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Lewis Dot Structures of Covalent Compounds

Lewis Dot Structures of Covalent Compounds. Atoms are made up of protons, neutrons, and electrons. The protons and neutrons are located at the center of the atom, the nucleus. These electrons can be divided into core electrons and valence electrons. The valence electrons are

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Lewis Dot Structures of Covalent Compounds

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  1. Lewis Dot Structures of Covalent Compounds Atoms are made up of protons, neutrons, and electrons. The protons and neutrons are located at the center of the atom, the nucleus. These electrons can be divided into core electrons and valence electrons. The valence electrons are the outermost electrons and are the ones involved in chemical reactions

  2. p+ p+ p+ p+ p+ p+ p+ e- e- e- e- e- e- e- e- e- e- Electrons Electrons occupy most of the volume of an atom They arrange themselves in ’shells’ at varying distances from the nucleus The Nucleus Protons and neutrons are located in the nucleus (center) of the atom n0 n0 n0 n0 n0 Valence electrons These are the outermost electrons and the ones In chemical reactions

  3. The number of valence electrons varies by element. For the Main Group elements, the number of valence electrons is equal to the Group Number that the elements belong to. For example, Sodium (Na) belongs to Group 1A and therefore has 1 valence electron.

  4. The Periodic Table

  5. For example, Bromine (Br) belongs to Group VIIA and therefore has 7 valence electrons. We can represent the valence electrons of an atom using a Lewis dot symbol, in which the element symbol is surrounded by dots representing the valence electrons. For example, Oxygen has six valence electrons, so its Lewis dot symbol is: Note the six dots representing the six valence electrons

  6. For example, neon has eight valence electrons, so its Lewis dot symbol is: For example, carbon has four valence electrons, so its Lewis dot symbol is :

  7. How many valence electrons does Potassium (K) have? 1 How many valence electrons does Antimony (Sb) have? 5 How many valence electrons does Phosphorus (P) have? 5 How many valence electrons does Magnesium (Mg) have? 2

  8. The Noble Gas elements in Group VIIIA have either two valence electrons (He) or eight valence electrons (Ne, Ar, Kr, Xe, and Rn). These elements are extremely stable because they have full valence shells- two electrons for He in the first row and eight electrons in each of the later rows. This is the basis for the Octet Rule - elements tend to react in a way to attain the electron configuration of Group VIIIA

  9. Metallic elements at the left side of the Periodic Table tend to lose one or more electrons to form positive ions, such as Na+ and Mg2+, each of which has the electron configurationof the Noble Gas that preceds it. Nonmetals at the right side of the Periodic Table tend to either gain electrons to form negative ions such as F-, O2-, and N3- or to share electrons in covalent bonds. This learning objective describes how this is done

  10. 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.

  11. 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.

  12. 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.

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

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

  15. 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.

  16. 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

  17. 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.

  18. 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.

  19. 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.

  20. 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

  21. 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

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

  23. 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.

  24. 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

  25. 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

  26. 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.

  27. 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.

  28. 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)

  29. 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)

  30. 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

  31. 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.

  32. 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.

  33. 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

  34. 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

  35. 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

  36. 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

  37. 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

  38. Example PF5 Rule 1 Phosphorus has 5 valence electrons and each fluorine has 7 valence electrons Total number of electrons • = 5 + 5(7) = 40 Rule 2 Phosporus is at the centre because it is less electronegative than fluorine

  39. Rule 3 Add the bonding electrons. This uses 6 of the 24 valence electrons. Rule 4 Add the remaining electrons to the outer atoms. Each Fluorine requires 6 more electrons

  40. Rule 5 This uses the remaining electrons leaving none to the central P atom Rule 6 Check the number of electrons around each atom. Each Fluorine atom has 8 electrons, but the phoshorus atom has 10. This is an exception to the Octet Rule.

  41. Rule 6 Check the number of electrons around each atom. Each fluorine atom has 8 electrons, but the phoshorus atom has 10 . This is an exception to the Octet Rule.

  42. How Elements Form Compounds Some atoms lose or gain electrons to become stable charged particles called ions When atoms loses electrons, they form positively charged ions called cations When atoms gain electrons, they form negatively charged ions called anions.

  43. Sodium chloride is a relatively harmless compound because the sodium and chlorine atoms have stable ions . The compound formed is called an ionic compound because it is made up of positive and negative ions that have resulted from the transfer of from a metal to a nonmetal. The positive and negative ions are attracted to each other because they have opposite charges.

  44. When ionic compounds are placed in water, the ions separate and are surrounded by water molecules. They are electrolytes. They are also conductive

  45. Ionic Charges and Chemical Families • Review • structure of the atom • How some atoms can form stable ions by gaining or losing electrons • You have also learned that the PT is a useful organizing tool for predicting behaviour of substances

  46. The location of the alkalis metals (dark green), the alkaline earth metals (light green), and the halogens (red) in the PT Activity

  47. Ionic Compounds There are over 100 elements in the PT Thousands of different compounds are formed when these elements combine. How can we name these compounds? How can we write formulas to represent them?

  48. We have seen from past discussions that The PT and a knowledge of the electronic structure could be used to predict ionic charge of elements Ionic charges (or valences) of some elements in the PT

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