Ionic and Covalent Bonding

1
Ionic and Covalent Bonding

• Chapters 6 7

2
Large Combination To Go
3
Bonding

• Why do atoms form bonds?
• Stability eight electrons in the outer levels
  makes an atom more stable.
• It has a filled outer orbital.
• Octet Rule
• Valence electrons

4
Energy in Reactions

• Nature works toward lower potential energy
• Rivers, center of gravity in a wire
• When two atoms for a bond, energy is released.
• Breaking bonds requires energy

5
Endothermic and Exothermic

• What is the difference between these two?
• How can you get an exothermic reaction?
• Endothermic?
• Weak bonds (non-stable) require little energy to
  break/make where strong bonds takes a lot.

6
Ions

• An ion is a charged particle
• What are the two types of charges?
• How can you get a charged particle?
• F-1

7
Ionic Bonds

• Electrons transferring from one atom to another
  can make both atoms more stable.
• When you lose/gain an e-, you become an ion.
• This is an ionic bond

8
Salts

• Salts are formed with completely ionic bonds
• Anion is the negative ion
• Cation is the positive ion
• These two ions are attracted to each other and
  thus an ionic bond is formed.

9
To Form or Not to Form
10
Meaningful Relationships
11
Naming Ionic Compounds

• NaCl
• Begin with the first element (cation)
• Sodium
• Drop the ending of the anion and add ide.
• Chloride
• Sodium Chloride

12
Practice

• LiF
• NaBr
• CaS
• SrO
• RbI
• Na2O
• CaCl2

13
Lattice Structure

• When Ionic Bonds form, they make a pattern.
• Show example NaCl
• These attractions between molecules give us a
  lattice structure.

14
A Different Twist On Things

• What do you breath in?
• O2
• How can oxygen bond with itself?
• Diatomic Bonds
• Which atoms?
• MEMORIZE
• Oxygen is a non-salt, which means it does not
  have completely ionic bonds.

15
Molecular Compounds

• Do not use ionic bonds
• Do not have a lattice structure
• They are separate molecules
• Have Covalent Bonds

16
Covalent Bonds

• How can O and O combine and both be happy?
• They can share an e-
• This bond is not as stable, what does that tell
  you about the energy in it?
• Mg strip

17
Covalent Bond

• We begin with two atoms coming together.
• There are forces of attraction between protons
  and e-
• There are forces of repulsion between protons and
  protons and e- and e-
• As long as the attractions are greater than the
  repulsions, it will make a bond

e-


e-
18
Covalent Bonds

• They can be seen as a spring
• Bond Length is the average distance between the
  nuclei
• This is not fixed because the bonds vibrate back
  and forth
• Bond Energy The energy stored in the bond is
  proportional ot the bond length This is the
  energy required to break the bond.

19
Electronegativity

• Who wants the electron more?
• If you have O and O, who wants the electron more?
• If you have N and O, who wants it more?
• The difference in electronegativity determines
  the type of bond!!

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Electronegativity
Completely Covalent
Completely Ionic
Polar Covalent
0
0.4
2.1
4.0
Electronegativity Difference
21
Polar Covalent

• This is an uneven sharing of electrons
• One atom wants it more (has a higher
  electronegativity)
• You get slightly positive and slightly negative
  atoms

22
What kind of Bond is

• NaCl
• CO2 (the C-O bond)
• N2
• CaO
• TiCl2
• You can mainly tell by the chart.

23
Oxidation Numbers

• It is hard to tell when something covalent is
  going to bond. You dont have a charge
• We assign one to make our molecules.
• They work a lot like charges on ions.
• There are rules you need to follow.

24
Rules for Assigning Oxidation Numbers

• Pure elements have oxidation numbers of zero
• The more electronegative atom is assigned a
  number like it was an anion and the less
  electronegative is assigned as if it were a
  cation
• Fluorine is always 1
• Oxygen is always 2, unless a peroxide or with
  fluorine
• Hydrogen is always 1 with elements more
  electronegative than it, it is 1 with metals.
• The sum of all oxidation numbers has to be zero
• Unless polyatomic ions, in which case it is the
  charge
• We can also use oxidation numbers for ionic bonds

25
Practice

• HCl
• CF4
• PCl3
• SO2
• HNO3
• KH

26
Comparing Ionic and Covalent Compounds
27
Practice

• H2O
• Cl2
• CH3F
• NH3
• CH3OH
• H2SO4
• C6H6
• NaOH

• Make these compounds.
• Draw these compounds
• Label each bond with the correct type

28
Lewis Structures

• Showing how atoms/molecules look
• Determine the number of valence electrons
• Arrange atom symbols and show bonding as dots
• Are the octets full?
• Change shared pairs to bonds, what are unshared
  pairs doing?

29
Practice

• HCl
• HBr
• CH2Cl2
• CH3OH
• C6H12
• CO2
• C2H2

30
Determining the Middle Atom

• H2 and halogens usually are on the ends
• Atoms with the smallest electronegativity are
  often in the center
• Atoms usually surround the middle atom.
• You can have double and even triple bonds
• CO2 or N2
• Sometimes no Lewis structure is correct
  Resonance (NO2)

31
Special Structures

• Polyatomic Ions can be left in their parentheses
• Benzene Ring

32
Molecular Geometry

• VESPR Theory repulsions between sets of valence
  level electrons causes them to be oriented as far
  apart as possible.
• Shapes
• Linear
• Bent or angular
• Trigonal planar
• Tetrahedral
• Trigonal Planar

33
Intermolecular Forces

• If you have a slightly positive and slightly
  negative atom, you can get attractions between
  them
• Not like ionic bonds, but still some attraction
• Dipole atoms Water
• These are weak forces

34
Hydrogen Bonding

• This is actually a fairly strong intermolecular
  force.
• Hydrogen bonds intermolecularly to non-shared
  electrons very well
• Water and surface tension

35
Metallic Bonding

• This the the bond that results from metal atoms
  bonding together.
• What are properties of metals?
• Electron ocean

36
Tin Oxide Lab