IONIC BONDING

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• CHAPTER 15
• IONIC BONDING

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• VALENCE ELECTRONS
• Knowing the electron configuration of elements
  will help you understand chemical bonding
  tendencies.
• The reason that elements share chemical
  properties within a given family is due to the
  similarity of their valence electrons.
• Valence electrons are the electrons that occupy
  the highest energy level.
• Valence electrons can simply be determined by
  looking at the group that they are in.
• All Group 1A elements have only one valence
  electron in their s orbital.

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• Valence electrons are usually the only electrons
  used in chemical bond formation.
• Electron dot structures are used to represent
  valence electrons of an atom.
• Table 15.1 on page 414 demonstrates a few
  elements and their dot strutures.
• All the elements within a group have the same
  number of dots, except for helium which only has
  2, as opposed to 8 for all the other elements in
  group 8A.

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• ELECTRON CONFIGURATION FOR CATIONS
• In 1913, an American chemist by the name of
  Gilbert Lewis, determined that when bonds form
  they generally end up achieving the electron
  configuration of the noble gases.
• This tendency is referred to as the octet rule.
• If you were to script out the electron
  configuration for each noble gas(except helium)
  you would discover that they all contain 8
  electrons in their highest energy level.
• Although there are exceptions, the octet rule
  applies to most atoms in compounds.

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• When an atom loses an electron it becomes a
  cation.
• Sodium, for example, loses one electron, and when
  it does its electron configuration resembles that
  of neon.
• For transition metals, the number of lost
  electrons can vary.
• Because atoms with more than 4 charges are not
  likely, exceptions to the octet rule exist.
• For example, silver would have to lose or gain
  too many electrons to obey the octet rule and as
  a result does not achieve an electron
  configuration like a noble gas.

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• ELECTRON CONFUGURATION OF ANIONS
• The gain of an electron produces an anion.
• Take note that the nonmetals in Group 7A all take
  on noble gas configurations that appear just next
  to them on the periodic table in Group 8A.
• Halide ions is the term given to the halogens in
  Group 7A when they gain an electron.
• At your seat answer questions 1 through 6 on page
  418

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• Answers to page 418
• 1. How can the periodic table be used to infer
  the number of valence electrons in an atom?
• The group number equals the number of valence
  electrons for representative elements.
• 2. Why do metals tend to form cations, and
  nonmetals tend to form anions?
• It is easier for a metal to lose an electron and
  a nonmetal to gain an electron to achieve the
  electron configuration of noble gases.

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• 3. How many valence electrons does each atom
  have?
• a. potassium 1
• b. carbon 4
• c. magnesium 2
• d. oxygen 6
• 4. Write the electron dot structure for each of
  the above.
• Refer to the blackboard.

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• 5. Write the electron configuration for the 1
  ion of copper and the 2 ion of cadmium?
• Cu 1s22s22p63s23p63d10
• Cd 2 1s22s22p63s23p63d104s24p64d10
• Note Because they end up with full highest
  energy electron configurations but not identical
  to a noble gas(S2P6) they are said to have
  pseudo-noble gas configurations.
• 6. How many electrons will each element gain or
  lose in forming an ion?
• a. calcium lose 2
• b. fluorine gain 1
• c. aluminum lose 3
• d. oxygen gain 2

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• 15.2 FORMATION OF IONIC COMPOUNDS
• Electrostatic forces attract oppositely charged
  ions toward each other.
• Ionic bonds are the type of bonds that hold these
  oppositely charged atoms into a formula unit.
• When sodium and chlorine react to form the
  compound sodium chloride, sodium transfers the
  electron it lost over to chlorine.
• As a result the compound exists uncharged.

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• PROPERTIES OF IONIC COMPOUNDS
• At room temperature most ionic compounds exist as
  crystalline solids.
• The resulting 3 dimensional patterns are strong
  and stable, and they generally have high melting
  points.
• The coordination number is the number of
  oppositely charged ions that surround an ion.
• If you look at figure 15.9 on page 423 in your
  text book, you can see how the different ionic
  salts and the way the ions arrange themselves,
  form different crystalline patterns. The size of
  the ion and how their charge relate to each other
  govern the structure of the resulting crystal.

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• Internal structures are examined by a technique
  called x-ray diffraction crystallography.
• X-rays that pass through a crystal are recorded
  on film.
• The pattern on the exposed film shows how the
  ions in the crystal deflect the X-rays and help
  define the structure of the crystal.

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• Ionic compounds can conduct electricity when
  heated.
• The orderly crystal structure of a crystal is
  broken down once the temperature reaches its
  melting point.
• If a voltage is applied to a molten(melted) mass,
  carrying a positive electrode at one end and a
  negative electrode at the other end, the charged
  ions will migrate to opposite electrodes and
  produce a flow of electricity between the
  electrodes
• Ionic compounds can conduct electricity in their
  dissolved state as well, since ions can move
  about freely when they are dissolved.

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• AT YOUR SEAT WORK ON QUESTIONS 9 THROUGH 14 ON
  PAGE 425

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• Answers for page 425
• 9. What are the characteristics of ionic bonds?
• Characterized by attraction between oppositely
  charged ions and electron transfer.
• 10. Explain why ionic compounds can conduct
  electricity when melted and when in aqueous
  solutions.
• They conduct electricity because of the free
  movement of ions.

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• 11. Write the correct chemical formula for the
  compounds formed from each pair of ions.
• a. K, S2-
• K2S
• b. Ca2 , O2-
• CaO
• c. Na, SO42-
• Na2SO4
• d. Al3, PO43-
• AlPO4

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• 12. Write the formula for each compound.
• a. potassium nitrate
• KNO3
• b. barium chloride
• BaCl2
• c. magnesium sulfate
• MgSO4
• d. lithium oxide
• Li2O
• e. ammonium carbonate
• (NH4)2CO3
• f. calcium phosphate
• Ca3(PO4)2

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• 13. Which pairs of elements are likely to form
  ionic compounds?
• Lithium and chlorine
• Iodine and sodium
• Chlorine and bromine are two nonmetals
• Helium is an inert gas
• 14. What determines the crystal structure of an
  ionic compound?
• The charge and relative sizes of the ions.

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• METALLIC BONDS AND METALLIC PROPERTIES
• Metals are made up of closely packed cations
  rather than neutral atoms.
• The cations are surrounded by mobile valence
  electrons that move about freely within the
  metal.
• A metallic bond is a result of an attraction of
  these free- floating valence electrons and the
  positively charged metallic ions.
• Metallic bonding gives a metal its physical
  properties.
• For example
• 1. good conductors of electricity

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• 2. ductile- can be made into wires
• 3. malleable- can be hammered and forced into
  shapes.
• The sea of electrons is what allows metals to
  have these properties because the electrons
  insulate the metal cations from one another.
• Conversely, a crystal with its lattice structure
  that keeps the like charges away from each other,
  if they get subjected to force and the opposite
  forces are drawn closer together, the repulsion
  of the like charges, cause the explosive
  shattering that is witnessed.

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• CRYSTALLINE STRUCTURE OF METALS 15.3
• Metal atoms are arranged in very compact and
  orderly patterns.
• For spheres of metallic atoms, they can take on 3
  different possible types of arrangements
• 1. body-centered cubic
• 2. face-centered cubic.
• 3. hexagonal close-packed.
• Figure 15.16 on page 428 illustrates the three
  types.
• Of the three, the hexagonal is the most closely
  packed arrangement of atoms.

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• ALLOYS
• Very few metallic substances that you come across
  in your daily life are composed of pure metal.
• Most are mixtures.
• An alloy is a mixture that is composed of two or
  more elements where at least one of the elements
  is a metal.
• Alloys are important because they take on
  properties superior to the individual components
  that make up the alloy.
• Steels, for example is made up of a mixture of
  carbon and iron as well as any variety of boron,
  chromium, manganese, molydbdenunm, nickel,
  tungston, and vanadium.

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• Table 15.3 on page 429 list a few common alloys.
• Properties of steel include resistance to
  corrosion,
• Ductility, malleabililty, hardness, and
  toughness.
• If atoms of a component are about the same size,
• They replace each other in the crystal.
• This type of alloy is called a substitutional
  alloy.
• If the atomic size is quite different, the
  smaller atom can fit into the spaces between the
  larger atoms.
• This is called an interstitial alloy.
• For example, in the various types of steel,
  carbon atoms fill in between the iron atoms.

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• WORK ON QUESTIONS 15 THROUGH 19 ON PAGE 429

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• Answers for page 429
• 15. Use metallic bonding theory to explain the
  physical properties of metals.
• The sea of free-flowing electrons conduct
  electricity and heat they also shield cations
  from repulsive during physical stress.
• 16. Describe the arrangement of atoms in metals
• Solid containing tightly packed atoms in a sea
  of electrons.
• 17.Define metallic bond.
• A metallic bond is a stationary positive metal
  ion attracted to mobile electrons

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• 18 . What is meant by ductile and malleable?
• Ductile can be drawn into wires malleable can
  be hammered into different shapes.
• 19. Why is it possible to bend metal but not
  ionic crystals?
• In an ionic compound, ions of like charge do not
  have mobile electrons as insulation. When acted
  upon by a force, the like charges become in
  contact with each other and repulsion causes the
  shattering of the crystal.
• END OF CHAPTER NOTES

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• QUIZ TOMORROW
• HOMEWORK PACKETS FOR CHAPTER 15 DUE WEDNESDAY
• TEST ON TUESDAY
• APRIL 14, 2009