen = Nitride • Oxygen = Oxide • Fluorine = Fluoride • Phosphorus = Phosphide • Sulfur = Sulfide • Chlorine = Chloride • Selenium = Selenide • Bromine = Bromide • Iodine = Iodide

  16. How Many Atoms in a Molecule? • Diatomic Molecules = a molecule consisting of two atoms of the same or different elements • CO • Polyatomic Molecules = a molecule consisting of more than two atoms of the same or different elements • NH3

  17. Covalent Bonding • Covalent Bond = the attractive forces between two atoms that results when electrons are shared by the atoms • A simultaneous attraction of two nuclei for a shared pair of electrons • In Lewis Diagrams – the shared pairs of electrons are shown as lines and the lone pairs as dots

  18. Octet Rule Still Applies! • The shared pair of electrons is considered to be a pair of electrons that make both atoms have an octet 8 8

  19. Lewis Dot Diagrams – Covalent Bonds

  20. The Lone Pair • Lone Pair = a pair of valence electrons not involved in bonding

  21. Bonding Capacity • Bonding Capacity = the number of electrons lost, gained or shared by an atom when it bonds chemically • Allows us to predict how many bonds an atom can form

  22. Bonding Capacity

  23. Choosing the Central Atom for Polyatomic Molecules • The central position… • Is usually occupied by the element with the highest bonding capacity • C and N are often in the central position • The least electronegative atom is usually the central atom • Hydrogen is NEVER the central atom • Oxygen and Halogens are usually not the central atom

  24. Covalent Bonds = Strong • A large amount of energy is needed to separate the atoms that make up molecules • The stronger the bond the greater the amount of energy needed to break the bond • Single bond = strong • Double bond = stronger • Triple bond = strongest

  25. Single, Double & Triple Bonds

  26. Polar Covalent Bonds • Polar Covalent Bonds = a covalent bond formed between atoms with significantly different electronegativities; a bond with some ionic characteristics • When electrons are shared between two atoms = covalent bond • In a bond between identical atoms the electrons are shared equally • In a bond between two different atoms the sharing is unequal

  27. Non-Polar versus Polar Covalent

  28. Comparison…

  29. Difference in Electronegativity… If the electronegativity difference is: • less than 0.2 = bond is pure covalent • is between 0.2 and 1.6 = bond is polar covalent • is greater than 1.7 = bond is ionic

  30. Polar Molecules • Polar Molecules = a molecule that is slightly positively charges at one end and slightly negatively charged at the other because of electronegativity differences

  31. Types of Forces • Intramolecular Force = the attractive forces between atoms and ions within a compound • Ionic • Polar Covalent • Non-polar Covalent • Intermolecular Force = the attractive force between molecules

  32. IntRA versus IntER-molecular Forces

  33. Some Intermolecular Forces • 3 major types of Intermolecular Forces: • Dipole-dipole forces • London dispersion forces • Hydrogen bonding • The first two are known as van der Waals Forces • London dispersion forces and dipole-dipole forces

  34. van der Waals Force • Dipole-dipole force = an attractive force acting between polar molecules • Attraction between oppositely charged ends of polar molecules

  35. van der Waals Force • London Dispersion force = an attractive force acting between all molecules including nonpolar molecules • A result of temporary displacements of the electron “cloud” around the atoms in a molecule resulting in a extremely short-lived dipole

  36. London Dispersion Force

  37. Hydrogen Bonding • Hydrogen Bonding = a relatively strong dipole-dipole force between a positive hydrogen atom of one molecule and a highly elecgtronegative atom (F, O or N) in another molecule

  38. Hydrogen Bonding

  39. T • M

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