COVALENT COMPOUNDS. Back to Lewis Dot Structures and Valence Electrons!

Transcription

1 COVALENT COMPOUNDS Back to Lewis Dot Structures and Valence Electrons!

2 Review of Lewis Dot Structures Electron Dot Structures contain: Element s Symbol: representing the atom s nucleus and inner electrons Dots: representing all of the valence electrons n 8 maximum Examples: N O Cl

3 Covalent Bonding A joining of two atoms through the sharing of valence electrons. Most common with non-metals. Octet Rule will apply. They share electrons to become more stable.

4 Sharing Electrons Atoms want to become stable and look like the noble gas (in terms of their electron configuration) Because non-metals are close to having 8 electrons, they will not give up electrons! In order to look like the noble gas, they must share their valence electrons with each other. Both atoms get to count the shared electron(s) toward the noble gas configuration

5 Example: Fluorine One atom of Fluorine has 7 valence electrons. It needs 1 more electron to become stable. A second fluorine atom also has 7 valence electrons. Again, it needs only 1 more electron to become stable. By sharing electrons both atoms will end with full orbitals.

6 Types of Covalent Bonds Single Covalent Bond: Two atoms are held together by sharing one pair of electrons Usually represented with a single line These are the longest bonds and also the weakest Example: Cl-Cl

7 Types of Covalent Bonds Double Covalent Bond: Two atoms held together by sharing two pairs of electrons. Usually represented with a double line Example: O=O

8 Types of Covalent Bonds Triple Covalent Bond: Two atoms held together by sharing three pairs of electrons Usually represented with a triple line These are the shortest and strongest bond

9 Electrons Involved in Bonding Depending on the number of valence electrons available: Most atoms can have up to 4 covalent bonds. This is based on unshared electron pairs or Lone Pairs of electrons These are valence electrons that are NOT involved in a covalent bond.

10 Example: Oxygen Lone Pairs of Electrons

11 Rules for Drawing Lewis Dot Structures Rule 1: Add together the number of valence electrons for EACH atom in the molecule Example: CF 4 Carbon has 4 valence electrons Each Fluorine has 7 valence electrons Therefore, the total number of valence electrons = 4 + 4(7) = 32

12 Rules for Drawing Lewis Dot Structures Rule 2: Write out the elements of the molecule so that the least electronegative element is in the center surrounded by the other elements n TIP: Fluorine is the most electronegative. The closer the elements are to Fluorine, the higher the electronegativity! Example: CH 4 F C F F F

13 Rules for Drawing Lewis Dot Structures Rule 3: Place a covalent bond between the central atom and the outside atoms. n Covalent bonds are represented by a line. n Remember: each covalent bond contains two electrons. F Example: CF 4 F C F F

14 Rules for Drawing Lewis Dot Structures Rule 4: Add electrons to the outer atoms as lone pairs to satisfy the Octet Rule. Example: CF 4 We started out with 32 electrons Each line you drew connecting Carbon to Fluorine was 2 electrons. So in total you used 8 electrons. There are now 24 valence electrons remaining.

15 F F C F F

16 Practice Problem #1 Write the Lewis Dot Structure for TeH 2 What is the total number of valence electrons in this molecule? n Te: 6 valence electrons n H: 1 valence electron n Total: 6 + 2(1) = 8 Which atom is at the center of the molecule? n The least electronegative is Te

17 Practice Problem #1 Draw Te in the middle and H on opposite sides of Te. H Te H How many valence electrons do you have left? You have 4 valence electrons left. You started with 8 electrons and use four for the bonds n Two bonds each with two electrons

18 Practice Problem #1 Will the remaining 4 electrons go on Te or H? Te! H Te H You have now written the Lewis Dot structure for TeH 2

19 Practice Problem #2 Write the Lewis Dot Structure for CO 2 What is the total number of valence electrons in this molecule? n Carbon: 4 valence electrons n Oxygen: 6 valence electrons n Total: 4 + 2(6) = 16 valence electrons Which atom is at the center of the molecule? n The least electronegative element is Carbon

20 Practice Problem #2 Draw Carbon in the middle and O on opposite sides O C O How many valence electrons do you have left? You have 12 valence electrons left. You started with 16 electrons and use four for the bonds n Two bonds each with two electrons

21 Practice Problem #2 Will the remaining 16 electrons go on C or O? They will be split amongst both, C and O, in order to give each atom a complete octet! O C O I have now used up my remaining 12 electrons. Do they form a complete octet for each atom? NO.

22 Practice Problem #2 Each atom does not have a complete octet and we need to fix this problem. O C O To fix the problem, you need pair up any electrons that are not paired

23 Practice Problem #2 To pair them together, just draw another line between the C and O. This will form a double bond! O C O Now, does each atom contain a full octet? Yes!

24 COVALENT BONDS Naming Compounds, Writing Formulas

25 Properties of Covalent Bonds Bond Formation Type of Structure Physical State Melting Point Electrons are shared between atoms Molecules Can be liquids, gases, or brittle solids Low Soluble in Water? Most are not, BUT there are a few that are Electrical Conductivity Other No Some have odors

26 Strength of Covalent Bonds Strength depends on the distance between 2 nuclei (or bond length) As length increases, strength decreases!

27 Molecular Formulas Two or more atoms joined together by covalent bonds are called a Molecule Molecular Formula: A system for telling you the type and number of each element in a molecule H 2 O; CO 2 ; C 6 H 12 O 6 Remember: A formula unit for an ionic compound has the same style as a molecular formula. However, the unit tells you the ratio of cations to anions in an ionic compound

28 Diatomic Molecules Molecules are individual particles In comparison, Ionic Compounds exist as large clusters of ions arranged in a crystal lattice Diatomic Elements: Seven elements that, when in their pure form, will bond in pairs. n They are more stable in pairs than alone.

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30 Naming Covalent Compounds The following are the naming rules for a compound that is composed of all non-metals We know the number of atoms from each element by the prefixes used in the name of the compound If there is more than one atom of an element used, then a prefix is needed.

31 Covalent Pre-fixes Number of Atoms Prefix 1 Mono (use only w/ Oxygen) 2 di- 3 tri- 4 tetra- 5 penta- 6 hexa- 7 hepta- 8 octa- 9 nona- 10 deca-

32 Naming Covalent Compounds The second element ends in ide as it did for naming ionic compounds If the vowel combo is o-o or a-o, omit the first vowel (ex: monoxide) IF there is only one atom of the first element, then it DOES NOT get a prefix Example: NCl 3 n nitrogen trichloride

33 Practice P 2 O 5 Diphosphorous pentoxide CCl 4 Carbon tetrachloride CO Carbon monoxide N 2 O Dinitrogen monoxide CF 4 Carbon tetrafluoride SF 6 Sulfur hexafluoride

34 Writing Formula Names Use the prefixes in the names of each element to determine how many of element Write the element s symbol and use subscripts to indicate the number of each element Remember: If the first element has only one, it WILL NOT use mono- Example: Carbon Tetrachloride n CCl 4

35 General Names Some compounds get more common names because they are used often Example: H 2 O (water) CH 4 (Methane) NH 3 (Ammonia) C 6 H 6 (Benzene)

36 Practice Arsenic trichloride AsCl 3 Dinitrogen pentoxide N 2 O 5 Tetraphosphorus decoxide P 4 O 10