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Chapter 7 and 8

Bonding. Chapter 7 and 8. Ch. 7 – Ionic and Metallic Bonds. Elements within each group of the periodic table behave similarly because they have same number of valence electrons. Valence electrons : the number of electrons in the highest occupied energy level. Valence Electrons.

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Chapter 7 and 8

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  1. Bonding Chapter 7 and 8

  2. Ch. 7 – Ionic and Metallic Bonds • Elements within each group of the periodic table behave similarly because they have same number of valence electrons. • Valence electrons: the number of electrons in the highest occupied energy level

  3. Valence Electrons • Key Point: the number of valence electrons = group # A • Electron Dot Diagram: Shows the number of Valence electrons • Ex. • H C O Ar

  4. Octet Rule • In forming compounds, atoms tend to achieve the electron configuration of a noble gas • To Meet the octet rule elements will form either cations (+) or anions (-) • Ex. Na O

  5. 7.2 Ionic Bonds and Ionic Compounds • Ionic compounds are composed of cations and anions, however they are electrically neutral • The electrostatic force that holds ions together in ionic compounds are called ionic bonds • Show ionic bonds using electron dot diagrams

  6. Properties • Between metals and nonmetals • Most are crystalline solids • Stable structures • High melting points

  7. 7.3 Bonding in Metals • Metallic Bonds • Attraction of free floating electrons from positive ions • Between two metals • Properties • Malleable • Ductile • Good conductors

  8. Alloys • Mixture composed of 2 more elements at least one is a metal • Properties are superior to the elements • Prepared by melting then mixing • ex: Brass is an alloy of copper and zinc

  9. Chapter 8 – Covalent Bonding • Define • Covalent Bond: the joining of atoms held together by the sharing of electrons • Molecule: a neutral group of atoms joined together by a covalent bond • Molecular Compound: a compound composed of molecules.

  10. Definitions continued • Diatomic Molecule: a molecule consisting of 2 atoms when found in nature • They include Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, Iodine and Hydrogen • Form a 7 on the periodic table • Molecular Formula: the chemical formula of a molecular compound • Shows how many atoms of each element

  11. Summary Ionic Covalent Electrons shared 2 nonmetals Does not require ions Ex. H2O • Electrons transferred (gained or lost) • A metal (+) and nonmetal (-) • Ions have to be formed • Ex. NaCl

  12. The Octet Rule • Key Point: in forming covalent bonds, electron sharing usually occurs so that atoms attain the electron configuration of a noble gas. • Lewis Structures • Used to show this in covalent bonds

  13. Lewis Structures • Shared Pair: a pair of electrons that is shared between two atoms • Represented by a dash “ “ • Unshared Pair: also know a lone pair or nonbonding pair • Represented by two dots (dots = electrons) “..”

  14. Steps for Lewis Structures • Figure out the number of electrons available for bonding. (Group # x how many atoms) • Pick the central atom. (Usually written first, never H) • Connect the other atoms using shared pair (a dashed line) • Fill outer octets first (use dots, exception H) • Place any left over on central. (Using dots)

  15. Examples • CH4 • NH3

  16. VSEPR • Valence Shell Electron Pair Repulsion • Valence shell electrons repel as far apart as possible to give molecules different geometries (shapes) • Determine the number of shapes from the number of shared and unshared pairs from Lewis Structure

  17. Molecular Geometries

  18. Polarity • Is seen in covalent compounds (shared electrons) • Depends on • Geometry (shape) • Electronegativity

  19. Electronegativity • An atoms attraction for electrons when bonded to another atom • Trend: Bottom to Top, Left to Right • Due to differences in Electronegativity, elements don’t share electrons equally so they get positive or negative areas.

  20. Polar Molecules • Key Point: If an atom has 2 ends a positive and negative, it said to be polar. • More electronegative pulls electrons in closer so it negative • Less electronegative would be positive

  21. Polarity

  22. Attractions Between Molecules • Key Point: Intermolecular attractions are weaker than either ionic or covalent bonds. • Two Main Attractions: • Van der Waals • Hydrogen

  23. Van der Waals Forces • Are the weakest between molecules • Dipole Interaction: When polar molecules are attracted to one another. The negative end is attracted to the positive end of another molecule.

  24. Van der Waals • Dispersion: The weakest • Occurs between non-polar molecules • Caused by moving electrons • Explains the diatomic molecules

  25. Hydrogen Bonds • When a hydrogen is covalently bonded to a very electronegative atom causes a weak attraction to an unshared pair of another atom • Ex. Water

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