Chapter 7 Chemical Formulas and Bonding

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Chapter 7Chemical Formulas and Bonding

• How it all sticks together.

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Some Questions to Consider

• Why are so few elements (such as Au, S, N, O, Ag)
  found in Nature in their free atomic state?
• Why do atoms of different elements react to form
  compounds?
• What is happening in this process?
• How can we explain the tremendous number of
  compounds that are known today?
• Many of the answers will be found in Chapter 6
  (Chemical Formulas and Bonding).

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Chapter 7 Objectives

• Describe the characteristics of an ionic bond.
• State and use the Octet Rule.
• Learn how to use Lewis Dot diagrams.
• Learn the types of ions.
• Describe the characteristics of a covalent bond.
• Describe the difference between polar and
  nonpolar covalent bonds.
• Write names for ionic compounds, molecular
  compounds and acids.

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7-1 Ionic Bonding

• Recall how ions form
• Metals lose electrons to become positive ions
  (cations). (Which ones do they lose?)
• M ? M1 e1-
• Nonmetals gain electrons to become negative ions
  (anions). (Where do they go?)
• X e1- ? X1-
• Positive ions are attracted to negative ions.
• Opposites attract.
• Ionic Compound one that is composed entirely of
  ions.
• Total charge balances to zero.
• That is, total () charges total (-) charges
  Zero

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Ionic Bonding Example

• Sodium (Na) is a poisonous, very reactive metal.
• Chlorine (Cl2) is a poisonous, very reactive
  nonmetal.
• They combine explosively to form salt, NaCl.
• NaCl is composed of Na1 and Cl1- ions, and it is
  harmless.
• Na ? Na1 e1-
• Cl e1- ? Cl-
• Overall Na Cl ? NaCl

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The Octet Rule

• Atoms tend to gain, lose or share electrons in
  order to acquire a full set (8) of valence
  electrons.
• EXAMPLE Look at the sodium and chlorine atoms
  in forming salt. (Fig. 7-5, p. 228)

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The Role of Valence Electrons

• Note that the valence electrons were involved in
  this change, NOT the core electrons.
• Why? (Which orbitals electrons are encountered
  first when two atoms interact?)
• Chemists focus on the valence electrons (outer
  electrons) to understand the chemistry of atoms.
• To aid us, we use shorthand diagrams, called
  Lewis Dot Diagrams, where dots represent the
  valence electrons around an atom.
• Lets do some examples.

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Lewis Dot Diagrams

• Write the element symbol.
• Use dots to show the valence electrons (alone or
  in pairs) around the symbol.
• Sodium would be Na with one dot.
• Chlorine would be Cl with seven dots.
• The reaction of sodium with chlorine would be
  written as
• Na. .Cl ? Na. .Cl ? Na1 .Cl1-

.



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Lewis Dot Diagrams (Practice)
Practice doing this!
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Types of Ions

• Monoatomic Cations
• Na, Mg2, Al3
• Fe2 Iron(II), Fe3 Iron(III)
• Monoatomic Anions
• F-, Cl-, Br-
• Polyatomic Ions
• NH4, OH-, NO3-, SO42-, CO32-, PO43-
• See list of ions you MUST learn!
• Pages 231 232

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Facts About Ionic Compounds

• Binary Ionic Compound - contains ions of only two
  elements. (e.g. NaCl)
• Empirical Formula the formula of a compound
  with the lowest whole-number ratio of the
  elements.
• NaCl (NOT Na2Cl2 or Na3Cl3 or Na100Cl100)
• The net charge of a neutral compound must equal
  zero, which tells us the ion ratio. (Ca2 Cl1-
  needs CaCl2 as the correct formula.)
• Crisscross method helps write ionic formulas.
• Ba2 Br1- becomes BaBr2
• Al3 NO31- becomes Al(NO3)3
• PRACTICE, PRACTICE, PRACTICE!

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Crisscross Method for Ionic Compounds

• Ionic compounds must have a net ionic charge of
  zero (neutral).
• The total and charges must cancel.
• Keep polyatomic ions intact!
• Use crisscross method to write formulas
• The charge superscript becomes the subscript of
  the opposite ion, indicating the number of ions.
• Ba2 Br1- becomes BaBr2 2 with 2(1-) 0
• Al3 NO31- becomes Al(NO3)3 3 with 3(1-) 0
• NH41 and SO42- becomes (NH4)2SO4 2(1) with 2-
  0
• PRACTICE, PRACTICE, PRACTICE!

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Naming Ionic Compounds

• Chemists name compounds on the basis of the atoms
  and bonds present.
• Ionic compounds are named from their elements or
  polyatomic ions.
• Cations () are named first (usually an element
  name).
• If it can have more than one charge, use Roman
  numerals to indicate which ion is actually
  present.
• FeCl3 is iron(III) chloride FeCl2 is iron(II)
  chloride.
• Change the ending of the anion to ide (unless a
  polyatomic ion is present).
• NaCl is sodium chloride.
• Al2O3 is aluminum oxide.
• Ba(NO3)2 is barium nitrate.
• K2SO4 is potassium sulfate.
• What is NiBr2? Sr3(PO4)2? FeI2?

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Hydrates

• Hydrate Ionic compound that absorbs water into
  its solid form.
• Recall the blue copper sulfate lab when we
  studied chemical/physical changes?
• Anhydrous Water-free substance.
• Name these ionic compounds to reflect the water
  of hydration.
• Name the compound in the normal way.
• Add the word hydrate and a prefix term to show
  the number of water molecules (degree of
  hydration).
• See Fig. 7-24 on page 246.
• Di-, tri- tetra-, penta- etc.
• MgSO4 7 H2O is magnesium sulfate heptahydrate.
• What is the formula for copper(II) sulfate
  pentahydrate?

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Properties of Ionic Compounds

• High melting points (usually).
• NaF (996 C), NaCl (801 C)
• This indicates very strong ionic bonding.
• Very brittle.
• Shatter, or cleave, in fixed paths.
• Example Rock salt.
• Water soluble (usually).
• Water breaks the ionic bonds.
• Aqueous solutions conduct electricity because the
  ions are free to move about in the water.
• Conduct electricity when molten (liquid).
• Ions are freed from the crystal structure
  (lattice).
• Do not conduct electricity when solid.
• Ions are held firmly in place, so they simply
  vibrate.

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7-2 Covalent Bonding

• A covalent bond is formed by a shared pair of
  electrons between two atoms.
• Molecule group of atoms united by a covalent
  bond.
• Molecular Substance a material made up of
  molecules.
• Molecular formula chemical description of a
  molecular compound or molecule.
• Structural Formula a formula that specifies
  which atoms are bonded to each other in a
  molecule.
• Lewis Structures molecular structure based on
  Lewis Dot diagrams.

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Describing Covalent Bonds

• Draw Lewis diagrams, including unshared pairs of
  electrons.
• Use a dash for each pair of electrons in a
  bond.
• Examples ClCl becomes Cl-Cl.
• Single covalent bonds
• CC or simply C-C (Note the dash.)
• Double covalent bonds
• CC or simply CC (Note the double dash.)
• Triple covalent bonds
• CC or simply C?C (Note the triple dash.)

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Properties of Covalent Compounds

• Low melting points (usually).
• Methane, (CH4) is a gas at room temperature
  oils are liquids at room temperature wax melts
  at 100C.
• This indicates very weak molecular association.
• Soft.
• Wax feels may be deformed even as a solid.
• Insoluble in water (usually).
• Water cannot break the covalent bonds.
• Aqueous solutions do not conduct electricity (no
  ions are free to move about in the water).
• Do not conduct electricity when molten (liquid).
• Again, there are no ions to move about.
• Do not conduct electricity when solid.
• No ions!

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Exceptions to the Octet Rule

• Atoms with less than an octet.
• Boron compounds.
• Atoms with more than an octet.
• Atoms with d-electrons, such as sulfur.
• Molecules with an odd number of electrons.
• So called Radicals like nitroxyl, NO.

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Properties of Covalent Bonds

• Remember electronegativity? (What is it?)
• The ability of an atom to attract electrons in a
  chemical bond.
• Fr has the lowest (0.7) and F has the highest
  (4.0) on the Pauling scale.
• Electronegativity differences (delta EN or
  ?EN) dictate which atom in a bond more strongly
  attracts the electrons.
• See Fig 7-20, page 242, and the following slide.
• Chemists use lower case Greek letter delta (d) to
  mean a partial or small difference.

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Bond Type by Electronegativity(Use the
electronegativity difference, ?EN, to predict the
bond type.)
Note that a large ?EN means that it is an ionic
bond. Electrons have transferred from one atom
to another.
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7-3 Naming Chemical Compounds

• Ionic compounds are named from their elements or
  polyatomic ions.
• Hydrates have water in their solid structure, but
  anhydrous substances do not.
• Molecular compounds are named using prefixes to
  indicate the number of atom in the formula.
• Acids have special names that must be memorized
  (Fig 7-27, pg 249).
• PRACTICE, PRACTICE, PRACTICE!

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Naming Molecular Compounds

• Use the element names and prefixes to indicate
  the number of atoms in the formula.
• Di-, tri-, tetra-, etc.
• CO is carbon monoxide. (Mono is not used for
  the first element generally.)
• CO2 is carbon dioxide.
• N2O is dinitrogen monoxide.
• N2O4 is dinitrogen tetroxide. (not usually
  tetraoxide because it is hard to say!)
• Name N2O5. SO3. BF3. PF5
• Many molecular compounds have common names.
• Dihydrogen monoxide is ______?
• Trihydrogen mononitride is ammonia.

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Naming Common Acids

• Acids are molecular substances that dissolve in
  water to produce hydrogen ions (H).
• Acids have special names that must be memorized
  (Fif. 7-27, page 249), but focus on these and
  their anions
• Hydrofluoric, hydrochloric, hydrobromic,
  hydroiodic,
• Nitric
• Sulfuric
• Carbonic
• Phosphoric
• Acetic

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Did we meet the Chapter 7 Objectives?

• Describe the characteristics of an ionic bond.
• State and use the Octet Rule.
• Learn how to use Lewis Dot diagrams.
• Learn the types of ions.
• Describe the characteristics of a covalent bond.
• Describe the difference between polar and
  nonpolar covalent bonds.
• Write names for ionic compounds, molecular
  compounds and acids.