1. Regents Chemistry Bonding
Unit 5

2. 2 Basic Principles
Systems at low energy are more stable than systems at high energy levels.
Chemical changes will occur if the changes lead to lower energy levels.

3. Chemical Bond
A chemical bond results from the simultaneous attraction of electrons to two nuclei.
Chemical energy is potential energy stored in chemical bonds. When atoms are bonded, they generally are at lower energy. When the bond is formed energy is released.
Energy must be added to break a chemical bond.
Energy will be released when a bond is formed.

4. Recall this information from the Periodic Table Unit
A noble gas configuration is more stable than other configurations.
In order to become more stable, all atoms would like to achieve a complete outer shell.
In order to complete their outer shell, an atom may gain, lose, or share electrons.
In order to remove an electron, energy must be added. (ionization energy)
When an atom gains an electron to complete its outer shell, energy will be released.
Define ion: an atom or group of atoms which has a charge because electrons have been lost or gained.

5. Types of Bonds 1. Ionic Bond 2. Polar Covalent Bond
3. Non-polar Covalent Bond
4. Metallic Bond
Coordinate Covalent Bond

6. Ionic Bond (electrovalent bond)
Bonds formed when electrons are transferred from metals to nonmetals.
Metals Nonmetals
Insert video

7. Covalent Bond Covalent bonds are formed when atoms share electrons.
Both nuclei will have an attraction for the electrons--they both have high electronegativities.
Non-metals form covalent bonds.
insert video

8. More review from the Periodic Table Unit
Electronegativity: measure of the ability of an atom to attract electrons.
First Ionization energy: amount of energy needed to remove the most loosely held electron.
An atom that is likely to gain electrons will have a high electronegativity and a high ionization energy.
Why? Strong forces keep the electron bound to the atom.
Most electronegative is F

9. Even More from the Periodic Table
Define ion: an atom or group of atoms which has a charge because electrons have been lost or gained.
An atom which gains electrons will form an ion with a negative charge because it has an excess of electrons.
Define anion: a negatively charged ion.
An atom which loses electrons will form an ion with a positive charge because it has a deficiency of electrons.
Define cation: a positively charged ion.
An atom that loses electrons will have a low electronegativity and low ionization energy.
Weak forces holding the electrons allow them to be easily removed.

10. Determining Bond Types
Calculate the electronegativity difference.
Table S is where you find electronegativity.
Nonpolar Covalent Bond--equal sharing of electrons, usually two identical atoms (H2, N2 and other diatomics).
Polar Covalent Bond--unequal sharing of electrons. The electron is more strongly attracted to one of the atoms in the bond.

11. Identifying Bond Types
Electronegativity Difference
Bond Type
Example
Nonpolar covalent bond H2
0 to 1.7
Polar covalent bond
HCl
1.7 and greater
Ionic bond
NaCl
Draw the Lewis Diagrams for each example.

12. Identifying Bond Types
Atoms
Electro-
negativity of each element
negativity difference
Bond Type
Li and Br Li Br
2.0
Ionic bond
H and O H O
1.4
Polar covalent bond
N and N N
Nonpolar covalent bond
C and O C
0.9

13. Identifying Bond Types 2
Atoms
Electro-
negativity of each element
negativity difference
Bond Type
K and F K F
3.2
Ionic bond
N and O N O
0.5
Polar covalent bond
Na and Cl Na Cl
2.3
S and O S
0.9

14. Metallic Bonds bonds between metals
Why: metals and metallic atoms have loosely held electrons that can be taken away fairly easily.
These electrons are more or less free to move from one atom to another.
Chemists often describe metals as:
metal ions floating in a “sea of mobile electrons” around a positive nucleus.

15. Properties of Metals High melting point
Solid (hard) at room temperature
Alloys: homogeneous mixture of metals
(ex. Steel)
Good conductors of heat and electricity
Ductile
Malleable
Have luster (shiny)

16. Metallic Bonding
Metallic bonding results in the formation of alloys rather than compounds because metal atoms do not combine in fixed ratios.
Only ionic and covalent bonding results in the formation of compounds

18. Determine the following Bond Types
HCl
NaCl

19. Bonding Unit -Learn to draw Lewis Structures.
-Define Ionic, Polar and Nonpolar Covalent Bonding, metallic bonding.
-Discuss ionic, covalent and metallic properties.

20. Lewis Dot Structure
In 1902 Gilbert Newton Lewis invented the valence bond theory.
Lewis came up with an easy way to represent electrons in the outer shells of ions. His invention is called “Lewis Dot Symbols”.
Lewis structures are used to visualize the valence electrons of elements.
Lewis was a famous American physical chemist.

21. Lewis' Paper of 1916
In this paper, Lewis begins by using cubes, but he moves away from them by the end of the paper. Here is how he visualized the elements lithium through fluorine:

22. Basic Chemical Bonding
Lewis Dot Structure
Basic Chemical Bonding

23. Lewis Dot Structures
Here are the Lewis dot structures for some elements. The transition metals, lanthanides and actinides are not displayed by Lewis Dot structures because they do not follow the octet rule. Their first outer orbital shell only has the capacity for two electrons.

24. Lewis Dot Structure
1)Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center.
2)Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.
3)Complete an octet for all atoms except hydrogen and helium.
4)If structure contains too many electrons, form double and triple bonds on central atom as needed.

25. Lewis Dot Structure Example:
Sulfur has 16 electrons and on the Regents Tables you can see the electron configuration is There are 6 electrons in the outer shell of sulfur.

26. Lewis Dot Structure
This is what sulfur looks like according to the Lewis Dot Diagram: S

27. Lewis Dot Structure
Now it’s your turn to try and draw some elements using the Lewis Dot Structure Potassium Germanium Phosphorus 4. Neon 5. Aluminum

28. Lewis Dot Structure

29. Four Types of Chemical Bonds
Ionic
Polar covalent
Nonpolar covalent
Metallic

30. Ionic Compounds
Ionic compounds are composed of both metals and nonmetals. The bond that is formed is based on electrostatic forces between negatively(anion) and positively(cation) charged ions.
Ionic bonding occurs by the transfer of one or more electrons from one atom to another
Electrostatic forces remember are the attractions between positive and negative ions that hold the bond together.

31. Properties usually solid at room temperature have high melting points
usually do not conduct electricity as a solid
usually dissolve in water
usually conduct electricity when in solution or molten state
1+2-Strong bonds are formed so it is hard to break them apart
3-ions can’t move, and when compound is formed it doesn’t want to give away any electrons
4-since water is a polar molecule it attracts ions

32. Ionic Bonds
An ion is an atom or group of atoms that have a charge. Atoms normally have a neutral charge because most often they have the same number of electrons and protons. They become ions by the loss or addition of one or more electrons. This process is called ionization.
To understand ionic bonding we will develop an understanding of ions.
An ion that has more electrons than protons is called an anion, and an ion that has fewer electrons than protons is called a cation.

33. Ionic Bonds
The interaction of ionic bonds is when atoms gain or lose electrons until the outer shell of electrons is full and stable with 8 electrons. This is part of the octet rule.
Recall octet rule:
When atoms combine to form molecules they generally each lose, gain, or share valence electrons until they attain or share eight and reach a noble gas electron configuration
Stress that the octet rule is useful but it does not work in every case, and there are exceptions to the rule.

34. Ionic Bonds
The number of electrons the atom gains, loses, or shares is called its Valence.
Nonmetals usually have four or more electrons in their outer shell. To make their outer shell full, it’s easier(it takes less ionization energy) for them to gain three or four electrons than to lose four or five electrons.
When you look at the metals, they usually have three or less electrons in their outer shell. Opposite from nonmetals, it is easier for metals to lose three or less electrons than to gain four or more.
Therefore it makes sense that metals and nonmetals bond together easily.

35. Ionic Bond
An ionic bond is the resulting attraction for an anion and a cation after an electron is transferred from the metal to the non-metal.

38. Ionic Bonds in a crystal of NaCl(S)

39. Covalent Compounds
Covalent compounds are made up of two nonmetals. A single covalent bond is formed when a pair of electrons is shared between two atoms
There are two types of covalent bonds: non-polar covalent and polar covalent.

40. Covalent Bonds
A covalent bond exists when two electrons are shared by two nonmetals.

41. Formation of H2 Nonpolar covalent equal sharing of electrons

42. Formation of I2

43. Formation of HCl polar covalent bonding unequal sharing of electrons

44. Polar Bond vs. NonPolar Bond
The nonpolar covalent bond has an equal sharing of electrons, while the nonpolar covalent bond has an unequal distribution of charge.
A ΔEN of 0 indicates a nonpolar covalent bond.

45. H2O has 2 polar covalent bonds
Polar molecules must have polar bonds, and they must be spaced so the molecule has an uneven distribution of charge.

46. Water is a bent molecule with polar covalent bonds

47. Carbon dioxide is a linear molecule with polar covalent bonds.

48. Example-HF F H
e- rich
e- poor d+ d- H F

49. Properties
simple molecular substances have low melting and boiling points
larger more complex compounds will have higher melting and boiling points
usually do not conduct electricity as a solid or when molten or in solution
usually do not dissolve in water
1-weak intermolecular bonds
2-many strong covalent bonds
3-no mobile electrons, no way to transfer electrons
4-molecules are not charged, thus polar water
molecules do not attract them

50. Non-Polar Covalent Bonding
Accounts for the bond that keeps two atoms of the same element together. (Cl2, H2) diatomic elements!
Atoms share electrons from ½ filled orbitals to achieve noble gas configurations
Shared electrons are attracted to both nuclei, which keeps atoms together
Electrons involved in bonding are called shared electron pairs, ones that are not are called lone electron pairs

51. > between 1.7 and

52. Polar Covalent Bonds Account for the bonding found in HF
One electron from each. atom is shared but not equally due to unequal attraction for shared electrons
The bond is referred to as polar because 2 poles are formed
(+ and -)
Electronegativity values allows us to determine which atom has a greater pull
The atom with the greater electronegativity becomes the negative end of the polar bond.
The atom with the lower electronegativity becomes the positive end of the polar bond

53. Electronegativity
Electronegativity is the tendency of an atom to draw or attract the electrons in a bond toward itself
Electronegativity is like a game of tug-of-war, atom's ability to pull determines what kind of bond it forms
To form a covalent bond, two or more atoms with similar electronegativities will share electrons
Values fall between a low of 0.7 for Fr and a high 4.0 for F
The greater the difference in electronegativity the more polar the bond.

55. Double Covalent Bonds
Compounds sometimes share two pairs of electrons and form a double bond.
Examples: O2, CO2

56. Triple Covalent Bond Same idea as single and double
Two atoms of the same element or two different elements share three pairs of electrons and form a triple bond
Example: N2

57. Methods to Classify Bond Type
Non polar covalent bonds=no difference in electronegativity
Polar covalent bonds=difference less than 1.7
Ionic Bonds= difference of 1.7 or more. _ +

58. Using Electronegativity
Use the electronegativity difference to classify these compounds as covalent or ionic:
NH3
NaCl O2
BaCl2
H2O
BaO
Polar COVALENT
IONIC
Nonpolar COVALENT
IONIC
Polar COVALENT
IONIC

59. Using Electronegativity 2
Use the electronegativity difference to classify these compounds as covalent or ionic: N2
KBr
PbO CO H2
HCl
Nonpolar COVALENT
IONIC
IONIC
Polar COVALENT
Nonpolar COVALENT
Polar Covalent

60. Using Electronegativity 3
Use the electronegativity difference to classify these compounds as covalent or ionic:
MgS
FeO
CH4
LiF F2
Ca I2
IONIC
IONIC
Polar COVALENT
IONIC
Nonpolar COVALENT
IONIC

61. Metallic Bonding
Valence electrons in metals are free to move throughout the entire substance. Metallic bonds are positive nuclei surrounded by a sea of mobile electrons.
The valence electrons of metals are said to be delocalized.
Delocalized electrons cause metals to conduct electricity very well. Delocalized electrons also explain metallic luster, malleability, and ductility.

62. SEA OF MOBILE VALENCE ELECTRONS
Metallic Bonding
SEA OF MOBILE VALENCE ELECTRONS

63. Comparison of Two Types of Bonding
Covalent Bonds
A sharing of two or more electrons between atoms.
When two nonmetal atoms combine they do so by the formation of covalent bonds.
Ionic Bonds
An attraction that occurs between ions of opposite charge.
Metals and nonmetals are held together with ionic bonds.
Polyatomic ions also form ionic bonds. (e.g. ammonium chloride)

64. Metals
What specific property of metals accounts for their unusual electrical conductivity?

65. MC
An unknown substance is an excellent electrical conductor in the solid state and is malleable. What type of bonding does this substance exhibit? a. Ionic b. Covalent c. Metallic d. Cannot be determined

66. An unknown substance conducts electricity when dissolved in water
An unknown substance conducts electricity when dissolved in water. It has a melting point of 1200 oC. What type of bonding does this substance exhibit? a. Ionic b. Covalent c. Metallic d. Cannot be determined

67. MC
A chemical bond results from themutual attraction of nuclei for a. Electrons b. Neutrons c. Protons d. Positrons

68. MC
A polar covalent bond is likely to form between two atoms that a. Are similar in electronegativity b. Are of similar size c. Differ in electronegativity d. Have the same number of electrons

69. MC
A bond formed between two elements that have a very large difference in electronegativity is called a. Covalent b. Polar covalent c. Ionic d. Double bond

70. MC
To attain a noble gas configuration, an atom of chlorine must a. Gain 2 electrons b. Lose 2 electrons c. Gain 1 electron d. Lose 1 electron

71. Energy and Chemical Bonds
Potential energy is stored in chemical bonds.
Energy is released when bonds are formed.
Energy is absorbed to break a chemical bond.

72. Part 2 Forces Between Molecules

73. Opposites attract H-Cl H is the positive end of the molecule
Cl is more negative (has more electrons)

74. Same charges repel
H2 molecules repel eachother, it explains why they are gases.
Hydrogen is a nonpolar molecule…both ends are the same, positive.

75. Same charges still repel
All the nitrogens are surrounded by a lot of electrons…each is slightly negative.

76. Hydrogen Bonding Hydrogen bonding explains many of waters properties.
High boiling point
Explains why ice floats
Explains why salts dissolve

77. Hydrogen bonds are strong forces of attraction.

78. Ionic Solids Dissociate in Water

79. The case of the Gecko Walk!
Spatulae
Check the Gecko feet’s microstructure
Seta
1 billion of them, all in the van der Waals domain
Satae
Feet microstructure enables the Gecko reach van der Waals domains of attraction, they scale walls with ease!

80. Can we possibly make the Spiderman costume?
Spatulae
Satae
Seta
Can we possibly make the Spiderman costume?

81. + - O H + - Property Ionic Covalent (simple) Metallic Bonding Type
Which types of atoms does it involve?
How are they structured?
An example of the bonding type is
What is the melting point or boiling point like?
Are they magnetic?
Are they soluble in water?
Do they conduct electricity when solid liquid, gas or dissolved?
Give an example substance with a use
Bonding Type
Metal and Non-metal atoms
Non-metal atoms only
Metal atoms only
Form giant lattices that are brittle so are easy to crush
Form molecules with weak forces of attraction so are mainly gases and liquids
Forms a strong lattice structure + - O H Na + Cl -
High melting point and boiling point
Low melting point and boiling point
High melting point and boiling point
Not magnetic
Not magnetic
Some are magnetic
Many are soluble in water
Are not soluble in water
Are not soluble in water
Will conduct electricity when molten or dissolved in water
Will conduct electricity when solid o molten
Will not conduct electricity
Copper
Used for wiring electrical appliances
Sodium chloride
Used to flavour food
Water
Needed for life

82. Bonding Type Property Ionic Covalent (simple) Metallic
Which types of atoms does it involve?
How are they structured?
An example of the bonding type is
What is the melting point or boiling point like?
Are they magnetic?
Are they soluble in water?
Do they conduct electricity when solid liquid, gas or dissolved?
Give an example substance with a use