Covalent Compounds

1. Covalent Compounds

2. Why do atoms bond? • When a + nucleus attracts electrons of another atom • Or oppositely charged ions attract( ionic bonds-metals and nonmetals) • We know the octet rule- atoms tend to gain or lose or share e- in order to acquire a full set of 8 valence e-

3. Properties • 1)  Covalent compounds generally have much lower melting and boiling points than ionic compounds.  • 2)  Covalent compounds are soft and squishy (compared to ionic compounds, anyway • On the other hand, covalent compounds have these molecules which can very easily move around each other, because there are no bonds between them.  As a result, covalent compounds are frequently flexible rather than hard.

4. More Properties • 3)  Covalent compounds tend to be more flammable than ionic compounds. • 4)  Covalent compounds don't conduct electricity in water. • 5)  Covalent compounds aren't usually very soluble in water.

5. More Properties • There's a saying that, "Like dissolves like".  This means that compounds tend to dissolve in other compounds that have similar properties (particularly polarity).  Since water is a polar solvent and most covalent compounds are fairly nonpolar, many covalent compounds don't dissolve in water.  Of course, this is a generalization and not set in stone - there are many covalent compounds that dissolve quite well in water.

6. Covalent Bond • Chemical bond that results from sharing valence electrons. (Generally occurs when elements are close to each other on the periodic table and between nonmetallic elements) • Molecule- a another name for covalent compound, is formed when two or more atoms bond covalently.

7. more • Diatomic molecules- are molecules that occur in nature not as single atoms because the molecules formed are more stable than the individual atoms. There are 7 of these: Hydrogen (H2), Nitrogen (N2), Oxygen (O2), Fluorine (F2), Chlorine (Cl2), Bromine (Br2), and Iodine (I2).

8. Naming Covalent Compounds • Rule 1. The element with the lower group number is written first in the name; the element with the higher group number is written second in the name. • Exception: when the compound contains oxygen and a halogen, the name of the halogen is the first word in the name. • Rule 2. If both elements are in the same group, the element with the higher period number is written first in the name.

9. More Naming Rules • Rule 3. The second element in the name is named as if it were an anion, i.e., by adding the suffix -ide to the name of the element.Rule 4. Greek prefixes are used to indicate the number of atoms of each nonmetal element in the chemical formula for the compound. • Exception: if the compound contains one atom of the element that is written first in the name, the prefix "mono-" is not used.

10. Greek Prefixes • mono-1di-2tri-3tetra-4penta-5hexa-6hepta-7octa-8nona-9deca-10

11. Covalent Bonds • When showing the bonding between atoms of covalent compounds, either a pair of dots or a line is the Lewis structure of a molecule. Lewis structures- use electron-dot diagrams to show how electrons are arranged in molecules

12. more • Group 7A atoms have 7 valence e- and need 1 e- to fulfill the octet rule and become stable. So they will form a single covalent bond • Do example • Group 6A need 2 e- and will form 2 covalent bonds • Group 5A need 3e- and will form 3 covalent bonds • Group 4A will form 4 covalent bonds. • Let’s practice!!!

13. Molecular Structures • Several models can be used to represent a molecular structure: Molecular formula, structural formula, Lewis structure and ball and stick. • Predicting the location of certain atoms: • Hydrogen is always terminal • Halogens usually terminal • Elements with more than one atom are usually terminal • Central atom- smallest electronegativity • Find total number of electrons for bonding9 total valence electrons) • Determine pairs (divide valence electrons by 2) • Draw Bonds from central atom to other atoms • Subtract bonded e- for total valence e- • Place remaining e- around to complete octets • If there are not enough electrons to give the atoms 8, except for hydrogen, make double and triple bonds. • C,N,O and S can form double and triple bonds with same element and others.

14. Single and Multiple Bonds • The sigma bond(σ)- single covalent bond (overlapping of the valence atomic orbitals resulting in the e- being in a bonding orbital between the two atoms) • Multiple Covalent Bonds- atoms can attain a noble gas configuration by sharing more than one pair of electrons between two atoms.

15. Atoms can have • Single - 1 pair of e- • Double – 2 pair of e- • Triple – 3 pair of e- • Try CO2, O2, N2

16. Electronegativity and Lewis Dot Structures • When these atoms share electrons they are often not equally shared. This is due to electronegativity. • Hydrogen and halogens usually bond to only one other atom and are usually on the outside of the molecule. • The atom with the smallest electronegativity is often the central atom of the compound. • When a molecule contains more atoms of one element than the others, these atoms often surround a central atom.

17. More • A pi bond π is formed when parallel orbitals (p orbitals )overlap to share electrons • Single bond- sigma bond • Double bond- sigma and 1 pi bond • Triple – sigma and 2 pi bonds

18. Strength of Covalent Bonds • Covalent bonds involve attractive and repulsive forces • Nuclei(+) and electrons(-) attract each other but • Nuclei(+) repel nuclei(+) and • (e-) repel (e-)

19. More • There is a balance of attractive and repulsive forces • When this balance is disrupted the bond could break • Factors controlling Bond Strength • 1. distance between nuclei= bond length

20. More • Bond length • Determined by size of atoms and number of shared electrons • Size- length decreases when number of bonds increases • Example- a triple bond ( 3 shared pairs of e-) is shorter than a double bond (2 shared pairs of electrons) • Strength- shorter the length –stronger the bond

21. Bond Energy • Energy change occurs when bonds are formed • Bond Dissociation Energy- the amount of energy required to break a specific covalent bond- this is always a + value because energy is added to break a bond • The sum of BDE for all bonds in a compound= the amt. of chemical PE available in one molecule of a compound • BDE= the strength of the bond • **as two atoms are bonded closer together the BDE increases • Formation- energy released-exothermic • Broken- energy added-endothermic

22. Examples of BDE and Bond Length BDE Bond Length • F2 F---F 159 kJ/mole short • O2 O=O 498 kJ/mole shorter • N2 NtripleN 945 kJ/mole shortest

23. Exceptions • Some molecules and ions do not obey the octet rule. • Reasons • 1. some group has an odd number of valence e- and can’t form an octet • Ex. NO2 • 2. some groups form with few than 8 e- present. This is very reactive and will share an entire pair of e- which is a coordinate covalent bond. • exBH3 • 3. central atom contains more than 8 e- (expanded octet)- usually occurs with elements in period 3 or higher. • Ex. PCl5 or SF6

24. Resonance Structures • A condition when more than one valid Lewis structure can be written for a molecule or ion. This occurs when a molecule or ion has a combination of single and double bonds. • Let’s try to draw the resonance structures of NO3- • Now you try for homework • SO3, SO2, O3, NO2-

25. Molecular Shapes • Valence Shell Electron Pair Repulsion • VSEPR model • Many reactions depend on the ability of two compounds contacting each other. The shape of the molecules determines whether they can get close enough to react. • The repulsion of electron pairs in a molecule result in atoms existing at fixed angles to each other- Bond angle

27. Electronegativity and Polarity • We know electronegativity( the attraction an atom has for electrons) increases as we go up and to the right of the periodic table. So Fluorine is the most electronegative element. • If we have any of the diatomic molecules, we have identical atoms bonded together, their electronegativities are the same so they are nonpolar molecules. There is equal sharing of electrons

28. More • Chemical bonds between elements are never completely ionic or covalent. • When the difference in electronegativity between atoms increases it is more ionic • When the difference in electronegativity between atoms is less it is more covalent • Unequal sharing of electrons results in a Polar Covalent Bond

29. More

30. More • Molecules are either polar or non polar • With polarity comes a partial (+) and partial (-) • Show how partial charges spread out on • H2O • CCl4

31. 3 resonance structures

32. Dichromate Ion • Lewis Structure • Cr2O72- 56 valence electrons