Types of Bonds Between Atoms

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Types of Bonds Between Atoms

• Ionic Bonds
• Electrons are on one atom more than other
• One end of bond is negative
• Bond is polar

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• Covalent Bond
• Electrons are shared between atoms
• Both ends are the same
• Bond is non-polar

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Polar Covalent Bond

• Electrons are partially shared
• One end is slightly more negative
• Bond is still covalent

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Electronegativity

• The energy holding electrons to an atom in a
  chemical bond
• The higher the difference in electronegativity,
  the more polar the bond
• Two identical atoms will have a difference of
  zero and a 100 covalent bond

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Diatomic Elements

• Hydrogen
• Nitrogen
• Oxygen
• Fluorine
• Chlorine
• Bromine
• Iodine

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Polyatomic Elements

• Phosphorus, P4
• Sulfur, S8

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Difference in Electronegativity

• Non-polar ?E 0.0 0.4
• Diatomic elements
• C-H (2.5 2.1 0.4)

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• Polar Covalent ?E 0.5 1.6
• C O (2.5 3.5) ?E 1.0

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• Ionic ?E gt 1.9
• Na Cl (0.9 3.0) ?E 2.1

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Ionic Compound Formation

• Elements will form ions using the octet rule.
• Find the Group Number at the top of the column.
• If there are less than 4 electrons in outer
  shell, the atom will lose electrons and the ion
  will be positive.
• If there are more than 4 electrons in outer
  shell, the atom will gain electrons and the ion
  will be negative.

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Group IA Alkali Metals

• One valence electron
• Always 1 ions
• Na gt Na 1e-

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Group IIA Alkaline Earth Metals

• Two valence electrons
• Always 2 ions
• Mg gt Mg 2 2 e-

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Group IIIA

• Three valence electrons
• Always 3 ions
• Al gt Al 3 3e-

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Group VA

• Five valence electrons
• Gains 3 electrons to form 3 ions
• N 3 e- gt N3-

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Group VIA

• Six valence electrons
• Gains two electrons to form 2 ions
• O 2 e- gt O2-

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Group VIIA - Halogens

• Seven valence electrons
• Gains one electron to form 1 ions
• Cl 1e- gt Cl-

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Group VIIIA Noble Gases

• Eight valence electrons
• Cannot gain or lose electrons
• Chemically inert

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Stable Noble Gas Configuation

• Ions form with the same number of electrons as
  the noble gases.
• Na (11 electrons) gt Na (10 electrons)
• Mg (12 electrons) gt Mg 2 (10 electrons)
• O (8 electrons) gt O2- (10 electrons)
• F (9 electrons) gt F- (10 electrons)

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Isoelectronic Series

• Ions form with Noble gas configurations
• The ions with 10 electrons are
• N3- O2- F- Ne Na Mg2 Al3
• What are the members of the argon isoelectronic
  series?

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Polyatomic Ions

• NH4 ammonium ion
• H3O hydronium ion
• OH- hydroxide ion
• NO3- nitrate ion
• CH3COO- acetate ion (C2H3O2-)
• HCO3- bicarbonate ion

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• CO32- carbonate ion
• SO42- sulfate ion
• HPO42- (mono)hydrogen phosphate ion
• H2PO4- dihydrogen phosphate ion
• PO43- phosphate ion

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Formulas for Ionic Compounds

• Formula must be electrically neutral
• The number of positive charges must equal the
  number of negative charges
• The subscripts tell how many of each atom are in
  the formula

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• Sodium and Chlorine
• NaCl
• Calcium and Chlorine
• CaCl2
• Aluminum and Chlorine
• AlCl3

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• Sodium and Sulfate (SO42-)
• Na2SO4
• Magnesium and Hydroxide
• Mg(OH)2
• Barium and Nitrate
• Ba(NO3)2

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• Potassium and Phosphate
• K3PO4
• Magnesium and Phosphate
• Mg3(PO4)2
• Sodium and Bicarbonate
• NaHCO3

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• Ammonium and Acetate
• NH4CH3COO
• Sodium and Carbonate
• Na2CO3
• Sodium and Oxygen
• Na2O

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Nomenclature

• Find position of first element on Periodic Table
• Metal with constant charge (Group IA, IIA, Al,
  Ag, Zn, Cd, and NH4
• Transition Metal with variable charge
• Non-Metal

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Metals with Constant Charge

• Cation is named as element. Anion is named by
  changing ending to ide
• NaCl sodium chloride
• MgCl2 magnesium chloride
• Na2O sodium oxide
• Mg3N2 magnesium nitride
• (NH4)2CO3 ammonium carbonate

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Transition Metals

• Roman numerals are used to designate charge on
  cation
• FeCl2 iron(II) chloride (ferrous)
• FeCl3 iron(III) chloride (ferric)
• Cu2CO3 copper(I) carbonate (cuprous)
• CuO copper(II) oxide (cupric)

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Non-Metals

• Prefixes are used to tell how many of each atom
  are present in formula
• 1-mono 6-hexa
• 2-di 7-hepta
• 3-tri 8-octa
• 4-tetra 9-nona
• 5-penta 10-deca

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• CO carbon monoxide
• CO2 carbon dioxide
• CCl4 carbon tetrachloride
• XeF6 xenon hexafluoride

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Covalent and Polar Covalent

• Electrons are shared between two atoms
• Bonds are very stable

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Types of Covalent Bonds

• Single Bond
• Share one pair of electrons
• Double Bond
• Share two pairs of electrons
• Triple Bond
• Share three pairs of electrons

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Typical Bonds

• Double
• CC CN CO PO
• Triple
• CC NN CN

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Normal Covalency

• Usual number of bonds
• Carbon 4
• Nitrogen 3 Phosphorus 3 or 5
• Oxygen 2 Sulfur 2 or 4
• Halogens 1
• Hydrogen 1

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VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory
• Shape of molecule using electron pairs
• 2 clouds linear molecule
• 3 clouds triangular molecule
• 4 clouds tetrahedral molecule

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Lewis Structures

• We indicate the number of valence electrons using
  a Lewis Structure
• There are four sides around the symbol
• Each side can hold two electrons
• There are a total of eight around the symbol

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Lewis Structures

• CH4

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• NH3

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• H2O

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• SO3

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• SO2

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• CO2

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• HCHO

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Molecular Polarity

• Symmetrical molecules are non-polar
• CH4
• SO3
• CO2

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• Asymmetrical molecules are polar
• NH3
• H2O
• SO2

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Resonance Hybrids

• When a molecule has two or more equivalent bonds,
  the electrons resonate between the positions.
• SO3
• SO2