A chemical bond is a bond by which the atoms in the molecule of a compound are combined and held together by a strong combining force. There are two types of chemical bonds; ionic bond and covalent bond.
1. Chemical Bonding
2. bond: forces that hold one atom to another in a compound
3. To break a bond requires energy to be put in to overcome the forces of attraction Bond breaking is endothermic.
4. Tomake bonds causes a release of energy. Bond making is exothermic.
5. Compounds have less energy (more stable) than the substances from which they form. Ex:water has less energy than the hydrogen and oxygen from which it formed. The energy stored in a bond is potential.
6. Three Types of Bonds 1. Metallic 2. Ionic 3. Covalent
7. Metallic Bonds Definition: bonds between atom in a metal; ions held together in a crystalline lattice in a “sea of mobile electrons”
8. Metallic Bonds Conduct electricity because of freely moving electrons. Any substance that has moving charged particles (typically either mobile electrons or ions) will conduct electricity. High MP and high BP because bonds are strong. malleable
9.
10. Ionic Bonding: What does the word “ionic” mean? Comes from the word ion meaning charged particle. Ions are formed when atoms gain or lose e-.
11. Electrical Attraction Opposites attract.
12. Ionic Bonds Formedwhen electrons are transferred between atoms.
13. Ionic Bonds: Metals and Non-metals Metals… lose e-. Non-metals… gain e-. They can exchange electrons to form a bond.
14. The Octet Rule Atoms are happy with a full valence shell. This achieved by gaining or losing electrons.
15. Metalstend to have low ionization energies, so they lose e- easily. Non-metals tend to have high electronegativities, so they gain e’ readily.
16. To add to notes… The two Elements involved in an ionic bond have a difference in electronegativity (E.N.D.) that is greater than or equally to 1.7 Ex: NaCl E.N.D.= 2.3 (How do you find this? Look up electronegativity of each element on Table S and subtract.) Na = 0.9 Cl = 3.2 3.2 – 0.9 = 2.3
17. sodium reacts with chlorine Na Cl Has one valence Has 7 valence electron, wants electrons, needs to get rid of it so one more to that its valence have a full shell is full valence shell This reactions makes sodium chloride or table salt. Na+1[ Cl ]-1
18. See the reaction Reaction to form an ionic compound In your notes, describe the reaction. Are the reactive properties of these elements consistent with what we learned last chapter about the groups of the periodic table?
19. Summary of Concepts: Ionic Bonds Occur between metals (+) and non- metals (-) Involve a transfer of electrons from the metal to the nonmetal. The two elements have an electronegative difference of 1.7 or greater We call ionic compounds “salts”
20. Covalent Bonds vs. Ionic Bonds
21. Two Hydrogen Atoms The valence shells overlap and the electrons are shared making a more stable molecule.
22. Covalent Bonds Involves SHARING electrons because both elements have high electronegativities. Sharing of electrons can be equal (non-polar) or unequal (polar) Usually is between two NON-METALS (ex. H, C, O, N…) Covalent bonding like Ionic bonding results in a more stable compound, because the atoms involved meet the “octet rule”.
23. Covalent Bonding (cont.) A “shared” pair of electrons makes a SINGLE BOND (2 total e-). 2 “shared” pairs makes a DOUBLE BOND (4 total e-). 3 “shared” pairs makes a TRIPLE BOND (6 total e-). We call them MOLECULES. Covalent bonds are very strong bonds.
24. Ionic Bonding Involves a transfer of electrons because atoms have an E.N.D. greater than or equal to 1.7. These bonds are always polar because it involves a + and – ion. Usually involves a metal(+) and a nonmetal(-). (ex. NaCl) The oppositely charged ions create an electrical attraction which forms the bond. Ionic bonds result in compounds (not molecules) that are more stable because the atoms meet the octet rule.
25. Simple Molecules (covalent bonds) Diatomic molecules are atoms of the same element that covalently bond to meet the octet rule. The following elements are diatomic: H2, O2, F2 Br2, I2, N2, Cl2 Alittle trick to remember them: HOF BrINCl
26. Lewis Dot Diagram of Atoms and Ions (Review) Sodium atom Potassium ion Magnesium atom Calcium ion Chlorine atom Iodide ion Aluminum atom Oxide ion Sulfur atom Aluminum ion
27. Dot Diagrams for Ionic 1. NaI 2. Na2S 3. RbBr 4. CaF2 5. AlCl3 6. BaO 7. Li3N 8. K3P 9. MgO 10. BaCl2
28. Drawing Lewis Dot Diagrams for Covalently Bonded Molecules 1) Add up all the valence electrons for an atom in the molecule. 2) The “central atom” is the one that is most electronegative or there will only be one of this atom. 3) Add the electrons until you reach the total remembering that all the atoms must obey the octet rule except for hydrogen which obeys the duet rule. 4) Also remember that one pair of electrons between two atoms is a single bond two is a double and three pairs is a triple bond. These are known as bonded or shared pairs of electrons. 5) Not every electron has to be bonded. We call these the lone or non-bonded pairs because they are “lonely”.
29. H2 F2 Cl2 H2O NH3 CH4
30. Practice on your own …
31. Covalently Bonded-Molecular
32. Coordinate Covalent Bonds +1 +1 When a shared pair of electrons comes from only one atom, not two its known as a coordinate covalent bond. Polyatomic Ions have these.
33. Polyatomic Ions dot diagrams
34. Metal and a Polyatomic Ion Dot
35. Bond Polarity Polar Covalent Bonds (between two atoms) Covalent bonds unlike ionic bonds, are NOT composed of oppositely charged ions. However, we find that since the electrons are shared and each element has a different electronegativity, the shared electrons are often shared unequally. Meaning the electrons spend more time around the more electronegative atom. This causes one atom to be slightly negative and one atom to be slightly positive. This is known as a dipole moment. One positive end and one negative end 2 oppositely charged poles.
36. Bond Polarity Non-polar Covalent Bond (between two atoms) If both of the atoms involved in a bond have the same electronegativity then the electrons are shared equally. This will happen if a bond forms between two of the same atoms. Since the atoms are the same they will have the same electronegativity and share the electron(s) equally. When electrons are shared equally their will never be a dipole moment and the bond will be non-polar.
37. Molecule Polarity Polar Molecule A molecule that is polar (also known as a dipole) is asymmetrical which results in an uneven distribution of charge throughout the entire molecule. Ex: H2O which has a bent shape and is assymetrical
38. Molecule Polarity Non-Polar Molecules A molecule that is non-polar is symmetrical in shape which results in an even distribution of charge. Ex CH4 which is tetrahedral symmetrical in shape
39. Bond Polarity vs. Molecule Bond Polarity is based on the electronegativity between the bonded atoms. If they have different electronegativities (different atoms) the bond must be polar. If the have the same electronegativity (same atoms) the bond must be non-polar Molecule polarity is based on the shape symmetry of the entire molecule. If the molecule has a shape that is symmetrical then the charge distribution is even and the molecule is non-polar. If the molecule has a shape that is asymmetrical then the charge distribution is uneven and the molecule is polar.
40. Shapes of Molecules The shape of a molecule will help determine whether or not the molecule is polar or nonpolar Examples of Shapes of Molecules CO2
41. Shapes Continued CH4
42. Shapes Continued HCl
43. Shapes Continued NH3
44. Shapes Continued H2O
45. Summary of Shapes If the central atom has ... the shape is example 1 or 2 bonds only linear 2 bonded pairs and 2 lone bent 3 bonded pairs and 1 lone pyramidal 4 bonded pairs and 0 lone tetrahedral
46. Intramolecular and Intermolecular Forces Atoms involved in a bond are held together by what is known as an Intramolecular Force. An intramolecular force is just simply a bond so it can be either ionic, polar covalent or nonpolar covalent. When substances are in a solid or liquid state the compounds or molecules are held together by a force as well. The force that holds molecules together in a solid or a liquid sample is known as an Intermolecular Force.
47. Ionic Compounds (ionic bonds) For solid or liquid Ionic compounds the only intermolecular force that is possible is created by the attraction between the oppositely charged ions. This is known as a dipole-dipole attraction between ions. For example in a Salt Crystal (NaCl) the positive and negative ions will alternate as shown in the diagram below. It is the opposite charges of the ions that hold the crystal together.
48. Molecular Compounds (polar covalent or nonpolar covalent) For solid or liquid molecular the compounds there are three main intermolecular forces that hold the molecules together. Some of the intermolecular forces for molecules are strong, because like the ionic solids it is due to a dipole-dipole attraction. Some of the intermolecular forces are weak, because they are only based on the size of the molecules. Intermolecular forces for molecules that involve a dipole-dipole attraction only exist in polar molecules.
49. Polar Molecules Hydrogen bonding is one type of intermolecular force it is formed between a hydrogen atom in one molecule and a nitrogen, oxygen, or fluorine atom in another molecule. Hydrogen bonding is the strongest intermolecular force for all molecular substances. Substances with H-bonding have relatively high boiling points, because the force is so strong. The most important example of this is found in water. The strength of Hydrogen bonding is the reason why water has a relatively high boiling point compared to other molecular compounds.
50. Polar Molecules (Continued) Ifwater did not have H-bonding it would not exist on earth as a liquid only a vapor. The diagram below shows how the H-bonding occurs. H-bonding also occurs between ammonia molecules and Hydrogen Fluoride molecules.
51. Polar Molecules (Continued) Allother intermolecular forces for polar molecules are due to the dipole-dipole attractions between the molecules, and the force is not quite as strong as Hydrogen bonding. Remember H-Bonding only occurs between the HYDROGEN atom of one molecule and either an OXYGEN, NITROGEN or FLUORINE atom of another molecule.
52. Nonpolar Molecules For nonpolar molecules there is no dipole so the force of attraction between the molecules is based on the molecular mass of the molecules. The main rule for non-polar molecules is the smaller the molecules the weaker the intermolecular forces, and the larger the molecules the stronger the intermolecular forces. Smaller non-polar molecules tend to be gases at room temperature, because the weak intermolecular forces give them low boiling points. Examples: methane (CH4) and Hydrogen (H2)are gases at room temp. Larger non-polar molecules will have stronger intermolecular forces so they will usually be liquids or solids at room temperature. Examples: Glucose (C6H12O6) and Gasoline (C8H18)
53. 1 Intramolecular forces are chemical bonds between the atoms that make up an individual molecule. 2. Intermolecular forces are attractions between two or more molecules. 3. H-bonding is one type of intermolecular force that occurs between the hydrogen of one molecule and the oxygen, nitrogen or fluorine of another molecule. 4. H-bonding is the strongest intermolecular force for polar molecules, which is why these substances will have high boiling points. 5. All other polar molecules have a dipole-dipole attraction.
54. Summary (continued) 6. Nonpolar molecules have intermolecular forces that depend on the mass of the molecules. 7. Molecules with a high molecular mass will have strong intermolecular force. 8. Molecules with a low molecular mass will have weak intermolecular force. 9. A substances boiling point and its intermolecular forces are directly related. 10. Substances with high boiling points have strong intermolecular forces 11. Substances with low boiling points have weak intermolecular forces.
55. Properties of Bonds Metallic, covalent and ionic bonds all have varying properties. These properties can be used to identify the bonding of unknown substances. The three main properties that can be used to distinguish bond type are Melting & Boiling Points, Hardness in the solid state, and Conductivity in the solid, liquid and aqueous states.
56. Properties of Metallic, Ionic and Covalent Bonds Summary Bond Melting & Boiling Hardness Conductivity Type Points Solid Liquid Aqueous Metallic High (except Hg) Hard Yes Yes Yes Covalent Low Soft No No No Ionic High Hard No Yes Yes
57. Properties of Metallic, Ionic and Covalent Bonds Summary Conductivity Bond MP Hardness Type & Solid Liquid Aqueous BP State State Metallic High (except Hard Yes Yes Yes Hg) Ionic High Hard No Yes Yes Covalent Low Soft No No No
58. Why do some substance have conductivity and others do not? In order for a substance to conduct electricity and heat the substance must have charged particles that are free to move or mobile. Metallic substances can always conduct electricity and heat because the bonds are created by free flowing or mobile valence electrons between the metal atoms. These mobile electrons are what move the electrical current through the substance. Ionic substance are made up of positive and negative ions (charged particles) but they are not free to move in the solid phase so they can not conduct as soilds. When Ionic substances are melted or dissolved in water the ions are free to move so they have the ability to conduct in the liquid and aqueous states, because of the mobile ions. Molecular substances are made up of atoms sharing electrons in a covalent bond and since they do not have any charged particles they can NEVER conduct electricity.
59. Shapes of Molecules CO2 – linear, space-filling model vs. ball and stick Polar bonds (C-O), Non-polar molecule b/c it is symmetrical
60. Shapes of Molecules CH4 – tetrahedral shape, ball and stick model Polar bonds (C-H), Non-polar molecule b/c it is symmetrical
61. Shapes of Molecules HCl – linear shape, ball and stick model Polar bonds- Polar Molecule, asymmetrical
62. Shapes of Molecules NH3 – trigonal pyramidal shape Polar bonds (N-H) and polar molecule b/c asymmetrical.
63. Shapes of Molecules H2O – bent shape Polarbonds (O-H) and polar molecule b/c asymmetrical.
64. Solid State Chemistry- Crystalline Solids
65. Ionic Solids Repeating particles in crystals are ions (+/-). Examples: NaCl, KNO3, CaCl2, etc. Diagram
66. Ionic Solids Properties: High melting point (strong bonds) Electrically conductive ONLY when dissolved in water or in liquid form (mobile charged particles) Water soluble (exceptions on Table F) Relatively hard
67. Molecular Solids: Repeating particles are atoms or molecules. Examples: CO2, H2O, S8 Diagram:
68. Molecular Solids Properties: Relativelylow MP and BP (based on Intermolecular forces) Soft Non-conducting Only polar molecules are soluble in water
69. Covalent-Network Solids: Repeating particles are atoms held together by covalent bonds NOT intermolecular forces. Examples: Diamond, quartz, SiO2, SiC Diagram
70. Covalent Network Solids Properties: Very strong covalent bonds Very high melting point Very hard Not soluble in water
71. Graphite is weird!
72. Allotrope: different molecular structure of the same element; also has different properties Examples: Carbon (diamond vs. graphite (C8) vs. buckminster fullerenes (C60), Oxygen vs. Ozone
73. Metallic Solids Repeating particles are metal atoms with mobile valence electrons moving throughout the crystal; a “sea of mobile electrons” Examples: Ag, Au, Na, Cu, Zn Diagram
74. Metallic Solids Properties: Some are hard, others are soft Conduct electricity and heat Malleable/ductile insoluble
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