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Chemical Bonding. Chem. Bonding Asgns:. 30 - 6.1: 177/1-6 6.2: 189/1-5 31 - 6.3: 194/1-5 6.4: 196/1-3 32 - 6.5: 207/1-6 33 - 209/1-7 34 - 209/8-15 35 - 209/16-24 36 - 210/25-32 37 - 210/33-42 38 - 211/43-49. Review. Ionization Energy.

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chem bonding asgns
Chem. Bonding Asgns:

30 - 6.1: 177/1-6 6.2: 189/1-5

31 - 6.3: 194/1-5 6.4: 196/1-3

32 - 6.5: 207/1-6

33 - 209/1-7

34 - 209/8-15

35 - 209/16-24

36 - 210/25-32

37 - 210/33-42

38 - 211/43-49

ionization energy
Ionization Energy
  • Ionization Energy – energy needed to remove an e- from the outermost shell of a neutral atom
    • Low i.e. means easier removal of e- and a resultant positively charge cation will be formed
    • Trend: greater i.e. up and to the right on the periodic chart
electron affinity
Electron Affinity
  • The energy released or absorbed by a neutral atom from the acquisition of an electron to its outer shell.
    • High e.a. means it’s easier to accept an e- and a resultant negatively charged anion will be formed.
    • Trend: highest e.a. at top of group because of increased nuclear attraction for acquired e- and highest at right of each period because atoms reach stability by achieving stability because of the acquisition
lattice energy
Lattice Energy
  • The energy released when two elements combine during the formation of a compound due to electrostatic interactions forming the molecule’s physical structure.
  • Ions of unlike charge are attracted to one another.
  • Equilibrium is reached and a lattice is formed (as in the picture at upper right).
  • Energy released in formation of the compound is equal to the energy needed to break apart the compound into its component elements.


ionization energy1
Ionization Energy
  • A measure of the ability of an atom or ion to hold onto electrons.
  • Trend: i.e. increases up and to the right.
    • Electrons are held more tightly by positive ions.
    • Electrons are held less tightly by negative ions.
types of chemical bonds
Types of Chemical Bonds
  • Covalent – electrons are shared
    • Non-polar – e’s shared equally
  • Polar covalent – unequal sharing
  • Ionic – electrons are transferred
  • Coordinate covalent
  • Radical
  • Metallic
1 non polar covalent bonds
1. Non-Polar Covalent Bonds
  • These occur primarily between two non-metallic elements, especially the diatomic gases (H, N, O, and the Halogen family)
  • There is an equal sharing of the valence electrons so that both atoms achieve octet and chemical stability
  • e. n. falls between 0.0 and 0.4


2 polar covalent bonds
2. Polar Covalent Bonds
  • These occur primarily between two active non-metals or between a moderately active metal and a non-metal. There is an unequal sharing of the valence electrons.
  • e.n. difference is 0.5 to 1.6 range.

Electrons are held closer to the Oxygen because O has greater e.n.


another way to show polar covalent
Another way to show polar covalent

d means slightly – or +


main bond types
Main Bond Types


3 ionic bonds
3. Ionic Bonds
  • These occur primarily between active metals and active non-metals. The lesser electronegative element actually transfers one or more valence electrons to another atom – the more e.n. element.
  • Range of e.n. difference is 1.7 to 4.0.
  • See prior slide for sample.
ionic bond sample
Ionic Bond Sample
  • Source:
4 coordinate covalent bonds forming polyatomic ions
4. Coordinate Covalent Bonds forming Polyatomic Ions
  • These occur primarily between two non-metals one of which is usually oxygen. The lesser e.n. element provides both shared electrons.
  • Always results in the formation of a polyatomic ion – usually an anion or radical.

nitrate ion, NO31-

phosphate ion, PO43-

5 radical bonds
5. Radical Bonds
  • Occur mainly between a metallic cation and a radical anion. Cation transfers e- to the central atom of the radical which shares these electrons with its combining atoms by coordinate covalent bonds.
  • Contains both ionic and covalent bonds.
  • Always forms a stable multi-atomic compound of at least 3 different elements.
radical bond samples
Radical Bond Samples

Sodium Phosphate


Sodium Nitrate


6 metallic bonds
6. Metallic Bonds

Occur only between metals during the formation of alloys. The metallic kernel, composed of metallic nuclei and their inner shell electrons, is surrounded by a “sea” of valence electrons that flow across and about the kernel.

Metal cation

- electrons that are delocalized or free to move about



Metals are shiny because they absorb light, exciting electrons to higher energy levels. The e’s immediately fall to lower levels, emitting light energy and making the metals appear shiny.

  • Ductile
  • Malleable

Relationship Between Electronegativity(See page 4 of note handout

Difference and Ionic Characteror page 161 in book)

Electronegativity Percentage of

Difference Ionic Character

0.2 nonpolar 1

0.4 covalent 4



0.6 9

0.8 polar 15

1.0 covalent 22

1.2 bond 30

1.4 39

1.6 47

1.8 55

2.0 63

2.2 70

2.4 ionic 76

2.6 bond 82

2.8 86

3.0 89

3.2 92

sample e n problems
Sample e.n. problems

See pages 4, 5, & 6 of note handout or frames 20 and 21

Given: As2S3

e.n. difference:

2.44 – 2.20 = 0.24

Bond Type:

Non-polar covalent

Given: CaF2

e.n. diff.:

4.10 – 1.04 = 3.06

Bond Type:


  • #30 177/1-6 and #33 207/1-7
covalent bond characteristics
Covalent Bond Characteristics
  • Atoms bond to become more stable by getting a full outer energy level (Octet for all except H and He).

From table at left notice that shorter bond lengths require a greater energy to break the bond and form neutral isolated atoms.

Notice how in the figure at left how a bond is formed between 2 hydrogen atoms to form a stable 1s2 configuration of H2.

Source: Modern Chemistry, 2006 ed.


There are exceptions to the octet rule: some compounds formed with F, O, and Cl. This is called expanded valence.

Source: Modern Chemistry, 2006 ed.

electron dot notation
Electron-Dot Notation

Dots are placed around the symbol of an element to represent its number of valence electrons.

Source: Modern Chemistry, 2006 ed.

lewis structures
Lewis Structures
  • Formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes represent electron pairs in covalent bonds. Adjacent dots represent unshared electrons. (Structural formulas show dashes only: F-F or H-Cl)
drawing a lewis structure
Drawing a Lewis structure
  • Determine type and number of atoms in the molecule
  • Write electron-dot notation for each atom type
  • Determine total numbr of valence electrons
  • Arrange atoms. C is central atom if present; or least e-n atom is central (not H).
  • Arrange electrons to get 8 around each atom (except H)
nh 3 co 21



  • #30 189/1-5
  • #34 209/8-15
  • #35 209/16-24
  • #31 194/1-5 and 196/1-3
resonance structures
Resonance Structures

Molecules or ions that cannot be correctly represented by a single Lewis structure. Below are three possibilities for NO3-1 - the nitrate ion.

ionic bonding ionic compounds
Ionic Bonding & Ionic Compounds
  • Formed by electrons being transferred. The number of positive and negative ions are equal (no charge). Simplest form is a formula unit.




info on ionic bonds
Info on Ionic Bonds
  • They form a crystal lattice which minimizes their potential energy.
  • Stronger bonds than in covalent compounds producing higher boiling or melting points than cov.
  • Hard but brittle.
  • Do not conduct in solid state, but do conduct in molten state or when dissolved in water.
hybridization and molecular geometry
Hybridization and Molecular Geometry
  • VSEPR theory is used to predict shapes of molecules based on the fact that unshared electron pairs strongly repel each other
  • Hybridization theory is used to predict the shapes of molecules based on the fact that orbitals within an atom can mix to form orbitals of equal energy.
  • Bottom line: bonding in atoms depends on the central atom’s ability to maximize the spread of its valence electrons into adjacent orbitals within the same energy level.

VSEPR – valence shell electron pair repulsion – determines the geometry (shape) of molecules.

You have this sheet (p. 7) in your Ch 6 handout…

It gives you compound samples such as BeF2 and BF3 and shows you electron pairs that are shared and unshared, hybrid type, angles and geometry.


Page 9 of your Ch6 handout:

Molecular Type – AXE format

A = central atom

X = shared pairs of electrons (shown as B in your book)

E = unshared pairs of electrons

Shared pairs = number of bonding atoms united with the central atom of the molecule

3 D Diagram meanings

shared pairs

above, below or to the side of the central atom

behind of central atom

in front of central atom

.. lone pairs of electrons

dipole dipole forces
Dipole-Dipole Forces
  • A dipole is created by equal but opposite forces that are separated by a short distance


Hydrogen bonding – between H of one molecule of water and O of another water molecule. Shown by dashed line.

  • Source:
  • H



A polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons. Making O2 dissolve in H20, for instance. See page 206 in book.

  • London dispersion forces cause intermolecular attractions as a result of constant motion of electrons and creation of instantaneous dipoles.
  • #32 207/1-5
  • #36 210/25-32
  • #37 210/33-42
  • #38 211/43-49
  • Chem review from Prentice Hall.