Ionic Compounds

1
Ionic Compounds

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• Mr. Chan
• Northwestern University

2
Day 8/15

• 30 Discuss HW,Labs,Quiz
• 30 QUIZ
• 30 Ionic Compounds
• 30 Go Fish
• 60 Lunch
• 45 Covalent Compounds
• 15 Exit Slips
• 105 Labs

3
Determining Number of Valence Electrons and
Drawing Electron Dot Structures

• Valence Electrons
• Electrons in outermost energy level, (highest
  occupied energy level of element)
• Count number of electrons in highest energy level
  (n1,2,3,4)
• Use placement in group to determine valence
  electrons
• Electron Dot Structures
• Represent element, core electrons (symbol), and
  valence electrons (dots)

4
Practice

• How many valence electrons?
• Potassium
• Carbon
• Magnesium
• Oxygen
• Now, draw electron dot structures

5
Formation of Cations and Anions

• Remember those charges? How do we get those?
• Octet Rule (from observations/experiments)
• Atoms tend to achieve the electron configuration
  of a noble gas
• Duet rule hydrogen and helium
• Either gaining or losing electrons not protons
  or neutrons!

6
Cations and Anions

• Cations
• Lose valence electrons to achieve octet
• SHOW electron configuration of noble gas
• Ions for transition metals
• Do not acquire noble gas configurations
• But, can achieve ones that are similar or
  reasonably full
• Anions
• Gain electrons to achieve octet

7
Examples

• Number of electrons gain or lose in forming an
  ion
• Na, F, Al, Ba, O
• Now, what is the formula of the ion?

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Characteristics of Ionic Bonds

• Ionic bonds
• Forces of attraction that bind oppositely charge
  ions
• Ionic compounds also called salts
• Show with electron dot structures (Na, Cl)
• Both ions achieve stable octets
• Note connections to chemical formulas
• Formula unit smallest sample of ionic compound
  that has the composition of the compound.

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Examples

• Using electron dot formulas to predict formulas
  of ionic compounds formed from
• Potassium and iodine
• Calcium and sulfur
• Sodium and phosphorus
• Aluminum and oxygen

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Properties of Ionic Compounds

• Most are solids at room temperatures
• Atoms or ions are arranged in repeating 3-D
  patterns
• Coordination number
• Number of ions of opposite charge that surround
  each ion in a crystal
• Very stable compounds
• high melting temperatures (SHOW)
• Conduct electric current (SHOW)
• As liquids, or in aqueous solutions

11
Metallic Bonding to Explain Physical Properties
of Metals

• Metals
• Closely packed cations and mobile valence
  electrons
• Metallic Bonds
• Attraction of free valence electrons for
  positively charged metal ions
• Properties explained
• Conductors of electrical current (SHOW)
• Ductility (SHOW)
• Malleability (SHOW)

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Alloys

• Alloys
• Mixtures of 2 or more elements, 1 metal
• Sterling silver (92.5 Ag, 7.5 Cu
• Bronze (7 Cu 1 Sn)
• Steel (Iron, Carbon)
• Stainless Steel (80.6 Fe, 18Cr, 0.4C, 1Ni)
• Amalgams (Dental Filling Hg, Ag, Zn liquid
  that hardens quickly)

13
Go Fish!

• Card Game
• RULES

14
Exit Slips

• Midterm Self Assessment
• Give yourself 2 grades
• 1) Class grade (estimate using HW, LABS,
  QUIZZES, and TESTS)
• 2) Effort grade
• Write a few sentences justifying your grades and
  any discrepancies..
• Additional feedback regarding class/changes/sugges
  tions

15
Covalent Compounds

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• Mr. Chan
• Northwestern University

16
Covalent Bonding

• Single covalent bond
• Formed when a pair of electrons is shared between
  two atoms
• Structural formulas
• Show arrangement of atoms in molecules and
  polyatomic ions
• Molecular formulas vs. Formula units
• Lewis Dot Structures
• Sharing of electrons occurs if atoms can acquire
  the electron configurations of noble gases (OCTET
  RULE!)
• Shared and unshared pairs of electrons

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Double and Triple Bonds

• 2 or 3 shared pairs of electrons
• Still obeys octet rule
• Examples
• CO2
• N2, HCN

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Coordinate Covalent Bonding/Ions

• Recall CO example
• Coordinate covalent bond
• Formed when one atom contributes both bonding
  electrons in a covalent bond
• Dot Structures for Polyatomic Ions
• NH4 (loss of electron)
• SO3 2- (gain of electrons)

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Resonance

• Draw electron dot structure for ozone, O3
• How do you know which structure is correct?
• Experiments show two bonds in ozone are same
  length
• Average of the two electron dot structure
  (hybrid, not resonating)
• Resonance
• Two or more valid electron dot formulas can be
  written for a molecule

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Exceptions to the Octet Rule

• Odd Numbers
• NO2
• Paramagnetic
• Relatively strong attraction to an external
  magnetic field
• Diamagnetic
• All of the electrons are paired
• Weakly repelled by an external magnetic field
• Incomplete Octet
• BF3
• Exceeding the Octet
• PCl5
• SF6

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Relationship between Covalent bonding and orbital
diagrams

• H2O
• Show O and H orbital diagrams
• S from H and P from O overlap
• CO2
• S and P together leads into hybrid orbitals
  discussion
• Practice Drawing
• H2S
• PH3
• ClF
• HCN

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Predicting the Shape of Molecules

• Using 3-D model as guide (molecular modeling
  lab), we could see how atoms arrange themselves
  outside of 2-dimensional drawings on paper
• CH4
• Tetrahedral, 109.5 bond angles

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VSEPR Theory

• Valence Shell Electron Pair Repulsion
• Molecular shape adjusts so the valence electron
  pairs are as far apart as possible.
• Tetrahedral, as opposed to a square for CH4
• Confirmed by experiment
• NH3
• Pyramidal, 107 degrees (lone pair repels bonding
  pairs)
• H2O
• Bent (planar), 105 degrees
• CO2
• Linear, 180 degrees no unshared pairs of
  electrons
• List of possible shapes
• Examples
• CO2, SiCL4, SO3, SCL2, CO

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Classifying as Polar, Nonpolar, or Ionic

• Not all covalent bonds are the same
• Nonpolar covalent bond
• Bonding electrons shared equally
• Polar covalent bond
• Bonding electrons shared unequally
• Two different elements may or may not be polar
• Recall electronegativity
• atoms strength of attracting electrons when
  shared

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Using electronegativity difference

• 0-0.4 nonpolar covalent
• 0.4-1.0 moderately polar covalent
• 1.0-2.0 very polar covalent
• 2.0 Ionic
• Note continuum ranges are approximate, not
  necessarily exact
• Examples
• What type of bond? (LAB)
• Mg-Cl, NH, K-I, F-F

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Polar/Nonpolar molecules

• Effect of polar bonds on polarity of entire
  molecule
• Dipole moment
• One end partial positive, one end partial
  negative (HBr)
• Nonpolar molecules
• Polar bonds in opposite directions cancel out
  net nonpolar (CCl4)
• Like Dissolves Like

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Intermolecular Forces

• Van der Waals forces
• Weakest attractions between molecules
• 1) Dispersion forces
• Formed from temporary dipoles
• Br2, I2
• 2) Dipole-Dipole forces
• Occur when polar molecules attracted to one
  another
• Electrostatic attractions between partial
  positive and negative charges
• 3) Hydrogen bonding
• Hydrogen covalently bonded to very
  electronegative atom (F,O,N)
• Using IM Forces to explain molecular properties
• IM forces still weaker than ionic compounds
• Network solids all atoms covalently bonded to
  each other
• Diamonds (synthetic diamonds)

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Molecular Orbital Theory

• Molecular orbitals hybridization
• Overlap of atomic orbitals to form hybrid
  orbitals
• Methane (CH4)
• 1 2s and 3 2p orbitals form four sp3 orbitals
• Each sp3 orbital overlap with 1s orbital of H
  atom
• Ethene (C2H4)
• Sp2 orbitals formed from 1 2s and 2 2p
• 3 sp2 orbitals single bonds between C-C and C-H
• Nonhybridized 2p orbitals (1 each C) overlap to
  form pi bond double bond
• Ethyne
• Try as example
• Sigma and Pi bonds
• Sigma symmetrical to axis between atoms
• Part of single bonds, first part of double bonds
• Pi bonds weaker than sigma bonds
• Part of double and triple bonds
• Bonding and Antibonding orbitals
• Bonding orbital molecular orbital with an
  energy lower than atomic orbitals formed from
• Antibonding orbital molecular orbital with
  energy higher than atomic orbitals formed from

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Lab 25 Molecular Modeling

• Objectives
• Form 3-D models of covalent compounds
• Visualize the geometric shape of molecules
• Technique Check structures
• Safety Watch out for those atoms!
• Questions