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Covalent Bonding

Covalent Bonding. Chapter 9. The Covalent Bond Section 9.1. Why do atoms bond? To achieve full outer electron shells Octet Rule- atoms gain, lose, or share electrons to achieve the electron configuration of noble gases Gain and Lose  IONIC BONDING Share  COVALENT BONDING.

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Covalent Bonding

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  1. Covalent Bonding Chapter 9

  2. The Covalent BondSection 9.1 • Why do atoms bond? • To achieve full outer electron shells • Octet Rule- atoms gain, lose, or share electrons to achieve the electron configuration of noble gases • Gain and Lose IONIC BONDING • Share  COVALENT BONDING

  3. What is a covalent bond? • A chemical bond that results in the SHARING of valence electrons • Occurs between 2 or more nonmetals

  4. A molecule is formed when two or more atoms bond covalently • Examples: sugars, DNA, proteins, fats, carbohydrates, cotton, synthetic fibers

  5. Formation of a covalent bond • REMEMBER: Hydrogen (H2), Nitrogen (N2), Oxygen (O2), Fluorine (F2),Chlorine (Cl2) Bromine (Br2) and Iodine (I2) occur in nature as diatomic molecules

  6. Covalent Bonding • An attractive force occurs between the protons of one atom and the electrons of the other atom • When a single pair of electrons is shared, such as in the hydrogen molecule, a single covalent bond forms

  7. Lewis Structures • Use electron-dot diagrams to show how electrons are arranged in molecules

  8. Group 7A (Halogens) have 7 VE • One more VE is necessary • A single covalent bond will form • Group 6A have 6 VE • Two more VE are necessary • Two covalent bonds will form • Group 5A have 5 VE • Three more VE are necessary • Three covalent bonds will form • Group 4A have 4 VE • Four more VE are necessary • Four covalent bonds will form

  9. Sigma Bonds- single covalent bonds- when electron pairs are centered between two atoms • Multiple Bonds • In many molecules, atoms attain noble gas configuration by sharing more than one pair of electrons between two atoms • Carbon, Nitrogen, Oxygen, and Sulfur most often form multiple bonds

  10. Strength of Covalent Bonds • The strength of covalent bonds depends on how much distance separates both nuclei • The distance between the two bonding nuclei at the position of maximum attraction is called bond length • Determined by the size of the atoms and how many electron pairs are shared • Bond length decreases as the number of bonds increases (triple bond has a shorter bond length than a single bond)

  11. Energy Changes • An energy change accompanies the forming or breaking of a bond between atoms in a molecule. • Energy is released when a bond forms • Energy must be added to break the bonds of a molecule • The amount of energy required to break a specific covalent bond is called bond dissociation energy

  12. Bond dissociation energy indicates the strength of a chemical bond because a direct relationship exists between bond energy and bond length • In chemical reactions, bonds in reactant molecules are broken and new bonds are formed as product molecules form

  13. Endothermic reactions occur when a greater amount of energy is required to break the existing bonds in the reactants than is released when the new bonds from in the product molecules • Exothermicreactions occur when more energy is released forming new bonds than is required to break bonds in the initial reactants

  14. Checkpoint • What is a covalent bond? How does it differ from an ionic bond? • What type of elements form covalent bonds? • Draw the Lewis Structures for each of these molecules: • PH3 • H2S • CCl4

  15. Naming Molecules Section 9.2 • Naming Binary Molecular Compounds • The first element in the formula is always named first, using the entire element name. • The second element in the formula is named using the root of the element and adding the suffix – ide. • Prefixes are used to indicate the number of atoms of each type that are present in the compound.

  16. Practice Naming • CCl4 • As2O3 • CO • SO2 • NF3

  17. Naming Acids • Binary Acids • Use the prefix hydro- to name the hydrogen part of the compound • The rest of the name consists of a form of the root of the second element plus the suffix –ic, followed by the word acid. • Examples: • HCl  Hydrochloric Acid • HCN Hydrocyanic acid (even though there are more than 2 elements present, if no oxygen is present- the acid is named as a binary)

  18. Oxyacids • Acids that contain an oxyanion (polyatomic ion that contains oxygen) • The name of the oxyacids consists of a form of the root of the anion, a suffix, and the word acid • If the anion suffix is –ate, it is replaced with the suffix –ic • If the anion suffix is –ite, it is replaced with the suffix –ous.

  19. Practice Problems • HI • HClO3 • HClO2 • H2SO4 • H2S

  20. Checkpoint • Write the molecular formula for each of the following compounds • Disulfur trioxide • Iodic acid • Dinitrogen monoxide • Hydrofluoric acid • Phosphorus pentachloride • What is the difference between a binary acid and oxyacid? • Complete the following table 

  21. Molecular Structures Section 9.3 • Structural Formula- uses letter symbols and bonds to show relative positions of atoms • Lewis Structure Procedure 1. Predict the location of certain atoms • Hydrogen is always a terminal, or end, atom. Because it can share only one pair of electrons, hydrogen can be connected to only one other atom • The atom with the least attraction of shared electrons in the molecule is the central atom. This element usually is the one closer to the left on the periodic table. The central atom is located in the center of the molecule, and all other atoms become terminal atoms.

  22. 2. Find the total number of electrons available for bonding. This total is the number of valence electrons in the atoms in the molecule. 3. Determine the number of bonding pairs by dividing the number of electrons available for bonding by two. 4. Place one bonding pair (single bond) between the central atom and each of the terminal atoms.

  23. 5. Subtract the number of pairs you used in step 4 from the number of bonding pairs you determined in step 3. The remaining electron pairs include lone pairs as well as pairs used in double and triple bonds. Place lone pairs around each terminal atom bonded to the central atom to satisfy the octet rule. Any remaining pairs are assigned to the central atom.

  24. 6. If the central atom is not surrounded by 4 electron pairs, it does not have an octet. You must convert one or two of the lone pairs on the terminal atoms to a double bond or a triple bond between the terminal atom and the central atom. These pairs are still associated with the terminal atom as well as with the central atom. Remember that, in general, carbon, nitrogen, oxygen, and sulfur can form double or triple bonds with the same element or with another element.

  25. Space-filling Structure Lewis Structure Ball-and-stick molecular model

  26. NH3 PO4-3 CO2 BH3 Examples

  27. Resonance Structures • Resonance is a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion • Nitrate Ion Resonance Structures:

  28. Exceptions to the Octet Rule • A small group of molecules has an odd number of valence electrons and cannot form an octet around each atom. (i.e. NO2 , ClO2, and NO) • Some compounds form with fewer than eight electrons present around an atom. (Example: Boron) • When one atom donates a pair of electrons to be shared with an atom or ion that needs two electrons to become stable, a coordinate covalent bond forms

  29. 3. Some central atoms contain more than eight valence electrons (expanded octet) • Examples: PCl5, SF6, and XeF4

  30. Molecular ShapeSection 9.4 • The shape of the molecule determines many of its physical and chemical properties. • Molecular shape is determined by the overlap of orbitals that share electrons • Valence Shell Electron Pair Repulsion model or VSEPR model • Based on an arrangement that minimizes the repulsion of shared and unshared pairs of electrons around the central atom

  31. VSEPR Model • The angle formed by any two terminal atoms and the central atom is a bond angle. • Shared electron pairs repel one another • Lone pairs of electrons occupy a slightly larger orbital than shared electrons • Shared bonding orbitals are pushed slightly together by lone pairs

  32. Look at the VSEPR cheat sheet 

  33. Hybridization • A hybrid results from combining two of the same type of object, and it has characteristics of both. • Hybridization- a process in which atomic orbitals are mixed to form new, identical hybrid orbitals.

  34. Practice Problems • Determine the molecular geometry and bond angle for the following: • BF3 • NH4+ • OCl2 • CF4

  35. Electronegativity and PolaritySection 9.5 • Electron affinity is a measure of the tendency of an atom to accept an electron

  36. The character and type of a chemical bond can be predicted using the electronegativity difference of the elements that are bonded • Polar Covalent- Unequal sharing • Nonpolar Covalent- Equal sharing Generally- Ionic bond form when the electronegativity difference is greater than 1.70

  37. Polar Covalent Bonds • Polar covalent bonds form because not all atoms that share electrons attract them equally • The shared pair of electrons is pulled toward one of the atoms • Partial charges occur at the ends of the bond • Partially negative • Partially positive  • The resulting polar bond is referred to as a dipole (two poles)

  38. Molecular Polarity • Molecules are either polar or nonpolar, depending on the location and nature of the covalent bonds they contain. • A polar molecule has a partial negative charge on one side, while the other side of the molecule has a partial positive charge • Note:symmetric molecules are usual nonpolar and molecules that are asymmetric are usually polar

  39. Polar Molecule or not? • Compare H2O and CCl4

  40. Solubility of polar molecules • The ability of a substance to dissolve in another substance is known as the physical property solubility • The bond type and the shape of the molecules present determine solubility • “Likes dissolve likes” • Polar compounds are usually soluble in polar substances • Nonpolar molecules only dissolve in nonpolar substances

  41. Properties of Covalent Compounds • Lower melting and boiling points (indicating weak bond strength) • Many are liquids or gases at room temperature • Do not conduct electricity • Many do not dissolve in water (polar)

  42. Practice Problems • Decide whether each of the following molecules is polar or nonpolar • SCl2 • H2S • CF4 • CS2

  43. Structural formula molecule VSEPR model Coordinate covalent bond hybridization oxyacid electronegativity polar covalent covalent bond resonance endothermic exothermic terminal atom Sigma bond Vocabulary on Test

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