Covalent Bonding and Nomenclature Notes

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Covalent Bonding and Nomenclature Notes

• I. Writing Formulas for Binary Molecular
  Compounds-those containing 2 nonmetals. Prefix
  naming system - know theses prefixes
• mono one di two tri three
• tetra four penta - five hexa six
• hepta seven octa - eight nona nine
• deca ten

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• Simply write what it says.
• Ex phosphorus pentachloride PCl5
• dihydrogen monoxide H2O
• Practice
• nitrogen tetrasulfide ______________
• carbon dioxide ________________
• oxygen monofluoride _____________
• sulfur hexachloride __________________
• trioxygen decanitride ______________
• tetrafluorine monophosphide ___________
• hexafluorine nonasulfide ___________
• heptabromine octanitride ____________

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• II. Writing Names for Binary Molecular Compounds
• 1. The less electronegative element is given
  first. It is given a prefix only if it
  contributes more than one atom to a molecule of
  the compound. (All this means is that you will
  never start with mono-)
• 2. The second element is named by combining a
  prefix indicating the number of atoms contributed
  by the element to the root of the name of the
  second element and then adding ide to the end.
• The o or a at the end of a prefix is usually
  dropped when the word following the prefix begins
  with another vowel. (monoxide or pentoxide)
• Common Roots
• H hydr C carb N nitr O ox
• F flor Si silic P phosph S sul
• Cl chlor Br brom I iod

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• Practice
• CCl4 _________________________
• NF3 _______________________
• PBr5_________________________
• SF6_____________________________
• SO3 _________________________
• PCl5 _______________________
• N2O_________________________
• PF6_____________________________

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Covalent Bonds
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Remember what happens when an ionic bond forms?

• One or more electrons from 1 atom are removed and
  attached to another atom, resulting in a cation
  and an ion which attract each other

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Write this down!!

• Ionic Compounds
• never exist as individual
• formula units, are solids
• Molecular Compounds
• -can exist as an individual
• molecule, are usually liquids or gases
• http//web.jjay.cuny.edu/acarpi/NSC/6-react.htm

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III. Formation of Covalent Bonds and Molecular
Compounds

• A. Covalent Bonds a bond in which electrons
  are shared. Which compounds have covalent
  bonding?
• Molecular (or covalent) compounds - these are two
  NON-METALS. These compounds always have covalent
  bonding
• Polyatomic ions (PO-3, NO-1, CN-1). These ions
  are held together with covalent bonds.

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• B. Type of Covalent Bonds
• l. Nonpolar covalent bond-a covalent bond in
  which the bonding electrons are shared equally by
  the bonded atoms, resulting in an evenly balanced
  charge. If the difference in electronegativity
  between two bonded atoms is less than 0.3 a
  nonpolar bond will exist.

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Nonpolar Covalent Bond (equal sharing)
Oxygen Atom
Oxygen Atom

Oxygen Molecule (O2)
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• 2. Polar covalent bond-a bond is which the bonded
  atoms do NOT share the bonding electrons equally.
  A polar covalent bond is a bond in which the
  atoms have an unevenly balanced charge. If the
  difference in electronegativity between two
  bonded atoms is from 0.3 to 1.7, a polar bond
  will exist If the difference in EN is less than
  0.3 then the bond is nonpolar covalent. The atom
  with the greater electronegativity will pull the
  electrons toward it, giving that atom a slightly
  negative charge. A partial negative charge is
  shown by ?- and the less electronegative atom
  will have a partial positive charge, designated
  ?.

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Polar Covalent Bonds Unevenly matched, but
willing to share.
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The H-O bonds in water are polar covalent because
oxygen is more electronegative than hydrogen, and
therefore electrons are pulled closer to oxygen.
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• Practice Find the differences in
  electronegativity (EN page 151 or on calculator)
  in the following pairs of atoms. Designate which,
  if any, atom is partially negative and partially
  positive.
• a. H and Cl
• b. F and Br
• c. S and I
• d. O and H

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Another Example cesium-fluorine bonding

• Cs EN 0.7 F EN 4
• 4 - 0.7 3.3
• Ionic Bond

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D. The Octet Rule and Dot Structures -chemical
compounds tend to form so that each atom, by
gaining, losing, or sharing electrons, has an
octet of valence electrons. Electron dot
structures (also known as Lewis dot diagrams)
show valence electrons as dots around the
elements symbol. Dot structures for molecules
show atoms sharing dots (covalent bonds).
Ar
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• Covalent bonds are single, double, or triple
• single bond-two atoms share one pair of electrons
  (1 sigma bond)
• double bond-two atoms share two pair of electrons
  (1 sigma and 1 pi bond)
• triple bond-two atoms share three pair of
  electrons (1 sigma and 2 pi bonds)

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Rules for correctly illustrating the dot
structure of a molecule

• 1. Add up the TOTAL number of valence electrons
  in the substance
• be sure to subtract 1 electron if it is a
  positively-charged ion (NH41)
• be sure to add electrons for each negative
  charge on an ion (SO4-2)
• 2. Decide what is the central atom. The central
  atom is the one that is least represented. (or
  the least electronegative)
• 3. Hook the particles together using a short
  straight line (or 2 dots) to indicate a covalent
  bond between atoms. Each of these "bonds"
  represents 2 shared electrons.
• 4. Subtract the number of electrons used in
  "hooking" the atoms together from the total
  valence electrons.
• 5. Use the "leftover" electrons, if any, to fill
  the octets of the peripheral atoms.
• 6. Place anymore "leftover" electrons on the
  central atom (in pairs).
• http//chemsite.lsrhs.net/d_bonding/flashLewis.htm
  l

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Practice Draw the electron dot structures for
the following

• Practice Draw the electron dot structures for
  the following
• 1. carbon tetrachloride (CCl4)
• 2. F2O
• 3. NF41
• 4. PCl3
• 5. CO2
• 6. N2

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E. Exceptions to the Octet Rule

• Some atoms have less then an octet.
• Example Hydrogen only needs 2 electrons
  surrounding it and boron only needs 6.
• H2
• BF3
• b. Some atoms have more than an octet (One
  reason is because of bonding d orbitals as well
  as s and p orbitals.) Example Sulfur can have
  up to 12 electrons surrounding it.
• SF6

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• Resonance a concept in which two or more Lewis
  structures for the same arrangement of atoms
  (resonance structures) are used to describe the
  bonding in a molecule or ion. To show resonance,
  a double-headed arrow is placed between a
  molecules resonance structures.
• Example Ozone
• Coordinate Covalent Bond is formed when one atom
  contributes BOTH bonding electrons in a covalent
  bond.
• Examples
• carbon monoxide
• SO42-
• HCN

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Covalent Bonding and Nomenclature Notes Part 2

• I. Metallic Bonds-a third type of bond. This is
  what holds pure metal atoms together.
• What happens to form a metallic bond?
• 1. each metal donates its valence
  electron(s) to form an electron cloud
• 2. this leaves positive particles which
  are "cemented" together with the negative
  electron cloud, often called a sea of electrons.

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Metallic Bonds Mellow dogs with plenty of bones
to go around.
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A Sea of Electrons
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Metals Form Alloys
Metals do not combine with metals. They form
Alloys which is a solution of a metal in a
metal. Examples are steel, brass, bronze and
pewter.
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II. Polarity - Polar and nonpolar molecules - if
a molecule contains a polar bond, is the molecule
itself polar also? It depends!!
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• 1. Polar molecules - a polar molecule is
  positive at one point and negative at another
  point. For example, HBr contains a polar bond.
  As a result the hydrogen side of the molecule is
  partially positive and the bromine side of the
  molecule is partially negative. It "acts like a
  magnet".
• Water is a molecule with two polar bonds.
  A molecule of water is also polar because there
  is an area of positive charge on the hydrogen
  atoms and an area of partially negative charge on
  the oxygen. Its bent shape allows it to act
  somewhat like a magnet. The presence of 2
  unshared pair of electrons is a main factor for
  it being a polar molecule.

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• 2. Nonpolar molecules - Carbon tetrachloride has
  four C-Cl bonds. Each bond is a polar covalent
  bond.
• The molecule itself is nonpolar because of
  its
• 1.) shape. It is perfectly symmetrical, and
• 2.) the partially-positive carbon in the
  center which is covered by the 4 partially
  negative chlorine atoms. It cannot act like a
  magnet.

Cl
Cl
Cl
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• 3. Helpful hints and practice
• A. Hints to help you decide if a molecule
  is POLAR
• 1. Does it have at least one polar
  bond? If so, it's probably polar.
• 2. Does it have any unshared pairs of
  electrons around the central atom? If so, it is
  probably polar.
• 3. Can the molecule act like a magnet?
  If so, it is probably polar.
• B. Practice Which of the following
  molecules are polar and which ones are nonpolar
  molecules?
• If the molecule is polar, tell why it
  is polar.
• 1.) SO2
• 2.) H2S
• 3.) CO2
• 4.) BF3
• 5.) CH4
• 6.) ClO2-1
• 7.) CH3Cl
• 8.) PO4-
• 9.) MgCl2

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• III. Hydrates Some compounds trap water inside
  their crystal structure and are known as
  hydrates. You will not be able to predict which
  compounds will form hydrates. CuSO4 5H20 is an
  example of a hydrate. This says that one formula
  unit of cupric sulfate will trap 5 molecules of
  water inside its crystal.
• Hydrates are named by naming the ionic compound
  by the regular rules and then adding (as a second
  word) a prefix indicating the number of water
  molecules. You will use the word hydrate to
  indicate water. The above compound would be
  called cupric sulfate pentahydrate.
• To find the formula mass of a hydrate, simply
  find the mass of the ionic compound by itself and
  then ADD the mass of water molecule(s) to that
  mass.
• Practice What is the formula mass of barium
  chloride dihydrate?
• What is the formula mass of aluminum sulfate
  octahydrate?