Covalent Compounds

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Covalent Compounds
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Why do atoms bond?

• When a nucleus attracts electrons of another
  atom
• Or oppositely charged ions attract( ionic
  bonds-metals and nonmetals)
• We know the octet rule- atoms tend to gain or
  lose or share e- in order to acquire a full set
  of 8 valence e-

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Properties

• 1)  Covalent compounds generally have much lower
  melting and boiling points than ionic compounds. 
  
• 2)  Covalent compounds are soft and squishy
  (compared to ionic compounds, anyway
• On the other hand, covalent compounds have these
  molecules which can very easily move around each
  other, because there are no bonds between them. 
  As a result, covalent compounds are frequently
  flexible rather than hard.

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More Properties

• 3)  Covalent compounds tend to be more flammable
  than ionic compounds.
• 4)  Covalent compounds don't conduct electricity
  in water.
• 5)  Covalent compounds aren't usually very
  soluble in water.

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More Properties

• There's a saying that, "Like dissolves like". 
  This means that compounds tend to dissolve in
  other compounds that have similar properties
  (particularly polarity).  Since water is a polar
  solvent and most covalent compounds are fairly
  nonpolar, many covalent compounds don't dissolve
  in water.  Of course, this is a generalization
  and not set in stone - there are many covalent
  compounds that dissolve quite well in water.

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Covalent Bond

• Chemical bond that results from sharing valence
  electrons. (Generally occurs when elements are
  close to each other on the periodic table and
  between nonmetallic elements)
• Molecule- a another name for covalent compound,
  is formed when two or more atoms bond covalently.

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more

• Diatomic molecules- are molecules that occur in
  nature not as single atoms because the molecules
  formed are more stable than the individual atoms.
  There are 7 of these Hydrogen (H2), Nitrogen
  (N2), Oxygen (O2), Fluorine (F2), Chlorine (Cl2),
  Bromine (Br2), and Iodine (I2).

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Naming Covalent Compounds

• Rule 1. The element with the lower group number
  is written first in the name the element with
  the higher group number is written second in the
  name.
• Exception when the compound contains oxygen and
  a halogen, the name of the halogen is the first
  word in the name.
• Rule 2. If both elements are in the same group,
  the element with the higher period number is
  written first in the name.

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More Naming Rules

• Rule 3. The second element in the name is named
  as if it were an anion, i.e., by adding the
  suffix -ide to the name of the element.Rule 4.
  Greek prefixes are used to indicate the number of
  atoms of each nonmetal element in the chemical
  formula for the compound.
• Exception if the compound contains one atom of
  the element that is written first in the name,
  the prefix "mono-" is not used.

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Greek Prefixes

• mono-1di-2tri-3tetra-4penta-5hexa-6he
  pta-7octa-8nona-9deca-10

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Covalent Bonds

• When showing the bonding between atoms of
  covalent compounds, either a pair of dots or a
  line is the Lewis structure of a molecule.
  Lewis structures- use electron-dot diagrams to
  show how electrons are arranged in molecules

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more

• Group 7A atoms have 7 valence e- and need 1 e- to
  fulfill the octet rule and become stable. So they
  will form a single covalent bond
• Do example
• Group 6A need 2 e- and will form 2 covalent bonds
• Group 5A need 3e- and will form 3 covalent bonds
• Group 4A will form 4 covalent bonds.
• Lets practice!!!

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Molecular Structures

• Several models can be used to represent a
  molecular structure Molecular formula,
  structural formula, Lewis structure and ball and
  stick.
• Predicting the location of certain atoms
• Hydrogen is always terminal
• Halogens usually terminal
• Elements with more than one atom are usually
  terminal
• Central atom- smallest electronegativity
• Find total number of electrons for bonding9 total
  valence electrons)
• Determine pairs (divide valence electrons by 2)
• Draw Bonds from central atom to other atoms
• Subtract bonded e- for total valence e-
• Place remaining e- around to complete octets
• If there are not enough electrons to give the
  atoms 8, except for hydrogen, make double and
  triple bonds.
• C,N,O and S can form double and triple bonds with
  same element and others.

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Single and Multiple Bonds

• The sigma bond(s)- single covalent bond
  (overlapping of the valence atomic orbitals
  resulting in the e- being in a bonding orbital
  between the two atoms)
• Multiple Covalent Bonds- atoms can attain a noble
  gas configuration by sharing more than one pair
  of electrons between two atoms.

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Atoms can have

• Single - 1 pair of e-
• Double 2 pair of e-
• Triple 3 pair of e-
• Try CO2, O2, N2

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Electronegativity and Lewis Dot Structures

• When these atoms share electrons they are often
  not equally shared. This is due to
  electronegativity.
• Hydrogen and halogens usually bond to only one
  other atom and are usually on the outside of the
  molecule.
• The atom with the smallest electronegativity is
  often the central atom of the compound.
• When a molecule contains more atoms of one
  element than the others, these atoms often
  surround a central atom.

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More

• A pi bond p is formed when parallel orbitals (p
  orbitals )overlap to share electrons
• Single bond- sigma bond
• Double bond- sigma and 1 pi bond
• Triple sigma and 2 pi bonds

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Strength of Covalent Bonds

• Covalent bonds involve attractive and repulsive
  forces
• Nuclei() and electrons(-) attract each other but
  
• Nuclei() repel nuclei() and
• (e-) repel (e-)

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More

• There is a balance of attractive and repulsive
  forces
• When this balance is disrupted the bond could
  break
• Factors controlling Bond Strength
• 1. distance between nuclei bond length

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More

• Bond length
• Determined by size of atoms and number of shared
  electrons
• Size- length decreases when number of bonds
  increases
• Example- a triple bond ( 3 shared pairs of e-) is
  shorter than a double bond (2 shared pairs of
  electrons)
• Strength- shorter the length stronger the bond

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Bond Energy

• Energy change occurs when bonds are formed
• Bond Dissociation Energy- the amount of energy
  required to break a specific covalent bond- this
  is always a value because energy is added to
  break a bond
• The sum of BDE for all bonds in a compound the
  amt. of chemical PE available in one molecule of
  a compound
• BDE the strength of the bond
• as two atoms are bonded closer together the BDE
  increases
• Formation- energy released-exothermic
• Broken- energy added-endothermic

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Examples of BDE and Bond Length

• BDE
  Bond
• 
  Length
• F2 F---F 159 kJ/mole short
• O2 OO 498 kJ/mole shorter
• N2 NtripleN 945 kJ/mole shortest

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Exceptions

• Some molecules and ions do not obey the octet
  rule.
• Reasons
• 1. some group has an odd number of valence e- and
  cant form an octet
• Ex. NO2
• 2. some groups form with few than 8 e- present.
  This is very reactive and will share an entire
  pair of e- which is a coordinate covalent bond.
• exBH3
• 3. central atom contains more than 8 e-
  (expanded octet)- usually occurs with elements in
  period 3 or higher.
• Ex. PCl5 or SF6

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Resonance Structures

• A condition when more than one valid Lewis
  structure can be written for a molecule or ion.
  This occurs when a molecule or ion has a
  combination of single and double bonds.
• Lets try to draw the resonance structures of
  NO3-
• Now you try for homework
• SO3, SO2, O3, NO2-

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Molecular Shapes

• Valence Shell Electron Pair Repulsion
• VSEPR model
• Many reactions depend on the ability of two
  compounds contacting each other. The shape of
  the molecules determines whether they can get
  close enough to react.
• The repulsion of electron pairs in a molecule
  result in atoms existing at fixed angles to each
  other- Bond angle

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Electronegativity and Polarity

• We know electronegativity( the attraction an atom
  has for electrons) increases as we go up and to
  the right of the periodic table. So Fluorine is
  the most electronegative element.
• If we have any of the diatomic molecules, we have
  identical atoms bonded together, their
  electronegativities are the same so they are
  nonpolar molecules. There is equal sharing of
  electrons

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More

• Chemical bonds between elements are never
  completely ionic or covalent.
• When the difference in electronegativity between
  atoms increases it is more ionic
• When the difference in electronegativity between
  atoms is less it is more covalent
• Unequal sharing of electrons results in a Polar
  Covalent Bond

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More
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More

• Molecules are either polar or non polar
• With polarity comes a partial () and partial (-)
  
• Show how partial charges spread out on
• H2O
• CCl4

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3 resonance structures
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Dichromate Ion

• Lewis Structure
• Cr2O72- 56 valence electrons