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COVALENT BONDING.

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Presentation on theme: "COVALENT BONDING."— Presentation transcript:

1 COVALENT BONDING

2 COVALENT BOND: Two atoms SHARE a pair of e-
1. Between atoms of similar electronegativities – (< 1.7 electronegativity difference ) 2. Between NONMETALS or NONMETALS & HYDROGEN 3. Hold together MOLECULES a. MOLECULE: smallest particle of a covalently bonded substance

3 H H HOW DO COVALENT BONDS FORM?
1) Molecules can be diatomic elements or compounds Ex. Cl2, C6H12O6, CO2, H2O 4. Formation of covalent bonds – Mostly EXOTHERMIC HOW DO COVALENT BONDS FORM? 1. Orbital overlap and half-filled 1s orbitals H H

4 1s2 and 1s2 He and He NO! Filled 1 s orbitals
One pair of shared electrons But HELIUM He and He 1s and s2 NO! Filled 1 s orbitals

5 NONPOLAR COVALENT BONDS
Electrons are shared EQUALLY between atoms of equal electronegativities - Electronegativity difference is less than .4 Ex: H2,N2, O2, F2, Cl2, Br2, I2

6 H H2 H H F F F2 F e- dot formula Molecular formula
Single covalent bond

7 More Diatomic Elements
e- dot formula Molecular formula O O2 O O Double covalent bond N You try N2 N2 N N Triple covalent bond

8 e- dot formula Molecular formula N N N2 N Triple covalent bond

9 POLAR COVALENT BOND Electrons are shared UNEQUALLY by atoms of DIFFERENT electronegativities - electronegativity difference: .4  1.7 Most bonds are between different elements EX. H2O HCl NH3 CO2 2.1 and 3.5 =

10 Polar Covalent Bonding
e- dot formula Molecular formula H HCl H Cl Cl LINEAR .9

11 Polar Covalent Bonding
e- dot formula Molecular formula H2O O H O 105° H 1.4 Bent 4 e- pairs, 2 bonds

12 Polar Covalent Bonding
e- dot formula Molecular formula NH3 N H N H N H .9 Trigonal pyramid 4 e- pairs 3 bonds

13 Valence-Shell Electron Pair Repulsion Theory (VSEPR)
VSEPR states that e- pairs around an atom try to get as far apart from one another as possible Shapes of molecules are based on this idea

14 Hybridization , , 2s2 2p2 2p3 2s1 (sp3)4
One s & three p orbitals combine to form four sp3 orbitals of equal energy Carbon forms 4 equivalent bonds Why? 2s2 2p2 2p3 2s1 , PROMOTION , (sp3)4 HYBRIDIZATION

15 sp3 hybrid orbitals

16 When bonding, C atoms hybridize:
Ex: Methane CH4 H C C & H 109.5° Shape: tetrahedron

17 Cl C C Cl C Cl EX. CCl4 e- dot formula 4 e-pairs, 4 bonds
Carbon tetrachloride e- dot formula 4 e-pairs, 4 bonds Cl C C Cl C Cl

18 H Cl POLAR MOLECULES Cl – 3.0 H – 2.1 DIPOLE (polar molecule)
a. Define: Molecule with a asymmetrical (uneven) distribution of electric charge (two charged ends) b. Examples: HF, HCl, HBr, NH3, H20 Cl attracts electrons more H Cl Cl – 3.0 H – 2.1

19 c. Polar covalent bonding usually results in polar molecules
Case 1: Case 2: Polar covalent bond between 2 atoms gives a polar molecule Ex: HCl, HBr Molecule with unsymmetrical arrangement of polar bonds Ex: NH H2O PYRAMID H N O H BENT 105°

20 O C H C O O C C d. Nonpolar molecules Case 1:
Can result from polar bonds if there is a symmetrical distribution of charge (bonds) Ex: CH4, CCl4, CO2 O C H C O 2.5 -2.1 .4 O C C

21 Case 2: Nonpolar covalent bonds produce nonpolar molecules
Ex: H2, N2, O2, F2, Cl2, Br2, I2 Diatomic Elements O O Bonds to form….

22 RESONANCE Ex. SO2 S O Resonance occurs when more than one Lewis structure can be drawn for the SAME molecule. The actual structure is an INTERMEDIATE of the possibilities.

23 Be EXCEPTIONS TO THE OCTET RULE Boron compounds – BH3 B
Trigonal planar BeCl2 Be Cl Be Cl linear

24 NAMING COVALENT COMPOUNDS
Binary compounds end in IDE. When more than one atom of a particular element is present in the molecule, the prefixes below are used: di = 2 tri = 3 tetra = 4 penta = 5 hexa = 6 hepta = 7 octa = 8 nona = 9 deca = 10

25 HF(g) = hydrogen fluoride
N2O = dinitrogen oxide NO = N2O4 = N2O5 = NO2 = N2O3 = SO3 = SO2 = SF4 = SF6 = nitrogen oxide dinitrogen tetroxide dinitrogen pentoxide nitrogen dioxide dinitrogen trioxide sulfur trioxide sulfur dioxide sulfur tetrafluoride sulfur hexafluoride

26 COMMON NAMES H2O: CO: water H2O2: NH3: carbon monoxide PH3: N2O:
hydrogen peroxide ammonia phosphine nitrous oxide

27 NAMING ACIDS – BINARY When naming binary acids, the term HYDRO is attached to the front part of the negative ion, and the suffix IC is attached to the end of the negative ion. The word “acid” is then added. HCl(aq): Hydrochloric acid HBr(aq): Hydrobromic acid HF(aq): Hydrofluoric acid HI(aq): Hydroioidic acid

28 NAMING OXYACIDS (Naming acids containing oxygen in the negative group – Refer to Polyatomic ion table) ION ACID ClO- hypochlorite HClO: hypochlorous acid ClO2- chlorite HClO2: chlorous acid ClO3- chlorate HClO3: chloric acid ClO4- perchlorate HClO4: perchloric acid SO32- sulfite H2SO3: sulfurous acid SO42- sulfate H2SO4: sulfuric acid

29 PO33- Phosphite H3PO3 Phosphorous acid PO43- Phosphate H3PO4 Phosphoric acid CO32- Carbonate H2CO3 Carbonic acid CH3COO- Acetate HCH3COO Acetic acid NO2- Nitrite HNO2 Nitrous acid NO3- Nitrate HNO3 Nitric acid

30 Se and Te are chemically similar to S.
If H2SO3 = Sulfurous acid , then H2SeO3 Selenous acid H2SeO4 Selenic acid H2TeO3 Tellurous acid H2TeO4 Telluric acid

31 Br and I are chemically similar to Cl
If HClO = Hypochlorous acid, then HBrO Hypbromous acid HBrO2 Bromous acid HBrO3 Bromic acid HBrO4 Perbromic acid

32 HIO Hypoiodous acid HIO2 Iodous acid HIO3 Iodic acid HIO4 Periodic acid

33 CHARACTERISTICS OF COVALENT (MOLECULAR) COMPOUNDS
1. Much lower melting and boiling points than ionic compounds Why? Weak intermolecular forces of attraction crystals, holding crystal together Whereas,

34 2. States – Exist as solids, liquids, or gases at room temperature
In ionic crystals, very strong forces must be overcome to melt the crystal. 2. States – Exist as solids, liquids, or gases at room temperature a. Again, forces of attraction between the molecules are weak, causing some substances to exist as liquids or gases

35 3. Conductivity – Do NOT conduct, either as liquids or as solids.
Why? Uncharged molecules make them up and are not attracted to electrodes Which dissolved substance is ionic? Which is molecular?

36 Symmetrical distribution
Nonmetals + Nonmetals Nonmetals + H Metals + Nonmetals MOLECULE IONS Molecular Empirical SHARED TRANSFERRED 1.7 or > <.4 .41.7 NONPOLAR POLAR Share e- equally Share e- unequally Symmetrical distribution of electrical charge Nonpolar Polar CH4, CCl4, CO2 H2O NH3 O2, H2

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