1. Unit 9 Bonding

2. CH4
methane gas molecule
The attractive forces between atoms leads to chemical bonds that result in chemical compounds.
Ask the students how the valence electrons are shared in methane molecule.
You may need to remind the students that H has one valence electron while C has 4 valence electrons.

3. Why do atoms form bonds?
Atoms form bonds due to the need to have the most stable configuration for its electrons.
Tell the students that in order for sodium chloride to form, sodium atom loses its valence electron while chlorine atom (being more electronegative) gains the electron from sodium.
Atoms lose, gain, or share valence electrons in order to achieve a lower energy state (stable).

4. How atoms bond with each other depends on:
Electronegativity
Ionization Energy
# Valence Electrons e
I want an electron e e e

5. Metallic bonding Metallic bonds are formed by metal atoms.
Ask the student what they remember about the electronegativity and ionization energies of metals. (Low electronegativity and low ionization energy);
The valence electrons of metals are easily displaced or removed or delocalized.

6. Metallic bonding
Metallic bonding is the strong attraction between closely packed positive metal ions and a 'sea' of delocalized electrons.
What are the properties of metals?
How does metallic bonding affect the properties of metals?
The attraction between the metal ions and the delocalized electrons must be overcome to melt or to boil a metal. Some of the attractions must be overcome to melt a metal and all of them must be overcome to boil it. These attractive forces are strong, so metals have high melting and boiling points.
The delocalized electrons are also able to move through the metal structure. When a potential difference is applied, the electrons will move together, allowing an electric current to flow through the metal. Delocalized electrons make metals malleable and ductile.
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7. Metallic bonds What property of metals is illustrated above?
The illustration shows the malleability of metals. The atoms in metals slide past each other and held tightly by free-moving (sea of) electrons.
What property of metals is illustrated above?

8. Ionic Bonding an electrostatic force
Electrostatic refers to the attraction between opposite charges
Stronger than metallic bonds because of the opposite charges

9. Ionic bond
Formed by the transfer of electrons between a metal and a nonmetal
# of e- lost by metal = # of e- gained by nonmetal

10. Charge
The number of electrons that need to be lost or gained by an atom so it has the same electronic configuration as a noble gas.
When an atom gains or loses electrons, it is called an ION.
+ ions are cations
- ions are anions
The charge of the ion is called the OXIDATION NUMBER.

11. Watch This
HAVE STUDENTS LABEL OXICATION NUMBERS ON PT

12. Charge based on “copy cat” principle Sulfur wants to be like Argon (1s22s22p63s23p6)
Sulfur Atom 1s22s22p63s23p4
16 Protons (+)
16 Electrons (-)
0 No charge
Sulfur has 6 valence e- and will gain 2 more to complete an octet. The result is a:
Sulfur Anion 1s22s22p63s23p6
16 Protons (+)
18 Electrons (-)
2- Charge

13. Ionic Bonds
Strong electrostatic (positive-negative) force in ionic compounds makes a strong ionic bond.
How does the strong ionic bond affect the properties of ionic compounds?
Tell the students that it requires a lot of energy (high temperature) to break an ionic solid (or ionic bonds) into its component ions. As a result ionic compounds have:
a. High melting and boiling points
b. Brittle; the substance breaks into smaller units when hammered or turns into a powder
Ionic compounds also conduct electricity when dissolved in water (electrolytes) or when molten. It does not carry electrical charges (conduct electricity) in solid state.
Formula unit – smallest unit of an ionic compound; lowest whole number ratio of ions represented in an ionic compound

14. The greater the difference between eN values of 2 atoms is, the more ionic the bond will be. We call this “ionic character”.
= 3.2 very ionic

15. Ionic Bond
Na+ Cl-
Show the students the different representations for the formation of an ionic bond.
Use the model to the left to show how metals lose electrons to nonmetals.
Use the model to the right to illustrate how ionic bonds are represented as Lewis dot structures.

16. Writing Lewis Dot Structures
Element symbol represents the kernel (core) of the atom (nucleus and inner e-)
Dots represent the valence e-

17. Writing Lewis Dot Structures - Ionic
Metals tend to lose e- while nonmetals tend to gain electrons
Illustrate the formation of NaCl from Na and Cl atoms. Point to the lost of an Na electron forming Na+ while Cl gained the electron forming Cl-.
Have the students explain the formation of MgO and CaCl2 using the next two representaitons.
Ionic bonds
hyperphysics.phy-astr.gsu.edu

18. Writing Lewis Dot Structures – Ionic Bonds
Metals tend to lose e- while nonmetals tend to gain electrons
Illustrate how to write the Lewis Dot Structure of CaCl2. Emphasize the writing of brackets [] and ionic charges.
chemistry58.wikispaces.com
Ionic bonds

19. Properties of ionic compounds (all related to the strong attraction between the + and – charges)
high melting points and boiling points
hard solids
good conductors – in aqueous solutions and when molten
have a crystal lattice structure
Ions are here

20. Really, we don’t hate you.

21. Covalent Bond
Covalent Bond –formed when two nonmetals share pairs of valence electrons in order to obtain the electron configuration of a noble gas
Molecule - formed when two or more atoms bond covalently.
(A molecule is to a covalent bond as a formula unit is to an ionic bond.)

22. How covalent atoms bond

23. Diatomic Molecules
HOFBrINCl
Share electrons when they bond together

24. Polyatomic Ions covalently bonded group of atoms, with a charge
Watch this
They are listed on your STAAR chart. You will not have a bad time.

25. Properties of Covalent Molecules
Can exist as gases, liquids, or solids depending on molecular mass and polarity
Usually have lower MP and BP than ionic compounds of the same mass
Do not usually dissociate (break apart into ions) in water
Do not conduct electricity

26. How to draw Lewis dot structures for covalent molecules.
Write the formula for the compound.
Count the total number of valence electrons.
Predict the location of the atoms:
Hydrogen is NEVER the central atom.
If carbon is present, it is ALWAYS the central atom.
If there is only 1 atom of an element, it is the central atom.
The least electronegative atom is generally the central atom.
Place one electron PAIR between the central atom and each ligand (side atom) to “hook” the atoms together.
Dot the remaining electrons in pairs around the compound to complete the octet. Start with the ligands.
Check that each atom has an octet. (H only needs a pair, not an octet.)
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27. Lewis Structures for Molecules
Draw the Lewis dot structure for these molecules:
Hydrogen + Bromine (HBr)
Carbon + Chlorine (CCl4)

28. Writing Lewis Dot Structures - Covalent Bonds
Bonding e- Pairs
Lone Pairs
(nonbonding electrons)
Identify and describe to the class the lone pairs (non bonding electrons) and the bonding electron pairs or ligands. Provide more examples to your L-classes, if necessary.
Covalent bonds

29. Number of bonds
Single Bonds - when one pair of e- is shared between atoms
Double bond – when atoms share 2 pairs of valence electrons; ex. O2
Triple bond – when atoms share 3 pairs of valence electrons; ex. N2

30. Describing bonds Sigma bond - the first bond between 2 atoms
A single bond is a sigma bond.
Pi bond - the second bond between 2 atoms
A double bond consists of a sigma bond and a pi bond.
A triple bond consists of a sigma bond and two pi bonds.

31. Carbon can form single, double and triple bonds with itself.

32. Why are molecular shapes important?
The shape of a molecule plays a very important role in determining its properties.
Molecular shapes determine the properties of a substance. For example, hemoglobin performs an important function in respiration (transport of O2 and CO2). Sickle cell anemia results from malformed hemoglobin resulting to a change in the molecular geometry of hemoglobin, restricting the normal function of hemoglobin.
Properties such as smell, taste, and proper targeting (of drugs) are all the result of molecular shape.

33. VSEPR Theory also called electron geometry
Electron groups around the central atom will be most stable when they are as far apart as possible. We call this valence shell electron pair repulsion theory.
Because electrons are negatively charged, they should be most stable when they are separated as much as possible.
The resulting geometric arrangement will allow us to predict the shapes and bond angles in the molecule.

34. TO DETERMINE VSEPR SHAPE (electron geometry)
1) Draw the Lewis dot structure for the molecule 2) Identify the central atom 3) Count the number of electron groups around the central atom. 4) Look up the VSEPR shape on the chart. **shapes with no lone pairs are symmetrical **shapes with lone pairs are assymmetrical

35. Electron Groups
Each lone pair of electrons = one electron group on a central atom.
Each bond = one electron group on a central atom (regardless of whether it is single, double, or triple.) O N •
There are three electron groups on N
One lone pair
One single bond
One double bond

36. Two Electron Groups: Linear Electron Geometry
When there are two electron groups around the central atom, they will occupy positions on opposite sides of the central atom.
This results in the electron groups taking a linear shape.

37. Linear Geometry

38. Three Electron Groups: Trigonal Planar Electron Geometry
When there are three electron groups around the central atom, they will occupy positions in the shape of a triangle around the central atom.
This results in the electron groups taking a trigonal planar shape.

39. Trigonal Planar Geometry

40. Four Electron Groups: Tetrahedral Electron Geometry
When there are four electron groups around the central atom, they will occupy positions in the shape of a tetrahedron around the central atom.
This results in the electron groups taking a tetrahedral shape.

41. Tetrahedral Geometry

42. Molecular Geometry
The actual geometry of the molecule may be different from the VSEPR shape.
Lone pairs repel bonded atoms which distorts the expected shape.

43. Bond Angle Distortion from Lone Pairs
Electron Geometry: Tetrahedral Tetrahedral Tetrahedral
Molecular Geometry: Tetrahedral Trigonal Pyramidal Bent

44. Watch This

45. Predicting the Shapes around Central Atoms
Draw the Lewis structure.
Determine the number of electron groups around the central atom.
Classify each electron group as a bonding or lone pair, and count each type.
Remember, multiple bonds count as one group.
Look it up

46. Practice:
Determine the shape NF SiCl4
3. H2O

47. Types of Bonds:
Nonpolar covalent equal sharing of electrons between atoms; occurs between the atoms in a diatomic molecule (HOFBrINCl) and between C and H; ex. CH4
Video is 10:24

48. Polar Covalent
unequal sharing of electrons between atoms; occurs between two nonmetals or a nonmetal and a metalloid; ex. H2O
Electrons

49. Ionic complete transfer of electrons occurs between m/nm (ex. NaCl)
m/PAI
PAI/nm
PAI/PAI

50. This is a continuum. It describes the “ionic character” of the bond.
Bond type
Non-Polar Covalent Polar Covalent Ionic
NPC PC I
Difference in electronegativity values
Distance between atoms on the periodic table
Small medium big
This is a continuum. It describes the “ionic character” of the bond.

51. Practice: What type of bond exists in each of the following?
1. HCl 2. CaO
3. H2O
Br2

52. Water is a POLAR molecule
The more electronegative atom will have a slight negative charge, the area around the least electronegative atom will have a slight positive charge.

53. Symmetric molecules tend to be nonpolar Asymmetric molecules with polar bonds are polar
Symmetric means there are no lone pair around the central atom.

54. Bonding determines some physical properties
Type of Compound
Elements involve in bonding (metal/non metal)
Valence electrons are…
Melting /Boiling point
Electrical conductivity
Other properties
Metallic
Metal-metal
delocalized
high
Conductor
Malleable, ductile, shiny
Ionic
Metal - nonmetal
Lost/gained
Conducts in solutions or molten
Brittle, solid at room temperature
Covalent
Nonmetal - nonmetal
shared
low
Non-conductor
Mostly liquid or gas at room temperature
Type of Compound
Elements involve in bonding (metal/non metal)
Valence electrons are…
Melting /Boiling point
Electrical conductivity
Other properties
Metallic
Ionic
Covalent