Outline Introduction: Multiple bonds, Bond. strength. Naming molecules Drawing Lewis Structures Molecular shapes and VSEPR theory Bond Polarity

Transcription

1 Covalent Bonding

2 Outline Introduction: Multiple bonds, Bond strength Naming molecules Drawing Lewis Structures Molecular shapes and VSEPR theory Bond Polarity

3 Why do atoms bond? Recall that noble gases do not form compounds. This is because their electron arrangements, which consists of a full outer energy level, are stable.

4 Why do atoms bond? Recall that metal atoms and non-metal atoms can gain stability by transferring (gaining or losing) electrons to form ions. Atoms can also share valence electrons to acquire the stable electron configuration of noble gases.

5 Covalent Bonding A covalent bond is the chemical bond that results from sharing valence electrons. Compounds made up of covalent bonding are called molecules. Covalent bonding generally occurs between non-metals, that are near each other in the periodic table.

6 Covalent Bonding Diatomic molecules: Eg. Hydrogen, fluorine, oxygen. Bonding and nonbonding pairs:

7 Covalent Bonding Lewis Structures: Show arrangement of electrons in a molecule. A line represents a single covalent bond in a Lewis structure. Eg.

8 Covalent Bonding Single bonds: Groups 17, 16, 15, 14 Textbook p.244 Q1-6

9 Covalent Bonding Single bonds are also called sigma bonds. Multiple Covalent Bonds: Double: two pairs of electrons are shared eg. Oxygen O 2 molecule Double covalent bonds are made of one sigma bond and one pi bond.

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11 Covalent Bonding Multiple Covalent Bonds: Triple: three pairs of electrons are shared eg. Nitrogen N 2 molecule Triple covalent bonds are made of one sigma bond and two pi bonds.

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13 Covalent Bonding Strength of covalent bonds: Depends on the distance between the bonded nuclei. The distance between the two bonded nuclei at the position of maximum attraction is called the bond length.

14 Covalent Bonding Strength of covalent bonds: More shared pairs = shorter bond length, stronger bond.

15 Covalent Bonding Molecule Bond Type Bond Length Bond-Dissociation Energy F 2 Single 1.43 x m 159 kj/mol O 2 Double 1.21 x m 498 kj/mol N 2 Triple 1.10 x m 945 kj/mol

16 Naming Molecules Binary Molecular Compounds: (2 nonmetals) 1. First element in the formula is named first, using entire name. 2. Second element has suffix ide 3. Prefixes are used to indicate the number of atoms in the compound.

17 Naming Molecules Common prefixes: p.248 Number of Atoms Prefix Number of Atoms Prefix 1 mono 6 hexa 2 di 7 hepta 3 tri 8 octa 4 tetra 9 nona 5 penta 10 deca

18 Naming Molecules Binary Molecular Compounds: (2 nonmetals) Exceptions: First element in the compound name does not use mono. Prefix that ends up with consecutive vowels: drop one vowel. P. 249 Q14-18

19 Naming Molecules Acids: If a compound produces hydrogen ions in solution, it is called an acid. Common acids: Write hydro- prefix in front of the first word. HX (X is a halogen) are all acids.

20 Naming Molecules Common acids: Oxyacids(acid containing a hydrogen atom and an oxyanion). ClO 3- : Chlorate ion HClO 3 : chloric acid ClO 2- : Chlorite ion HClO 2 : chlorous acid NO 3- : Nitrate ion NO 2- : Nitrite ion HNO 3 : Nitric acid HNO 2 : Nitrous acid

21 Naming Molecules Common acids: Oxyacids(acid containing a hydrogen atom and an oxyanion). CO 3 2- : carbonate ion H 2 CO 3 : carbonic acid SO 4 2- :sulfate ion S 2- : sulfide ion P. 251 Q19-30 H 2 SO 4 : sulfuric acid H 2 S : hydrosulfuric acid

22 Structural formulas Molecular formulas show the type and number of each atom in a compound. Structural formulas use the chemical symbols and lines to show bonds and the relative positions of atoms.

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24 Structural formulas We can predict the structural formula for many molecules by drawing the Lewis structure. Simple molecules are easy to draw but for more complicated molecules, it is better to follow a rule.

25 Structural formulas: Lewis Structure Rules 1. The central atom is usually the one more to the LEFT of the periodic table. 2. Write down the sum of all the valence electrons available to bond. 3. Place one bonding pair between the central atom and each of the terminal atoms.

26 Structural formulas: Lewis Structure Rules 4. Subtract the number of bonding electrons from the total number of available electrons. 5. Fill the octets for all the terminal atoms first. 6. Then place any remaining electrons on the central atom.

27 Structural formulas: Lewis Structure Rules 7. If there aren t enough electrons for the central atom to have an octet, then go back and convert one of the lone pairs on the terminal atoms into a double bond. Then check that everything still has an octet. C, N, O, and S often form double or triple bonds. P

28 Structural formulas: Lewis Structure Rules Polyatomic ions: The atoms within a polyatomic ion are covalently bonded. If charge: add electrons to the total number of valence electrons. If + charge: subtract electrons from the total number of valence electrons. Eg. Phosphate ion

29 Structural formulas: Lewis Structure Rules Resonance structures: Some molecules can have equivalent structures, eg. NO 3 - The position of the electron pairs in a resonance structure changes but not the position of the atoms. P.258 Q43-46

30 Structural formulas: Exceptions to the Octet Rule: Exceptions to the Octet Rule: 1. A small group of molecules have an odd number of valence electrons. NO 2, ClO 2, NO

31 Structural formulas: Exceptions to the Octet Rule: Exceptions to the Octet Rule: 2. BH 3 and some other molecules don t have an octet around the central atom.

32 Structural formulas: Exceptions to the Octet Rule: Sometimes, an atom with lone pairs can donate a pair of electrons to another that needs two electrons. Eg. BH 3 + NH 3 This type of covalent bond is called a coordinate covalent bond.

33 Structural formulas: Exceptions to the Octet Rule: Exceptions to the Octet Rule: 3. Some central atoms can form an expanded octet. This just means it can have more than eight valence electrons. Eg. PCl 5, SF 6 P. 260 #47-49, 54, 55

34 Molecular Shapes The shape of a molecule affects its physical properties and chemical properties. The molecular shape, or molecular geometry, can be found after drawing a Lewis structure and using the VSEPR model.

35 Molecular Shapes: VSEPR Model Look at the central atom in a molecule. Each area around it where there are electrons is called an electron domain, or electron density area.

36 Molecular Shapes: VSEPR Model The VSEPR model simply states that each electron domain repels each other as much as possible, and this repulsion is what creates molecular shapes.

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40 Molecular Shapes: Hybridized orbitals To explain the molecular geometry in terms of atomic orbitals, which have various shapes

41 Molecular Shapes: Hybridized orbitals To explain the molecular geometry in terms of atomic orbitals, which have various shapes

42 Molecular Shapes: Hybridized orbitals..we use another theory of atomic orbitals, called hybridized orbitals.

43 Molecular Shapes: Hybridized orbitals A central atom can hybridize its orbitals by combining its s and p valence orbitals. One s + one p = Two sp orbitals One s + two p = Three sp 2 orbitals One s + three p = Four sp 3 orbitals

44 Molecular Shapes: Hybridized orbitals To find out what type of hybridization the central atom has, follow these 3 steps. 1. Draw the Lewis structure. 2. Count how many electron domains there are domains: sp 3 domains: sp 2 4 domains: sp 3

45 Molecular Polarity: Bond polarity A chemical bond between two different elements is never completely ionic or covalent. The ionic or covalent character of a chemical bond depends on how strongly each of the bonded atoms attracts electrons. Recall the term electronegativity: the relative ability of an atom in a chemical bond to attract electrons.

46 Molecular Polarity: Bond polarity

47 Molecular Polarity: Bond polarity We can predict the bond character of a chemical bond by looking at the differences in electronegativity values of the two atoms in the bond. Eg. C: 2.55 O: 3.44 Electronegativity difference: 0.89

48 Molecular Polarity: Bond polarity Electronegativity Difference Bigger than 1.7 Bond Character Mostly ionic Polar covalent Smaller than 0.4 Mostly covalent 0 Nonpolar covalent

49 Molecular Polarity: Bond polarity When electrons in a bond are shared equally, (only between two atoms of the same type) it is called a nonpolar covalent bond. When electrons are shared unequally, it is called a polar covalent bond.

50 Molecular Polarity Reason: Symmetric molecules: bond diploes cancel, making the molecule nonpolar overall. Asymmetric molecules: bond dipoles do not cancel, so the molecule is still polar.

51 Molecular Polarity: Bond polarity We say there is a dipole for polar covalent bonds.

52 Molecular Polarity: Which? Molecules, which are made up of many covalent bonds, can be polar or nonpolar too. 1. If nonpolar covalent bonds only, then molecule is nonpolar. Eg. Cl 2 2. If one polar covalent bond only, then molecule is polar. Eg. HCl 3. If multiple polar covalent bonds and no lone pairs, then molecule is nonpolar. Eg. CCl 4 4. If multiple polar covalent bonds and one or more lone pairs, then molecule is polar. Eg. NH 3

53 Molecular Polarity: Which? Activity

54 Molecular Polarity Polar molecules and ionic compounds can dissolve in polar substances. Nonpolar molecules can only dissolve in nonpolar substances.

55 Properties of Covalent Compounds Weak intermolecular forces The attraction between one molecule and another molecule is weaker than the intramolecular forces. These intermolecular forces are called van der Waals forces.

56 Properties of Covalent Compounds Weak intermolecular forces means that covalent molecular compounds have relatively low melting and boiling points, and are usually gas at room temperature. (weak forces holding molecules with each other) Usually soft solids, eg. wax

57 Properties of Covalent Compounds Covalent network solids: A few solids that are made up of covalent bonds are very strong and brittle, hard, and do not conduct heat and electricity.

58 Properties of Covalent Compounds Type of van der Waals forces: 1. Between nonpolar molecules, weak attraction force = dispersion force. 2. Between polar molecules, stronger attraction force = dipole-dipole force. 3. Between hydrogen of one molecule and F, O, N of another molecule = hydrogen bond