1 / 32

In Summary……

In Summary……. Covalent bond is the sharing of two electrons between the 2 atoms. Covalent bonding . Two hydrogen atoms can share their valence electrons to attain the same electron configuration of the nearest Noble gas configuration, Helium. Covalent Bond

leif
Télécharger la présentation

In Summary……

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. In Summary……

  2. Covalent bond is the sharing of two electrons between the 2 atoms Covalent bonding Two hydrogen atoms can share their valence electrons to attain the same electron configuration of the nearest Noble gas configuration, Helium

  3. Covalent Bond When nonmetallic elements react with other nonmetallic elements, they share electrons in order to obtain eight valence electrons. Each fluorine atom has seven valence electrons. They each require one more electron to satisfy the Octet Rule.

  4. The left fluorine atom now has a total of eight electrons and the right fluorine atom now has a total of eight electrons around it. When nonmetallic elements react with other nonmetallic elements, they share electrons in order to obtain eight valence electrons.

  5. The two electrons that form the covalent bond are often Represented by a single line. The F2 molecule can be represented using a line and dots to show the bonding pair and the six lone pairs, respectively. This is called a Lewis dot structure.

  6. Multiple Covalent Bond Some atoms have to share more than one electron in order to satisfy the Octet Rule.

  7. Each oxygen atom has six valence electrons. They each require two more electrons to satisfy the Octet Rule.

  8. The left oxygen atom now has a total of eight electrons around it. The right oxygen atom now has a total of eight electrons around it.

  9. The four electrons shared by the oxygen atoms form a double bond. The double bond is represented by two single lines. Each line in the Lewis dot structure represents two electrons

  10. The element hydrogen is an exception to the Octet Rule. It only needs two electrons, rather than eight, to be stable. The hydrogen atom has one valence electron. It requires one more electron to be stable. The fluorine atom has seven valence electrons. It requires one more to satisfy the Octet rule.

  11. The hydrogen atom now has a total of two electrons around it and is stable. The fluorine atom now has a total of eight electrons around it and is stable.

  12. The Lewis dot structure of the HF molecule shows a line and 6 dots to represent the bonding pair and the 3 lone pairs of electrons, respectively.

  13. Rules for writing Lewis Dot structures • Rule 1 Add together the number of valence electrons for each atom in the molecule. For example, CF4 Carbon has four valence electrons and each fluorine has seven valence electrons = 4 + 4(7) = 32

  14. Rule 2 Write out the elements of the molecule so that the least electronegative elements is in the center surrounded by the other elements. For example, CF4

  15. Rule 3 Place a covalent bond between the central atom and the outside atoms. Remember each covalent bond contains two electrons.

  16. The four covalent bonds use eight of the 32 valence electrons in CF4 • This uses 24 electrons. There Are no electrons left, so this is The Lewis dot structure for CF4 • Rule 4 There are 24 valence electrons remaining. Add electrons to the outer atoms as lose pairs to satisfy the Octet Rule.

  17. Rule 5 for example, NH3 • First apply Rules 1-4 to the molecule • Rule 1: Count the valence electrons • Rule 2: Place the least electronegative element at the centre, except for H which is always an outer atom • Rule 3: Add covalent bonds between the centre atom and the outer atoms • Rule 4: Add lone pairs to the outer atoms • Rule 5: Add lone pairs to the centre atom

  18. Rule 1 Nitrogen has 5 valence electrons and each hydrogen has 1 valence electron The total number of valence electrons = 5 + 3 (1) = 8 Rule 2 Hydrogen is always an outer atom and is never at the centre of a molecule

  19. Rule 3 Add the bonding electrons. This uses 6 of the 8 valence electrons. Rule 4 The 2 remaining valence electrons are not added to the outer atoms, because each H has its maximum of 2 valence electrons.

  20. Rule 3 Add the bonding electrons. This uses 6 of the 8 valence electrons. • Rule 4 The 2 remaining valence electrons are not added to the outer atoms, because each H has its maximum of 2 valence electrons.

  21. Rule 5 Placethe remaining 2 Valence electrons on the central nitrogen atom This is the Lewis structure For NH3 Rule 6 Check all atoms in the molecule to ensure that each has 8 electrons(2 for hydrogen). If an atom has fewer than 8 electrons, create double or triple bonds. (Note: Double bonds only exist between C,N,O and S atoms)

  22. Apply rule 6 to the following; CH4, CF4, Hydrogen : 1 bond = 2 electrons (stable) Carbon : 4 bonds = 8 electrons (stable) • Fluorine : 1 bond + 3 lone pairs = 2 + 3 (2) = 8 electrons (stable) • Carbon : 4 bonds = 8 electrons (stable)

  23. Example; CH2O Apply Rules 1-5 to the molecule Rule 1: Count the valency electrons Rule 2: Place the least electronegative element at the centre, except for H, which is always an outer atom Rule 3: Add covalent bonds between the centre and the outer atoms Rule 4: Add lone pairs to the outer atoms Rule 5: Add lone pairs to the centre atom

  24. Rule 1 Carbon has 4 valence electrons, each hydrogen has 1 valence electron, and oxygen has 6 valence electrons. Total number of valence electrons : 4 + 2(1) + 6 = 12 Rule 2 Carbon is at the centre of the molecule because it is less electronegative than oxygen. Hydrogen is always an outer atom and is never at the centre of the molecule.

  25. Rule 3 Add the bonding electrons. This uses 6 of the 12 valence electrons • Rule 4 Add the remaining 6 lectrons to the outer atom. Hydrogen does not need any more electrons, but Oxygen needs 6 to complete its octet.

  26. Rule 5 There are no valence electrons left to add to the centre • Rule 6 Oxygen shares one of its lone pairs with C and O and give the desired 8 electron total This is the Lewis dot Structure for CH2O

  27. Exceptions to the Octet Rule The Octet Rule applies to Groups IVA through VIIA in the second row of the Periodic Table, but there are exceptions to the rule among some other elements. The following two cases are an example Example BF3 Rule 1 Boron has 3 valence electrons and each Fluorine has 7 valence electrons Total number of electrons = 3 + 3 (7) = 24

  28. Rule 2 Boron is at the centre of the molecule because it is less electronegative than fluorine Rule 3 Add the bonding electrons. This uses 6 of the 24 valence electrons

  29. Rule 4 Add the remaining electrons to the outer atoms. Each Fluorine has the required 8 electrons Rule 5 This uses the remaining electrons leaving none to add to the Boron central atom

  30. Rule 6 Check the number of electrons around each atom. Each Fluorine atom has 8 electrons, but the Boron Atom has only 6. This is an exception to the Octet Rule. A B=F bond is not an option, because double bonds exist only between C,N,O, and S atoms This is the Lewis dot structure BF3

  31. Metallic Bonding The valence (outermost) electrons are loosely held by the metal ions, so much so that they move away from the atom to form a positively charged ION. The electrons are free to move from one positively charged ION to the next (i.e. They are DELOCALISED) and are shared (just like in covalent bonding among the various metallic positively charged ions However, metals behave differently. Metallic bonding is similar to both covalent and ionic bonding The number of electrons = the number of protons. The metal is therefore electrically NEUTRAL Source: www.daviddarling.info/images/metallic_bond.jpg

  32. Compare and contrast types of bonding Similarities Differences • Metallic and ionic bonding involve electrostatic attractions between positive and negatively charged particles. • Metallic bonding shares electrons among the ions in a similar manner to how electrons are shared in covalent bonding. • Covalent bonding shares electrons rather than having electrostatic charges. • Ionic bonding will form compounds whereas covalent bonding can form a compound or element and metallic bonding is strictly found in elements

More Related