Chapter 7

1
Chapter 7Chemical Formulas and Bonding

• How it all sticks together.

T. Witherup 11/06
2
Some Questions to Consider.
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???

• Why are so few elements (such as Au, Ag, S, N, O)
  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 millions of compounds that
  are known today?
• Answers to these questions will be found in
  Chapter 7 (Chemical Formulas and Bonding).

3
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
  non-polar covalent bonds.
• Write names for ionic compounds, molecular
  compounds and acids.

4
7-1 Ionic Bonding

• Whats an ION?
• An atom or group of atoms having a charge.
• Do you remember how ions form?
• Metals lose electrons to become positive ions,
  called cations. (Which electrons do they lose?)
• M ? M1 e1-
• Nonmetals gain electrons to become negative ions,
  called anions. (Where do the new electrons go?)
• X e1- ? X1-

5
7-1 Ionic Bonding (contd)

• Positively charged ions are attracted to
  negatively charged ions.
• Why?
• Because opposites attract.
• Ionic Compound A substance that is composed
  entirely of ions.
• An ionic formula is the simplest whole-number
  ratio of the ions, so the total charge balances
  to zero.
• Total () charges total (-) charges Zero

6
Ionic Compound (General Example)

-
and
combine to form an Ionic Compound
Cations
Anions
These ions are held together in a solid by
electrostatic attraction


-
-


-
-


-
-
-

-

-
-




-
-
7
Ionic Bonding (Specific Example)

• Sodium (Na) is a poisonous, very reactive metal.
• Chlorine (Cl2) is a poisonous, very reactive
  nonmetal.
• They combine violently to form ordinary table
  salt, NaCl, which is relatively harmless.
• NaCl is composed of Na1 and Cl1- ions.
• Na ? Na1 e1-
• Cl e1- ? Cl-
• Overall Na Cl ? NaCl (Note the 11
  ratio.)

8
The Octet Rule

• Atoms tend to gain, lose or share electrons in
  order to acquire a full set (8) of valence
  electrons.

Na Ne3s1
Loses a 3s1 electron to form Na1 (Ne
electron core).
1
(Na1)
e1-
e1-
1-
Gains an electron in 3p to form Cl1- (3s23p6)
(Ar electron core).
Cl Ne3s23p5
e1-
e1-
e1-
e1-
(Cl1-)
e1-
e1-
e1-
e1-
e1-
e1-
e1-
e1-
e1-
e1-
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The Role of Valence Electrons

• Note that only 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 Diagram Method

• 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.
• Our previous reaction of sodium with chlorine
  would be written as
• Na. .Cl ? Na. .Cl ? Na1 .Cl1-

.



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Lewis Dot Diagrams (Practice)
Element Electron Configuration Lewis Dot Diagram
Li He2s1
Be He2s2
B He2s22p1
C He2s22p2
N He2s22p3
O He2s22p4
F He2s22p5
Ne He2s22p6
Al Ne3s23p1
P Ne3s23p3
Practice doing this! Remember, show only the
valence electrons.
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Types of Ions

• Monoatomic Cations
• Na1, Mg2, Al3
• Fe2 Iron(II), Fe3 Iron(III)
• Monoatomic Anions
• F1-, Cl1-, Br1-
• Polyatomic Ions
• NH41, OH1-, NO31-, SO42-, CO32-, PO43-
• See list of ions you MUST learn!
• Pages 231 232
• http//www.ausetute.com.au/wriiform.html

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Facts About Ionic Compounds

• Binary Ionic Compound - contains ions of only two
  elements. (e.g. NaCl, CaBr2)
• 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.)

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Rules for Writing Ionic Formulas

• Use the simplest whole number ratio of Cation and
  Anion.
• Since the net charge must be zero, balance the
  number of cations and anions so the total
  positive charge equals the total negative charge.
• Use subscripts after each ion to indicate how
  many are present. (Omit 1 though.)
• Use parentheses around polyatomic ions and
  indicate their number with a subscript outside
  the parenthesis.
• Crisscross method helps write ionic formulas.
• See the next slides.

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Crisscross Method for Writing Ionic Compound
Formulas

• Ionic compounds must have a net ionic charge of
  zero (neutral).
• The total and charges must cancel.
• Always 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

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Crisscross Method Examples
Barium bromide
BaBr2
Ba2
Br1-
becomes
Aluminum nitrate
Al3
NO31-
Al(NO3)3
becomes
Notice that 1 is not written, that the nitrate
ion is kept intact, and that the net charge is
zero. (For example, barium bromide, 1(2) 2(1-)
0)
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Crisscross Method More Examples
Ammonium sulfate
(NH4)2SO4
NH41
SO42-
becomes
Notice how the parentheses are used.
Aluminum oxide
O2-
Al2O3
Al3
becomes
Notice how the net charge is zero. 2(3)
3(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
  their crystals.
• Blue copper sulfate contains several water
  molecules in its crystal. We will do a lab about
  this.
• Anhydrous A water-free substance.
• These ionic compounds are named 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 rather than
  randomly.
• 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.
• Empirical Formula - the formula of a compound
  with the lowest whole-number ratio of the
  elements.
• 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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Covalent Bond Formation
Sharing of electrons, as in two chlorine atoms!
..
..
combines with
.Cl
Cl.
..
..
to form a Cl2 molecule by sharing electrons.
..
..
ClCl
..
..
This is a diatomic molecule, along with
molecules of fluorine, bromine, iodine,
hydrogen, nitrogen, and oxygen.
Professor BrINClHOF will help you remember them!
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Describing Covalent Bonds

• Draw Lewis dot diagrams, including unshared pairs
  of electrons.
• Use a dash for each pair of electrons in a
  bond.
• Examples Chlorine (Cl2) is written as 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 slippery and 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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Properties of Covalent Bonds (contd)

• 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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Polarity

• Refers to the unequal sharing of electrons in
  covalent bonds of compounds.
• When both atoms in a bond are identical, they
  form NONPOLAR bonds (e.g. Cl2 or F2) because
  there is, equal sharing.
• When one atom has higher electronegativity than
  the other, it forms a POLAR bond (e.g. HCl),
  which means the electrons are not shared equally.
• We use delta /- (d or d-) or arrows (?) to
  show polarity of a bond.

H-Cl

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Bond Type by Electronegativity(Use the
electronegativity difference, ?EN, to predict the
bond type.)
?EN Bond Type
2.0 Ionic
0.4 to 2.0 Polar Covalent
0.4 Pure Covalent (Non-polar Covalent)
Note that a large ?EN means that it is an ionic
bond. Electrons have transferred from one atom
to another.
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A Special Type of Bonding

• Metallic Bonding the force of attraction that
  holds metals together.
• Positive metal ions are in a sea of electrons
  (freely floating valence electrons) that are
  shared.
• This accounts for metallic properties, such as
  electrical conductivity, luster, ductility,
  malleability.
• Drifting electrons insulate the metal ions from
  one another, so the ions can easily slide past
  each other when stressed, unlike ionic solids,
  which shatter when stressed.

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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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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 these 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
  (Fig. 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
  non-polar covalent bonds.
• Write names for ionic compounds, molecular
  compounds and acids.