CHAPTER 3 CHEMICAL BONDING NUR FATHIN SUHANA BT AYOB SMK SULTAN ISMAIL, JB

CHAPTER 3 CHEMICAL BONDING NUR FATHIN SUHANA BT AYOB SMK SULTAN ISMAIL, JB

LEARNING OUTCOMES (ionic bonding) 1. Describe ionic (electrovalent) bonding such as NaCl and MgCl 2

LEARNING OUTCOMES (metallic bonding) 1. Explain metallic bonding in terms of electron sea model

LEARNING OUTCOMES (intermolecular forces : van der Waals forces and hydrogen bonding) 1. Describe the hydrogen bonding and van der Waals forces (permanent, temporary and induced dipole) 2. Deduce the effect of van der Waals between molecules on the physical properties of substances 3. Deduce the effect of hydrogen bonding (intermolecular and intramolecular)on the physical properties of substances

TYPES OF CHEMICAL BONDING 1. Metal and non-metal : Electron transfer and ionic bonding Metal atom (low IE) loses its valence electrons, non metal (high negative EA) gains electrons 2. Non-metal with non-metal : Electron sharing and covalent bonding A shared electron pair is considered to be localized between the two atoms 3. Metal with metal : Electron pooling and metallic bonding Electron sea model

Subtopic 4.1 LEWIS STRUCTURE

LEWIS SYMBOLS 1. When atoms react to form chemical bonds, only the electrons in the outermost valence shells are involved 2. Valence shell electrons of an atom represented either by cross ( X ) or a dot ( ) 3. It known as Lewis structures or electron-dot structures 4. E.g

LEWIS SYMBOLS

RELATIONSHIP BETWEEN GROUP AND VALENCE ELECTRON Group no. Example Electronic configuration 1 Sodium Lewis diagram 2 Magnesium 13 Aluminium 14 Silicon 15 Phosphorus 16 Sulphur 17 Chlorine 18 Argon

HOW TO WRITE LEWIS SYMBOL Eg. N (Z = 7) Electron configuration : 1s 2 2s 2 2p 3 (valence electron = 5) 1. Identify no. of valence electron 2. Place one dot at a time on the four side (top, bottom, right, left) 3. Pair up the dots until all are used

KEEP IN MIND! 1. Lewis symbols do not show the electron configuration of the valence electron 2. E.g C (Z =6) Electron configuration : 1s 2 2s 2 2p 2 3. C has 4 unpaired dots because it form 4 bonds

KEEP IN MIND! 4. Element in the same group : Similar valence electron configuration Similar Lewis dot symbols 5. E.g. N gains three electron to form N 3 (-3 charge) N can form three covalent bonds Cl gains one electron to form Cl (-1 charge) Cl can form one covalent bond

LEWIS SYMBOL EXERCISE 1 Write Lewis dot symbols for the following atoms: (a) K (b) Ca (c) Be (d) Ga (e) O (f) Br (g) N (h) I (i) As (j) F (k) Mg (l) S

LEWIS SYMBOL EXERCISE 2 Write Lewis dot symbols for the species according to the following electronic configuration: (a)p : 1s 2 2s 2 2p 6 3s 1 (b) Q : 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 (c) R : 1s 2 2s 2 2p 6 3s 2 3p 4 (d) S : 1s 2 2s 2 2p 3 (e) T : 1s 2 2s 2 2p 1 (f) U : 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 (g) V : 1s 2 2s 2 2p 6 3s 2 3p 2 (h) W : 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 14

OCTET RULE 1. An atom other than H tends to form bonds (by losing or gaining or sharing electron) until it is surrounded by eight valence electron 2. It can be achieved by ; E.g : Transfer of electrons Li (Z=3) + F (Z =9) 1s 2 2s 1 1s 2 2s 2 2p 5 Li + F 1s 2 1s 2 2s 2 2p 6

OCTET RULE E.g. sharing of electron

ELECTRON CONFIGURATION OF IONS Noble gas configuration ( 8 valence electron) Eg. 1s 2 2s 2 2p 6 (Ne) Pseudonoble gas configuration eg. [Kr] 4d 10 Half-filled orbitals e.g 3d 5

STABILITIES OF IONS Form stable ions (duplet/octet) Noble Gas Configuration Valence electronic configuration: ns 2 np 6 Example: 1) Na Na + + e - 1s 2 2s 2 2p 6 3s 1 2) F + e - F - 1s 2 2s 2 2p 5 18

NOBLE GAS CONFIGURATION 1. Atoms may lose or gain enough electron so as to forms stable ion with octet (or duplet) configuration (ns 2 np 6 ) 2. Eg. Na Na + + e 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 = [Ne] + e 1s 2 2s 2 2p 6 3s 2 3p 5 1s 2 2s 2 2p 6 3s 2 3p 6 = [Ar]

STABILITIES OF IONS A completely filled orbital but not the noble gas configuration Pseudo Noble Gas Configuration Example: Valence electronic configuration: ns 2 np 6 nd 10 or ns 2 np 6 nd 10 nf 14 1) 31 Ga : 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 1 Ga 3+ : 2) 29 Cu : 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 Cu + : 1 20

PSEUDO-NOBLE GAS CONFIGURATION 1. The (n-1) d 10 configuration of a p-block metal atom that empties its outer level 2. Eg. Sn 4+ + e [Kr] 4d 10 5s 2 5p 2 [Kr] 4d 10 Zn : 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 Zn 2+ : 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 (pseudonoble gas configuration )

STABILITIES OF IONS A special stability of half-filled d orbital Half-filled Orbital Configuration Example: Valence electronic configuration: nd 5 1) 25 Mn : 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 2 Mn 2+ : 2) 26 Fe : 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2 Fe 3+ : 22

HALF-FILLED ORBITALS 1. Some transition metal atoms form cations that have electron configuration associated with half-filled d orbital (d 5 ) 2. E.g. Mn Mn 2+ 2e [Ar] 3d 5 4s 2 [Ar] 3d 5 Fe : 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 Fe 3+ : 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 (stability of half-filled 3d orbital)

IONIC BONDING Electron transfer theory Strength of ionic bond Properties of ionic compound

IONIC BONDING 1. Attractive electrostatic force between positive and negative ions 2. Sometimes called : electrovalent bond 3. Ionic bonds are usually formed between metallic elements (Group 1, 2 and 13) and non-metallic elements (Group 15,16 and 17)

IONIC BONDING Metal atom (more electropositive) Non-metal atom (more electronegative)

FORMATION OF IONIC BONDS 1. By transferring electrons 2. Total number of electron lost by metal atoms = total number of electron gained by the nonmetal atoms 3. E.g. Li Li + + e ******************************************

HOW TO SHOW ELECTRON TRANSFER 1. Three ways : a. Electron configurations Eg. Li + F 1s 2 2s 1 1s 2 2s 2 2p 5 Li + 1s 2 F 1s 2 2s 2 2p 6

HOW TO SHOW ELECTRON TRANSFER 1. Three ways : b. Orbital diagram Eg. Li + F Li + F

HOW TO SHOW ELECTRON TRANSFER 1. Three ways : b. Lewis e-dot symnbol Eg. LiF + F [Li] + [ F ] -

Exercises 3: By using Lewis structure, show how the ionic bond is formed in the compounds below. ( a ) KF ( b ) BaO ( c ) Na 2 O

LEWIS SYMBOL EXERCISE 4 The element X has one electron and the element Y has six electrons in their outermost shell respectively. (a) (b) (c) What is the formula of the compound formed between the elements X and Y? Name the type of bond formed in (a) Draw the Lewis structure to show the formation of compound in (a) 32

Ionic bond is very strong, therefore ionic compounds: 1. Have very high melting and boiling points 2. Hard and brittle 3. Can conduct electricity when they are in molten form or aqueous solution because of the mobile ions

LEWIS STRUCTURE 1. Two dimensional structural formula consists of e-dot symbols that show each atom and its neighbors, the bonding pairs and the lone pairs that fill each atom s outer shell 2. E.g

WRITING LEWIS STRUCTURE 1. Step 1 Draw skeletal structure of compound showing what atoms are bonded to each other Put at least negative electron negative atom in the center F E.g F N F Electronegativity N = 3.0 F = 4.0 N = Central atom F = Surrounding atom

WRITING LEWIS STRUCTURE 2. Step 2 : Count total no. of valence electrons E.g. NF 3 F F N F Atom N X 1 F X 3 Total No. of valence electron 5e x 1 = 5e 7e x 3 = 21e 26 e

WRITING LEWIS STRUCTURE 2. Step 2 : For polynomials ions, add one electron for negative charge of the ion, or subtract one electron for each positive charge E.g NH 4 + Atom No. of valence electron N X 1 5e x 1 = 5e H X 4 1e x 4 = 4e + 1 charge 1e Total 8 e

WRITING LEWIS STRUCTURE 3. Step 3 : complete an octet (8 e ) for all atoms except H (2 e ) Complete the surrounding atoms first electrons not involved in bonding shown as lone pairs E.g Check : 8 e x 3 = 24 e + 2 e 26 e

DRAWING LEWIS STRUCTURE CH 4 1 Determine central atom & count valence e - Central atom : 2 Draw single bond & calculate the non-bonding e - (NBe - ) 3 Complete the octet of the terminal atom 4 Place any remaining e - at the central atom 5 Form double or triple bond if octet rule is not satisfied 39

DRAWING LEWIS STRUCTURE NO 2 + 1 Determine central atom & count valence e - Central atom : 2 Draw single bond & calculate the nonbonding e - (NBe - ) 3 Complete the octet of the terminal atom 4 Place any remaining e - at the central atom 5 Form double or triple bond if octet rule is not satisfied - 1 2 40

WRITING LEWIS STRUCTURE 4. Step 4 : If a central atom does have an octet, make a multiple bond by changing a lone pair from one of the surrounding atoms into a bonding pair to the central atom. E.g N = N N N

STRENGTH OF IONIC BONDS 1. The strength of an ionic bond is a measure of the electrostatic attraction between the ions 2. F Q+ Q d 2 Q + = charge + ve ion Q = charge ve ion d = distance between the ions F = force of attraction 3. The smaller the ions and/or the higher charge on ions > the stronger attraction between ions > the stronger the ionic bond 4. E.g Compound NaCl NaBr Melting point/ 801 750

STRENGTH OF IONIC BONDS 4. The melting point of sodium chloride is higher than that of sodium bromide. This shows that the ionic bond in NaCl is stronger than that in NaBr 5. This is because the Cl ion is smaller than that of Br ion Ion Na + Cl Br Ionic radius/nm 0.095 0.181 0.195 6. Electrostatic attraction between Na + and Cl is stronger.

STRENGTH OF IONIC BONDS 1. The melting point of sodium chloride and magnesium chloride are : Compound NaCl MgCl 2 Melting point/ 801 987 Cation radius/nm 0.095 0.065 2. The Mg 2+ ion is smaller in size than in Na + ion. On top of that, Mg 2+ has higher charge 3. As a result, the ionic bond in MgCl 2 is stronger than that in NaCl. This accounts for the higher melting of MgCl 2.

BOND LENGTH 2. For a given pair of atoms, Bond length : single > double > triple E.g. C C > C = C > C C (154 pm) (134 pm) (121 pm) Bond order increase Stronger bond Shorter bond As the number of bonds between the carbon increase, the bond length decreases because C are held more closely and tightly together As the number of bonds between two atoms increases, the bond grows shorter and stronger * Lebih panjang > lebih mudah break

LEARNING OUTCOMES (covalent bonding) 1. Draw the Lewis structure of covalent molecules (octet rule such as NH 3, CCl 4, H 2 O, CO 2, N 2 O 4, and exception to the octet rule such as BF 3, NO, NO 2, PCl 5, SF 6 ) 2. Explain the concept of overlapping and hybridisation of the s and p orbitals such as BeCl 2, BF 3, CH 4, N 2, HCN, NH 3, H 2 O molecules 3. Predict and explain the shapes of and bond angles in molecules and ions using the principle of valence valence shell electron pair repulsion, e.g. linear, trigonal planar, tetrahedral, trigonal bipyramid, octahedral, v-shaped, seesaw and pyramidal 4. Explain the existence of polar and non-polar bonds (including C- Cl, C-N, C-O, C-Mg) resulting in polar or/and non-polar molecules

LEARNING OUTCOMES (covalent bonding) 5. Relate bond lengths and bond strengths with respect to single, double and triple bonds 6. Explain the inertness of nitrogen molecule in terms of its strong triple bond and nonpolarity 7. Describe typical properties associated with ionic and covalent bonding in terms of bond strength, melting point and electrical conductivity 8. Explain the existence of covalent character in ionic compounds such as Al 2 O 3, All 3, and Lil 9. Explain the existence of coordinate (dative covalent) bonding such as H 3 O +, NH 4 +, Al 2 Cl 6, and [Fe (CN) 6 ]³ˉ

COVALENT BOND 1. Covalent bond is force of attraction between two adjacent nuclei and the electrons that are shared together between them 2. The covalent bond is usually formed between non-metallic elements 3. There are some exceptions. For example, beryllium and aluminium are metals, but they form covalent bonds with chlorine. E.g. BeCl 2, AlCl 3

COVALENT BOND Bonding pair electron Lone pair electron

Covalent compounds: Compounds may have these covalent bonds: i. Single bond ii. Double bond iii. Triple bond.

Lewis structure of water H + O + H H O H 2e - 8e - 2e - or single covalent bonds H O H

Double bond two atoms share two pairs of electrons O C O or O C O 8e - 8e - 8e - double bonds double bonds

Triple bond two atoms share three pairs of electrons N N or N N 8e - 8e - triple bond triple bond

RESONANCE STRUCTURE 1. Two or more Lewis structure for a single molecule that cannot be represented accurately by only one Lewis structure 2. E.g. Ozone (O 3 ) 6 e X 3 = 18e

FORMAL CHARGE 1. Difference between the valence electron in an isolated atom and the number of electron assigned to that atom in a Lewis structure 2. Formal charge of atom : No. of valence electron [ No. of lone pair electron + half of bonding electron] No. of valence electron [ No. of lone pair electron + No. of bonds]

FORMAL CHARGE Negative FC -- on more electronegative atom Positive FC -- on more electropositive atom FORMAL CHARGE (FC) Is used to find the most stable Lewis structure The sum of the FC of the atoms must equal the charge on the molecule or ion FC should be as small as possible 56

FORMAL CHARGE 1. Formal charge : O of O O = 6 6 1 = 1 : O of O = O = 6 4 2 = 0 : middle O = [6 2 3 ] = +1 + 1 + 1-1 0 0-1

SELECTING THE BEST RESONANCE STRUCTURE 1. Select the structure with : All zero formal charge Small formal charge Negative formal charges are placed on the more electronegative atoms

EXAMPLE Calculate the formal charge for each atom of the following compounds: (a) H I H C H I H (b) H N H I H (c) H O H (d) O O O (e) O S O (f) O N O O - 59

EXAMPLE 1) Draw all the possible Lewis structure of COCl 2. 2) Predict the most plausible structure.

SOLUTION 1) 2) The most plausible structure is (2) Formal charge is determined before completing a Lewis structure to predict the most stable structure because formal charge closest to zero.

EXCEPTION TO OCTET RULE 1. Molecular species that do not follow the octet rule fall under two categories ; Molecules in which atom has less than an octet (Incomplete octet) Molecules with an odd number of electrons Molecules in which an atom has more than an octet (Expanded octet)

EXCEPTION TO OCTET RULE 1. Incomplete octet: Electron - deficient molecules The central atoms have fewer than eight electrons around them E.g BeH 2 H Be H Be 2H Total *Be, B, Al incomplete 1 X 2e = 2e 2 x 1e = 2e 4e Other examples : BeCl 2, BCl 3, BF 3

EXCEPTION TO OCTET RULE 2. Odd electron molecules : free radicals Contain an unpaired electron The central atoms have fewer than eight electrons around them E.g NO N O Total 1 X 5e = 5e 1 x 6e = 6e 11e Most odd electrons molecules have a central atom from an-odd numbered group, such as N (Group 15) and Cl (Group 17) Other example : NO 2

EXCEPTION TO OCTET RULE 3. Expanded octet Central atoms have more than eight electrons around them Central atoms are normally elements of Period 3 or higher : d orbital available E.g SF 6 S F Total 1 X 6e = 6e 6 x 7e = 42e 48e Other examples : PCl 5, SO 4 2

EXERCISE Draw the Lewis structure of the following molecules and state the special features at the central atoms. (a) NO (b) TeCl 4 (c) AlBr 3 (d) XeF 2 66

Coordinate Covalent Bond (Dative Bond) 1. Coordinate bond is formed when one of the atom donates both electron (lone pair electron) 2. Known as dative covalent bond or dative bond 3. The atom contributes two electrons to form the coordinate bond is called the donor atom. 4. The atom which accepts the electron pair from the donor atom is called the acceptor atom

Coordinate Covalent Bond (Dative Bond) Example 1 : 1. Ammonium ion, NH 4 + 2. The nitrogen atom in an ammonia molecules has a lone pair of electrons 3. Thus, the nitrogen atom can act as a donor atom. In contrast, the hydrogen ion has an empty 1s orbital. Thus, hydrogen ion (H + ) can act as an acceptor atom. 4. The formation of a coordinate bond in the ammonium ion can be represented by Lewis structure as follows: H H H N H + H + H N H H +

: : Coordinate Covalent Bond (Dative Bond) Example 2: 1. Oxonium ion, H 3 O + 2. The oxygen atom in the water molecule has two lone pairs of electrons. 3. An oxonium ion is formed when one of the lone pairs of electrons is used to form a coordinate bond with a hydrogen ion. 4. The oxonium ion is called the hydroxonium ion or the hydronium ion H H O O + H + H H H +