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Chemical Bonding: Bonding Theory and Lewis Formulas

Chemical Bonding: Bonding Theory and Lewis Formulas. There are certain trends within the periodic table which affect reactivity and the ability to form bonds. . Periodicity. Periodic Law.

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Chemical Bonding: Bonding Theory and Lewis Formulas

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  1. Chemical Bonding: Bonding Theory and Lewis Formulas

  2. There are certain trends within the periodic table which affect reactivity and the ability to form bonds. Periodicity

  3. Periodic Law • When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals. Br Cl F

  4. Chemical Reactivity Families • Similar valence e- within a group result in similar chemical properties

  5. Chemical Reactivity • Alkali Metals • Alkaline Earth Metals • Transition Metals • Halogens • Noble Gases

  6. Properties that affect Reactivity • Atomic Radius • size of atom

  7. Atomic Radius K Halogens Na Li Ar Ne Noble Gases

  8. Atomic Radius Atomic radius increase as you move top to bottom. Atomic radius increase as you move right to left.

  9. Ionization Energy Again when you compare the ionization energy of elements there is a trend.. Ionization energy is the amount of energy needed to turn a stable element into an ion. He Ne Ar Li Na K

  10. Ionization Energy Ionization energy increases as you move right and down.

  11. Ionization Energy • Why are atomic radius and ionization energy opposite? • In small atoms, the negative electrons (e-)are closer to the positive nucleus where the attraction is stronger. 50 mμ 20 mμ Large Atoms Small Atoms

  12. Valence Electrons, Valence Energy Levels & Valence Orbitals • Periodic trends are related to the number of valence electrons an element has. • Valance electrons are those electrons occupying the highest energy level of an atom (the outside shell).

  13. How Many Valence Electrons? First Level: 2e- Second Level: 8e- Third Level: 8e- ***You can tell how many valence electrons are in each orbital by counting the number of elements on each row of the periodic table.***

  14. The maximum number of electrons in the first energy level is 2 e- • Ex. The magnesium atom has 12 protons and 12 electrons. • That leave 10 electrons left to place. • The next energy level cam only hold 8 electrons. • That only leaves 2 electrons that are TRUE valence electrons

  15. Orbital: a region in space in which an electron with a given energy is likely to be found. • There are four valence orbitals within the valence energy level of an atom (1s and 3p’s) • There are few exceptions to this rule Hydrogen Helium

  16. Electrons will occupy all valence orbitals before forming electron pairs. Empty bus seat rule • Normally a maximum of 8 electrons may occupy a valence energy level. This is known as the octet rule. • urinaltest

  17. 2 electrons in first energy level. \Leaving you with 4 more electrons to place. You place the other 4 electrons in each of the four orbital. Start at the top and go clockwise.

  18. Electron Dot Diagrams (Lewis Symbols) • Electron dot diagrams can represent atoms (neutral) or ions (charged). • ONLY show the atom’s valence electrons! – These are the only electrons involved in a chemical reaction (the electrons in the outer most ring!)

  19. Dot Diagram Trends

  20. Write the atomic symbol for the atom. This symbol represents the nucleus and the core electrons that do not participate in the chemical bonding. Dots () represent the electrons in the valence energy level of the atom. Arrange these dots around the atomic symbol. Follow these steps:

  21. One dot must be placed in each of the four orbitals before any electron pairing occurs. • Begin with the fifth electron to make lone pairs. (if you have to) • There is a maximum of 8 electrons that can be drawn.

  22. Calcium Oxygen Bromine Carbon Ca O Br C Lets try some…

  23. Bonding Electrons versus Lone Pairs • Bonding electrons are unpaired electrons that are involved in bond formation. • Paired electrons are called lone pairs and are generally not involved in bond formation.

  24. Bond Types • There are 3 types of bonds that can be formed • These are determined by which elements combine

  25. Types of Bonds 2. COVALENT 1. IONIC e- are transferred from metal to nonmetal e- are shared between two nonmetals Bond Formation Type of Structure true molecules crystal lattice Physical State liquid or gas solid Melting Point low high Solubility in Water yes usually not yes (solution or liquid) Electrical Conductivity no Other Properties odorous

  26. Metallic Bonding • 2 metals share electrons but no chemical reaction occurs • Valence electrons are free to move about between the atoms • Positive ions surrounded by a “sea” of mobile electrons • Allows metals to be formed into any shape

  27. 1. Ionic Bonding • A complete transfer of electrons occurs in an ionic bond. Valence Electrons

  28. 1. Ionic Bonding 8 8 1 7

  29. 2. Covalent Bonds

  30. 2. Covalent Bonding • Results in a mutual sharing of electrons between the two non-metals. • This Sharing can be: • Equal = nonpolar covalent • Unequal = polar covalent

  31. Nonpolar Covalent Bond • e- are shared equally between both nucleus. • Electron “Cloud” is symmetrical.

  32. - + Polar Covalent Bond • e- are shared unequally • asymmetrical e- density • results in partial charges (dipole)

  33. Polar Covalent Bond Example: H2O(l) 1.2 3.4 - 2.2 = - O H POLAR H +

  34. Bond Polarity • Most bonds are a blend of ionic and covalent characteristics. • Difference in electronegativity determines bond type.

  35. Electronegativity Electronegativity: • The attraction an atom has for a shared pair of electrons….aka….the strength an atom has to hold onto or take electrons.

  36. Trends in Electronegativity

  37. Name That Bond!!!!

  38. Bond Polarity Non-Polar ( - ) (+) Polar Ionic

  39. Bond Polarity Examples: • Cl2 • HCl • NaCl 3.2-3.2=0.0 Nonpolar 2.2-3.2=1.0 Polar 0.9-3.2=2.3 Ionic

  40. Covalent bonds are classified as: 1. single (sigma bond), 2. double (1sigma and 1pi bond), or 3. triple bonds (1sigma and 2 pi bonds) depending on the number of electrons shared between the two nuclei.

  41. Electron Dot Diagrams for Ionic Compounds • Electrons are transferred from the metal to the nonmetal. • results in a net negative charge for the nonmetal. • And a net positive charge for the metal.

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