1. Metallic & Ionic Bonding
Test Date:
Thurs, Feb 18

2. conductive malleable ductile Properties:
Metallic Bonds
Properties:
electricity flows through them: __________
can be hammered into sheets: _______
can be drawn into a wire: _______
conductive
malleable
ductile

3. Delocalized electrons Electrons do not belong to any one atom
Metallic Bonding
Attraction of a metallic cation (+ ion) and valence electrons that are free to move (delocalized)
Delocalized electrons
Electrons do not belong to any one atom
Sometimes called a “sea of electrons”

5. Metallic bond = a positive core surrounded by a “sea” of electrons

6. Properties of Ionic Compounds
Most form solid crystals (called “salts”) at room temperature
Have High melting points (strength of electrostatic attraction)
Form electrolytes (ions in solution) which can conduct electricity when dissolved in water (aqueous).
When dissolved the electrons are permitted to move freely (compare to ionic compound crystal structure)

7. Ionic vs Metallic = + + =

8. Ions are surrounded by ions of opposite charge.
A crystal lattice is formed.

9. loses positive cation gains negative anion
Ionic Bonding
Involves the transfer of electron(s) from a metal to a nonmetal.
The metal _____ electrons becoming ________ and is called a _______.
The non-metal _____ electrons becoming _______ and is called an _______.
loses
positive
cation
gains
negative
anion

10. Electrostatic Interactions
“Opposites attract”
Cations and anions are attracted to each other through strong electrostatic attractions – stronger than metallic bonds
The result is a stable, low energy ionic compound
Ionic compounds are neutral since the charges of the cation and anion cancel

11. Valence Electron Dot Structures
for Ionic Compounds
Example 1: Na & Cl
Determine the number of valence electrons for each element
From the valence count, draw the Lewis valence electron dot structures of Na and Cl
Show how the electron is transferred from Na to Cl to form ions

12. So what IS an ionic bond? Ionic compounds do not actually bond.
Electrons are completely gained
or lost (transferred), NOT
shared!!!!
The electrostatic attraction is the ionic “bond.”

13. Valence Electron Dot Structures for Ionic Compounds
Example 2: Mg & O Example 3: Ca & Cl

14. Why do atoms gain or lose electrons?
more stable
In order to become _________
What makes something stable?
It is lower in energy
It has a full valence shell
It’s electron configuration is the same as a _________ ________
noble gas

15. Every atom wants to be like a noble!
Stability: Atom vs. Ion
Which is more stable, Na or Na+1?
Every atom wants to be like a noble!

16. Electronegativity in Ionic Bonds
What does electronegativity mean?
Attraction for valence electrons in a bond
What element is the MOST electronegative?
Fluorine (F)
What element is the LEAST electronegative?
Francium (Fr)

17. Electronegativity in Ionic Bonds
The larger the difference between the EN values in a compound, the stronger the ionic bond.
What compound would have the STRONGEST ionic bond that is possible?
FrF
I’m strong; does that mean I’m ionic?

18. Electronegativity Values

19. What is a formula unit? One single unit of the ionic compound
Smallest ratio of cations to anions that form an electrically neutral compound
Ex. NaCl vs. MgCl2

20. Common Polyatomic Ions
Ammonium
Acetate
Carbonate
Hydroxide
Nitrate
Nitrite
Permanganate
Phosphate
Sulfate
Sulfite
(NH4)+1
(C2H3O2)-1
(CO3)-2
(OH)-1
(NO3)-1
(NO2)-1
(MnO4)-1
(PO4)-3
(SO4)-2
(SO3)-2

21. How does the charge work?
Ignore the subscripts and treat the group
as one thing:
(NO3) -1 total charge is -1
(SO4) total charge is -2
(PO4) total charge is -3

22. Write the formula for these compounds: 1. Ca2+ + (NO3) 2. Cr 4+ + Cl- 3. NH4 + O2- 4. Cu2+ + S2- 5. Ni2+ + N3- 6. Ag+ + O2-
Ca(NO3)2
CrCl4
(NH4)2O
CuS
Ni3N2
Ag2O

23. Ionic Bonding Recap Questions Copy the question and answer in complete sentences
What is the total charge on an ionic compound?
What holds the anions to the cations? In other words, how is an ionic “bond” formed?
List the pair(s) of elements below that are likely to form ionic compounds? Remember that ionic compounds are formed between metals and nonmetals. Write the balanced formula.
Cl, Br Li, Cl K, He I, Na
What properties characterize ionic compounds?
Why do ionic compounds conduct electricity when they are melted or dissolved in water?
What is a formula unit? What does it represent?

24. Covalent Bonding

25. Ionic Compound Properties (Review)
Exhibit a crystal structure
Exist as solids
Dissolve easily in water
Have high melting and boiling points
Conduct electricity in solutions
Have high electronegativity differences

26. Covalent Compound Properties
Properties of Covalent Compounds…
Have low melting and boiling points
Do not dissolve easily in water
Do not conduct electricity in solutions
Exist as gases, liquids, or solids
Have low electronegativity differences

27. Common Covalent Compounds
Nitrous oxide (N2O) is laughing gas used as an anesthetic and to boost auto engine power
Methane (CH4) is a flammable gas used as fuel and in homes for domestic heating and cooking purposes.
Sugars, like sucrose (C12H22O11) and glucose (C6H12O6), are used in food and energy production
Hydrogen peroxide (H2O2) is used as a bleaching agent, emetic (induces vomiting), and an antiseptic (clean cuts and scrapes)

28. Common Covalent Compounds
Carbon dioxide (CO2) is a gas used by plants during photosynthesis, produced in respiration; it is a byproduct of burning fossil fuels, and it has many uses as a fire extinguisher, refrigerant (dry ice), and carbonation in drinks

29. Common Covalent Compounds
Water (H2O) is a liquid that is vital for life. It covers about 71% of earth’s surface, and makes up about 60% of the human body. It is known as the universal solvent, and has many unique properties and uses.

30. Diatomic Elements
Some nonmetal elements on the periodic table exist in nature only as pairs called diatomic molecules; bonded covalently
H, O, F, Br, I, N, Cl
Hydrogen Molecule (H2)

31. Covalent Bonds
Occurs when valence electrons are shared between atoms – usually between nonmetallic elements
Covalently bonded compound known as a molecule
These shared electrons are part of the valences of all atoms involved (satisfies octet rule)
Since electrons are shared, no charges appear
Many combinations can occur between two nonmetals
Example: carbon and oxygen can form carbon monoxide (CO) and carbon dioxide (CO2)
Carbon Dioxide

32. Rules for Lewis Structures
1. Make certain that the bond is a covalent bond by checking the electronegativity difference. Then set up the skeleton structure as follows:
The atom with the lowest electronegativity will tend to go in middle
Place all the other atoms around this central atom
Attach these atoms to the central atom in reasonable fashion with single bonds

33. Rules for Lewis Structures
2. Sum valence electrons
3. Complete octets of peripheral atoms
4. Place leftover e- on central atom
5. If necessary use multiple bonds to fill the center atom's octet.

34. Write the structural formulas for:
Cl2
NH3
CH4
H2O

35. Polyatomic Ions (K ONLY)
Draw the electron dot structures for:
CO32-
NO3-
SO32-
SO42-

36. Single Covalent Bonds Ex: Methane (CH4)
Atoms of some elements attain a noble gas configuration by sharing one pair of electrons between two atoms
When drawing these molecules, a line can represent a pair of shared electrons
Ex: Methane (CH4)

37. Multiple Covalent Bonds
Sometimes, atoms share more than one pair of electrons between two atoms
When drawing these molecules, multiple lines can represent pairs of shared electrons
Ex: diatomic oxygen (O2) and nitrogen (N2)
How many pairs of electrons does each kind of bond share?
Single: _________ Double: _________ Triple: _________
1 pair
2 pairs
3 pairs

38. Draw the electron dot structures for the following
CO2
C2H4
NO2
SO2
CS2

39. Exceptions to Octet Rule
Beryllium
Has 2 valence electrons
Full with 4 valence electrons
Ex: BeI2
Aluminum
3 valence electrons
Full with 6 valence electrons
Ex: AlCl3
Boron
Ex: BH3

40. Exceptions to Octet Rule Cont’d
Expanded octets
Some atoms bond so they have more than 8 electrons in their valence
Occurs only around a central nonmetallic atom from period 3 or higher (most commonly sulfur & phosphorous)
Ex: PCl5 and SF6
Trigonal Bipyramidal
Octahedral

41. Polar & Nonpolar Covalent Bonds
Covalent bonds share electrons equally and unequally (“tug of war” with electrons)
Which molecule appears “balanced”?
Which one does not?
H2O O2
Water (H2O)
Oxygen (O2)

42. Polar Covalent Bonds Unequal sharing of electrons
Occurs when atoms that share electrons have different electronegativities
Shared electrons spend a greater amount of time at the more electronegative atom
Example: water (H2O)

43. Nonpolar Covalent Bonds
Equal sharing of electrons
Occurs in two scenarios…
No difference in electronegativity between atoms that are sharing (diatomic molecules)
Example: chlorine (Cl2)
Shape of molecule pulls electrons equally in all directions
Example: methane (CH4)
Is this even correct?

44. Bonding & Electronegativity
Atoms bond according to the difference in their electronegativities
Look at table of EN values and find the difference between the two atoms involved (absolute value)
If EN difference is…
< 0.5 = nonpolar covalent bond
0.5 – 1.7 = polar covalent bond
>1.7 = ionic bond
<0.5

45. Bonding & Electronegativity

46. Bonding & Electronegativity
Try to predict the bond type (ionic, polar covalent, nonpolar covalent) using your EN table:
Formula
EN values
EN Difference
Bond Type
LiCl
Li = 1.0
Cl = 3.0
2.0
ionic bonds
PCl3
MgO
CH4
NF3
P = 2.1 Cl = 3.0
0.9
polar covalent bonds
Mg = 1.2 O = 3.5
2.3
ionic bonds
C = 2.5 H = 2.1
0.4
nonpolar covalent bonds
N = 3.0 F = 4.0
1.0
polar covalent bonds

47. Molecular Shapes 2 linear 180° 3 trigonal planar 120° 1 bent <120°
# of Electron
Groups
Number of
Lone Pairs
Electron Pair
Arrangement
Molecular Geometry
Approximate
Bond Angles 2
linear
180° 3
trigonal planar
120° 1
bent
<120° 4
tetrahedral
109.5°
trigonal pyramid
<109.5° (~107°)
<109.5°(~105°)

48. Molecular Shape Try to predict the shape using VSEPR
and your shape chart:
1) BCl3 2) CH4 3) NH3
Lewis Dot
Central Atom
# e- pairs
# bonded pairs
# lone pairs
Shape

49. Molecular Shape Example: What is the shape of water (H2O)?
1) Central atom? oxygen
2) Draw Lewis Dot…
3) # of total electron pairs around central atom? 4
4) # of bonding pairs around central atom? 2
5) # of lone pairs around central atom? 2
6) Refer to chart to identify shape…angular (bent)

50. Molecular Shapes Applies to covalent compounds only.
Use VSEPR steps (valence shell electron pair repulsion):
1) Identify the central atom as the element that can form the most bonds
2) Draw the Lewis dot structure for the molecule
3) Count total # of electron pairs around the central atom
4) Count # of bonding pairs of electrons around the central atom
5) Count # of lone pairs of electrons around the central atom
6) Look at summary chart, identify shape

51. Polar & Nonpolar Molecules
Polar molecules have a partial positive and partial negative charge – called a dipole
Dependent upon location and nature of the covalent bonds they contain
Molecules are considered polar if they:
Have a polar covalent bond
- EN difference between
Have an asymmetrical shape
- trigonal pyramidal
- linear (with 2 different elements)
- angular (bent)
If both criteria are not met, it is either a nonpolar molecule or an ionic compound

52. Polar & Nonpolar Molecules
*IMPORTANT: Polar & nonpolar molecules are different from polar & nonpolar bonds!!!
Non-polar molecules have charges that are evenly distributed, due to shape (ex: any diatomic molecule, gasoline)
Polar molecules have a partial positive and a partial negative charge (ex: water)

53. Polar & Nonpolar Molecules
Evaluate and identify each molecule as polar covalent or nonpolar covalent:
Formula
Has a polar covalent bond
Has an asymmetrical shape
Identity
H2O 
polar molecule
CH4
BCl3
HCl x
nonpolar molecule
nonpolar molecule x 
polar molecule 

54. Property Ionic Bonds Covalent Bonds Electrons are: Transferred Shared
Difference in Electronegativity:
<0.5 nonpolar
0.5 to 1.7 polar
>1.7
A metal and nonmetal
Bond between:
2 nonmetals
State at room temperature
Solid
Solid, Liquid, or Gas
Particle Name:
Formula Unit
Molecule
High
Low
Melting Point:
Conducts Electricity?
Yes No
Dissolves in water?
Yes
No (usually)
Yes No
Flammable
Forms a more stable configuration?
Yes
Yes
Examples:
NaCl, MgI2
NH3, CHCl3

55. Types of Multiple Bonds
Sigma Bond: covalent chemical bond where ends of orbitals overlap; present in single and multiple bonds
Pi Bond: covalent chemical bond where parallel orbitals overlap to share electrons; only present in multiple bonds
Type of Bond
Sigma Bonds
Pi Bonds
Single 1
Double
Triple 2

56. Orbital Hybridization - Boron

57. Orbital Hybridization
Orbital Hybridization: 2 or more different types of orbitals (usually s, p, or d) mix to become the same number of identical orbitals around the central atom
2 bonds
sp hybridization
2 equivalent orbitals
3 bonds
sp2 hybridization
3 equivalent orbitals
4 bonds
sp3 hybridization
4 equivalent orbitals
5 bonds
sp3d hybridization
5 equivalent orbitals
6 bonds
sp3d2 hybridization
6 equivalent orbitals
*Lone pairs and sigma bonds count as equivalent orbitals

58. Orbital Hybridization - Carbon