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Chemical Bonds and Electron Shells: Understanding Atom Behavior

Learn about chemical bonds, electron shells, and the importance of electrons in atoms. Explore ionic and covalent bonds and understand how ions are formed.

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Chemical Bonds and Electron Shells: Understanding Atom Behavior

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  1. Chemical Bonds

  2. Atom – the smallest unit of matter “indivisible” Helium atom

  3. electron shells • Atomic number = number of Electrons • Electrons vary in the amount of energy they possess, and they occur at certain energy levels or electron shells. • Electron shells determine how an atom behaves when it encounters other atoms

  4. Electrons are placed in shells according to rules: • The 1st shell can hold up to two electrons, and each shell thereafter can hold up to 8 electrons.

  5. Octet Rule = atoms tend to gain, lose or share electrons so as to have 8 electrons Gain 4 electrons • C would like to • N would like to • O would like to Gain 3 electrons Gain 2 electrons

  6. Why are electrons important? • Elements have different electron configurations • different electron configurations mean different levels of bonding

  7. Electron Dot Structures Symbols of atoms with dots to represent the valence-shell electrons 1 2 13 14 15 16 17 18 H He:      LiBe B  C  N  O : F :Ne :            Na Mg AlSiPS:Cl  :Ar :    

  8. Chemical bonds: an attempt to fill electron shells • Ionic bonds • Covalent bonds • Metallic bonds

  9. Learning Check  A. X would be the electron dot formula for 1) Na 2) K 3) Al   B.  X  would be the electron dot formula  1) B 2) N 3) P

  10. IONIC BONDbond formed between two ions by the transfer of electrons

  11. Formation of Ions from Metals • Ionic compounds result when metals react with nonmetals • Metals loseelectrons to match the number of valence electrons of their nearest noble gas • Positive ionsform when the number of electrons are less than the number of protons Group 1 metals ion 1+ Group 2 metals ion 2+ • Group 13 metals ion 3+

  12. Formation of Sodium Ion Sodium atom Sodium ion Na  – e Na + 2-8-1 2-8 ( = Ne) 11 p+ 11 p+ 11 e- 10 e- 01+

  13. Formation of Magnesium Ion Magnesium atom Magnesium ion  Mg  – 2e Mg2+ 2-8-2 2-8 (=Ne) 12 p+ 12 p+ 12 e- 10 e- 0 2+

  14. Some Typical Ions with Positive Charges (Cations) Group 1 Group 2 Group 13 H+ Mg2+ Al3+ Li+ Ca2+ Na+ Sr2+ K+ Ba2+

  15. Learning Check A. Number of valence electrons in aluminum 1) 1 e- 2) 2 e- 3) 3 e- B. Change in electrons for octet 1) lose 3e- 2) gain 3 e- 3) gain 5 e- C. Ionic charge of aluminum 1) 3- 2) 5- 3) 3+

  16. Solution A. Number of valence electrons in aluminum 3) 3 e- B. Change in electrons for octet 1) lose 3e- C. Ionic charge of aluminum 3) 3+

  17. Learning Check Give the ionic charge for each of the following: A. 12 p+ and 10 e- 1) 0 2) 2+ 3) 2- B. 50p+ and 46 e- 1) 2+ 2) 4+ 3) 4- C. 15 p+ and 18e- 2) 3+ 2) 3- 3) 5-

  18. Ions from Nonmetal Ions • In ionic compounds, nonmetals in 15, 16, and 17 gain electrons from metals • Nonmetal add electrons to achieve the octet arrangement • Nonmetal ionic charge: 3-, 2-, or 1-

  19. Fluoride Ion unpaired electron octet 1 - : F  + e: F :  2-7 2-8 (= Ne) 9 p+ 9 p+ 9 e- 10 e- 0 1 - ionic charge

  20. Ionic Bond • Between atoms of metals and nonmetals with very different electronegativity • Most ionic compounds are called salts • Bond formed by transfer of electrons • Produce charged ions all states. Conductors and have high melting point. • Examples; NaCl, CaCl2, K2O

  21. Ionic Bond • Active metals readily loses its valence electrons, usually to a nonmetal atom • Oppositely changed ions are formed by this process of transferring electrons • Cation: positively charged ion • Anion: negatively charged ion

  22. Ionic Bond • Types of ions: • Monoatomic: ion formed from one atom • Polyatomic: ion made up of 2 or more atoms bonded together that acts like a single unit • Ionic Compound: composed of positive and negative ions so that the positive and negative charges equal in number

  23. Ionic Bonds: One Big Greedy Thief Dog!

  24. 1). Ionic bond – electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes (Na+) and the Cl becomes (Cl-), charged particles or ions.

  25. Properties of Ionic Compounds • Strong Bond • Form Crystal Lattice • Melting Point, Boiling Point, Hardness  due to strong bonds tend to be high • Good conductors of electricity when dissolved in water • Form electrolytes • Formation of ionic bonds is always exothermic

  26. Lattice energy • Energy required to separate 1 mol of ions in ionic cmpd. (Strength holding ions in place) • The more negative the lattice energy the stronger the force of attraction • Related to the size of the ions bonded (LiCl has higher lattice energy than LiBr since atoms are closer together)

  27. COVALENT BONDbond formed by the sharing of electrons

  28. Covalent Bond • Between nonmetallic elements of similar electronegativity. • Formed by sharing electron pairs • Stable non-ionizing MOLECULES, they are not conductors at any state • Examples; O2, CO2, C2H6, H2O, SiC

  29. Covalent Bond • May share 1 or more pairs of electrons • Single bond- one pair of shared electrons between 2 atoms • Double bond- two pairs of shared electrons between 2 atoms • Triple bond- three pairs of shared electrons between 2 atoms

  30. Covalent Bond • Atoms that are bonded covalently form stable particles called molecules. • Ex: H20, CO2, and C6H12O6 • 7 Diatomic molecules to know – H2, N2, O2, F2, Cl2, Br2, and I2 • Molecular compound: 2 or more atoms are covalently bonded, lowering their potential energy, than its individual atoms

  31. Formation of Covalent Bond • These bonds form due to ATTRACTIVE and REPULSIVE forces

  32. Covalent Bonds

  33. Bonds in all the polyatomic ions and diatomics are all covalent bonds

  34. NONPOLAR COVALENT BONDS when electrons are shared equally H2 or Cl2

  35. 2. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons. Oxygen Atom Oxygen Atom Oxygen Molecule (O2)

  36. POLAR COVALENT BONDS when electrons are shared but shared unequally H2O

  37. Polar Covalent Bonds: Unevenly matched, but willing to share.

  38. - water is a polarmolecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

  39. METALLIC BONDbond found in metals; holds metal atoms together very strongly

  40. Metallic Bond • Formed between atoms of metallic elements • Electron cloud around atoms • Good conductors at all states, lustrous, very high melting points • Examples; Na, Fe, Al, Au, Co

  41. Metallic Bonds: Mellow dogs with plenty of bones to go around.

  42. Metallic Bonds, A Sea of Electrons

  43. Metals Form Alloys Metals do not chemically combine with metals. They form alloys which is a solution of a metal in a metal. Examples are steel, brass, bronze and pewter.

  44. Formula Weights • Formula weight is the sum of the atomic masses. • Example- CO2 • Mass, C + O + O 12.011 + 15.994 + 15.994 43.999

  45. Practice • Compute the mass of the following compounds round to nearest tenth & state type of bond: • NaCl; • 23 + 35 = 58; Ionic Bond • C2H6; • 24 + 6 = 30; Covalent Bond • Na(CO3)2; • 23 + 2(12 + 3x16) = 123; Ionic & Covalent

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