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Chapter 5

Chapter 5. Atoms and Bonding. 5.1 Atoms, Bonding and the Periodic Table. Why do chemical reactions occur?. Valence Electrons are. The electrons that have the highest energy level and held the most loosely

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Chapter 5

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  1. Chapter 5 Atoms and Bonding

  2. 5.1 Atoms, Bonding and the Periodic Table • Why do chemical reactions occur?

  3. Valence Electrons are • The electrons that have the highest energy level and held the most loosely • The # of valence electrons determines properties of the element and how it will interact with other substances

  4. Electron Dot Diagrams • A way to represent the valence electrons in an atom or compound • Each dot represents one electron

  5. Chemical Bonds and Stability • Atoms are most stable, less likely to react, when they have 8 valence electrons to complete the outer shell. (A couple elements only need 2 valence electrons to complete the outer shell). • Atoms react to become more stable in 2 ways: 1) they give up the loosely held electrons or 2) they gain electrons until it has 8 • By gaining or losing electrons, atoms form a chemical bond: a force of attraction between 2 atoms that holds them together

  6. Valence Electrons and the PT • Group # = # of valence electrons for groups IA-VIIIA. Groups 3-12 vary

  7. Period # = # of energy levels

  8. Inert Gases • Group 18 (VIIIA) elements have a full outer shell of 8 already (except He, it has 2) • Therefore, they do not gain nor lose electrons and are unreactive

  9. Halogens Group (VIIA) • 7 valence electrons. They will gain 1 electron and become negatively charged, 1-

  10. Alkali Metals Group IA • Will lose 1 electron & become positively charged, 1+

  11. Alkaline Earth Metals, Group IIA • They will lose 2 electrons and become positively charged, 2+

  12. 5.2 Ionic Bonds • Electrons transfer from the metal atom to the non-metal atom. • The atom is now called an ion: a charged particle • The metal will have a positive charge (it lost electron(s))and the non-metal will have a negative charge (it gained electron(s)).

  13. Electron Transfer • http://www.google.com/imgres?q=writing+ionic+formulas&start=174&num=10&hl=en&safe=active&tbo=d&biw=1600&bih=775&tbm=isch&tbnid=1AmQI2z3YKBJfM:&imgrefurl=http://education.fcps.org/lhs/node/1154&docid=uTK47N-WGsTFjM&imgurl=http://education.fcps.org/lhs/sites/default/files/ann.andrex/images/ionicblue%252520animation.gif&w=343&h=234&ei=qcbIUJ7sDKquiQLa4oGwCA&zoom=1&iact=hc&vpx=645&vpy=192&dur=753&hovh=185&hovw=272&tx=147&ty=120&sig=103642743859352156590&page=5&tbnh=130&tbnw=212&ndsp=45&ved=1t:429,r:87,s:100,i:265&surl=1 • http://www.youtube.com/watch?v=pLSbkcSaYsg • http://www.youtube.com/watch?v=V3i4DKQmN5c only use 2:45-3:20 • http://www.youtube.com/watch?v=i6JhcolTR6U&feature=endscreen • 1-1:00 is ionic and 1-2 is covalent

  14. Ions…+ attracted to -

  15. Ionic Bond: • An attractive force between oppositely charged ions

  16. Chemical Names and Formula • Chemical Formula: a combination of chemical symbols that show the ratio of elements in a compound • Compounds are made so that the charges on the ions balance out • Total “+” charge = total “–” charge • Subscripts are used to 1)indicate the ratio of each element in the compound 2) indicate how many of each ion is needed

  17. Writing Chemical Formulas • Write the chemical symbol and charge of each element • Crisscross the charges and make them subscripts on the opposite ion • Write the formula

  18. Sample problems Magnesium Oxide Calcium Phosphide

  19. Naming Ionic Compounds • Name the metal • Name the non-metal with an “-ide” ending

  20. Sample problems KCl Al2O3 Be3P2 • Li2S

  21. Properties of Ionic Compounds • Ionic compounds form solids by building up repeating patterns of ions. • The ions form an orderly 3D arrangement called a crystal or crystalline lattice structure • http://www.youtube.com/watch?v=yzkA11j2sr4

  22. High Melting Points • Ionic bonds are strong & require a large energy input to break the attractive forces between ions • Therefore, ionic compounds have high melting points (solid →liquid)

  23. Electrical Conductivity • When ionic compounds dissolve in water, the ions move around independently and are able to conduct an electrical current. • Liquid ionic compounds also conduct an electrical current b/c the ions are free to move around • Ionic solids on the other hand Do Not conduct electricity b/c the ions are held rigidly in place and not able to move freely

  24. Electrical Conductivity • No ions ions held ions freely moving in place

  25. 5.3 Covalent Bonds • By sharing electrons, called a covalent bond, atoms become more stable • The force that holds atoms together in a covalent bond is the attraction of each atom’s nucleus for the shared pair(s) of electrons • When electrons are shared, molecules form, a neutral group of atoms joined by covalent bonds

  26. How many bonds are made? • The # of covalent bonds formed by non-metals, equals the # of electrons needed to make 8 (Exception: H only needs 2) • Here, chlorine has 7 valence electrons. 1 more is needed to get 8, so chlorine will make 1 bond

  27. Oxygen: has 6 valence e- • It needs 2 more to get 8, so it will make 2 bonds • Notice, the total # of e- you started with, is the same as you ended up with, just arranged differently. • →→

  28. Nitrogen: has 5 needs 3 more valence e-. So, it will make 3 bonds

  29. Molecular Compounds • Composed of molecules • Compared to ionic compounds, molecular compounds: • Have low melting and boiling points • Don’t conduct electricity… b/c there are no charged particles to pass the current along • Have weaker intermolecular forces; so, less energy is needed to cause a phase change

  30. Comparing Ionic and covalent bonding • http://www.schooltube.com/video/7870b1153b034ec08d7a/

  31. Boiling Point Comparison: Covalent vs Ionic

  32. 5.4 Bonding in Metals

  33. Metals and Alloys • Most metallic objects are alloys, a mixture of 2 or more elements, with at least 1 metal • Alloys are usually stronger and less reactive than the pure metals that they come from • Alloys can retain some of the physical properties of the pure metals • Chemical properties are enhanced or reduced in alloys

  34. Common Alloys • Stainless Steel: an alloy of iron, carbon, nickel, and chromium

  35. Brass: copper and zinc

  36. Gold Jewelry: gold & copper or silver

  37. Steel: iron and carbon (& sometimes other elements)

  38. Metallic Bonding • Most metals are crystalline solids; atoms exist as closely packed positive ions surrounded by a sea of electrons

  39. Metallic Bonding • Metal ions are held in place by Metallic Bonds: an attraction between a + metal ion and the many e- around it. • The more valence electrons an atom can add to the sea of electrons, the stronger the metallic bonds

  40. Metallic Properties • The “sea of electrons” model helps explain the properties of metals

  41. Malleability and Ductility • Most metals are flexible and can be re-shaped easily without breaking. • The ions can be pushed out of position without breaking the metallic bonds • (pg 210 Fig 24)

  42. Because metal ions move easily, they can change shape • Ductility: ability to be pulled into thin strands or wire • Malleability: the ability to bepounded or rolled into thin sheets

  43. Luster: shiny and reflective • Metals absorb light that strikes their surface and then release it.

  44. Thermal and Electrical Conductivity • Metals conduct heat and electricity b/c the electrons move freely among the atoms

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