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04 chemical bonds Presentation Transcript

  • 1. by CHEMISTRY DEPARTMENT CHEM-111 General Chemistry Unit four Chemical Bonds بسم الله الرحمن الرحيم
  • 2. CHEMICAL BONDS
    • Valence electrons
    • Electron dot structures
    • Molecules and Ions
    • Formation of Ions
    • Ionic bonds and Ionic compounds
    • Covalent bonds and molecular compounds
    • Coordinate Covalent bonds
    • Bond polarity
    • Attraction between molecules
    • Shapes of molecules
  • 3. Chemical Bonding
  • 4. I- Intramolecular Chemical Bonding
  • 5. I- Intramolecular Chemical Bonding Intramolecular chemical bonding refers to the attractive forces that holds the atoms together to form a molecule H 2 , O 2 , CH 4 NaCl, MgCl 2 etc. Major classes I) Ionic bonds II) Covalent bonds
  • 6. Because All atoms would like to attain (reach) stable electron configurations like noble gases by: losing electrons or sharing electrons or gaining electrons. For a stable configuration each atom must fill its outer energy level: 1 st : 2 electrons 2 nd : 8 electrons 3 rd : 8 electrons 4 th : 8 electrons Reason for Intramolecular Chemical Bonding
  • 7. Valence Electrons
    • The electrons in outer shell are called valence electrons. These electrons are very influential (important) as they determine chemical properties of the elements.
    • For representative elements, the number of valence electrons in an atom is equal to the group number .
    Group Number 6A IA 4A
  • 8. Core Electrons . Electrons present in shells below the outer shell are called core electrons Core electrons Valence electrons
  • 9. Group Number and Arrangement of Valence Electrons Group Selected Electron Number of Electron-Dot Number Elements Arrangement Valence Electrons Structure 1A Li 2-1 1 2A Ca 2-8-2 2 3A Al 2-8-3 3 4A C 2-4 4 5A N 2-5 5 6A S 2-8-6 6 7A Cl 2-8-7 7 8A Ne 2-8 8
  • 10. I-INTRAMOLECULAR BONDING 1- IONIC BONDS
  • 11. Types of intramolecular Bonding I-Ionic Bonding
    • Ionic bonds are formed due to the transfer of electrons from one atom (metal) to another atom
    • (nonmetal) to give oppositely charged particles called ions to attain a noble gas arrangement.
    Electrically charged atoms or groups of atoms are called ions. They are very reactive species H + , Li + , OH - , Br - etc
  • 12. Types of ions
    • 1- Positive Ions (Cations)
    • Positively charged ions: H+, Na+, K+, Ca+
    • They are formed by loss of an electron
    • (size of cation is smaller than atom BECAUSE A CATION LOSES A SHELL ).
    • 2-Negative Ions ( Anions )
    • Negatively charged ions: Cl- , OH- , I -
    • They are formed by gain of an electron
    • (size of anion is bigger than corresponding atom because entering electron and electrons already present, repel each other and shell get enlarged ).
  • 13. To form an ion, an element has to lose or gain electrons and form stable full outer shell of a noble gas Electronic configuration of lithium, carbon and fluorine Atomic number 3 Atomic number 6 Atomic number 9 Which elements form ions? Many non metals can gain electrons to form negatively charged ions Some non metals with partly filled electrons shell do not form ions All metals can lose electrons to form positively charged ions Lithium Electronic configuration 2,1 Carbon Electronic configuration 2,4 Fluorine Electronic configuration 2,7
  • 14. Group Number Ions formed by the representative elements + IA IIA IIIA IVA VA VIA VIIA Li + Be 2+ C 4- N 3- O 2- F - Na + Mg 2 Al 3+ Si 4- P 3- S 2- Cl - K + Ca 2+ Se 2- Br - Rb + Sr 2+ Te 2- I - Cs + Ba 2+
  • 15. A positive ion produced by loss of one or more electrons from metal atom. Li + , Na + ,K + , Mg 2+, Ca 2+ , Al 3+ , etc. Anion A negative ion formed by gain of one or more electrons by non metal. F 1- ,Cl 1- , O 2- , N 3- , C 4- etc X c Charge Element Ions formed by the Representative Elements Cation
  • 16. Ionic bonds are formed between a metal and a non-metal . Bonding occurs using valence electrons . Ionic Bonding and Octet Rule
    • Atoms that have 5, 6 or 7 electrons in their outer levels will tend to gain electrons from atoms with 1, 2 or 3 electrons in their outer levels
    • Atoms that have 1, 2 or 3 electrons in their outer levels will tend to lose them in interactions with atoms that have 5, 6 or 7 electrons in their outer levels.
    • Atoms try to get 8 electrons in the valence shell (octet rule)
  • 17.
    • Formation of sodium chloride or common salt
    • Na atom looses an electron and forms Na+ ion.
    • Cl atom gains an electron and forms Cl- ion.
    NaCl (Ionic compound) Oppositely Charged ions, Na + and Cl - attract each other and form ionic compound, NaCl.
  • 18. Ionic Giant Molecule of NaCl Ionic Giant Molecule
    • Ionic bond is the attraction between cations and anions
    • Na + and Cl - ions are stacked together
    • Lowest energy arrangement
    • The pattern is repeated throughout the crystal or Lattice
    Cl
  • 19. Loses 2 e - Each gains 1 e - One Magnesium ion Mg 2+ Two Chloride ions 2Cl - Magnesium and Chloride ions Combine to form MgCl 2 (Magnesium chloride) Formation of Magnesium chloride MgCl 2
  • 20. Lewis structure
      • Lewis symbols help us to find the valence electrons and predicts bond
    Magnesium chloride Sodium chloride
  • 21. PROPERTIES OF IONIC COMPOUND
    • a. They have high melting and boiling points due to presence of strong intermolecular force of attraction .
    • b. They are non-volatile : they do not easily turn into gas.
    • c. Generally soluble in water because they generally ionize in water .
    • d. Solids do not conduct electricity because the ions are held together by strong electrostatic force of attraction.
    • e. Conduct electricity when melted .
    • f. Aqueous solution conducts electricity because the ions are mobile in solutions.
    • g. They are generally insoluble in organic solvents, because organic solvents are covalent compounds.
  • 22. The noble gases are known for their chemical stability and existence as mono atomic molecules as they contain 8 electrons in outer most shell (valence shell) except Helium Nobel gases He Ne Ar Kr Xe Rn Valence electrons 2 8 8 8 8 8 Period no. 1 2 3 4 5 6 No. of shells 1 2 3 4 5 6 Other atoms also want electronic configurations just like noble gases (concept of duplet & octet rules) due to their Stability so they form ionic bond Why Ionic Bond are Formed? Reasons 1.Stability of noble Gases
  • 23. Li 1 1 H 2 3 He C 6 12 4 7 Be B C N O F Ne 4 5 6 7 8 9 10 9 11 12 16 14 19 20 Na Mg Al Si P S Cl Ar 11 12 13 14 15 16 17 18 23 24 27 28 31 32 35.5 40 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 19 20 21 22 23 24 25 26 27 28 29 35 36 30 31 32 33 34 39 40 45 48 51 52 55 56 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Ti Pb Bi Po At Rn Fr Ra Lr Rf Db Sg Bh Hs Mt Uun Uuu Uub Uut Uuq Uuh IA IIB IIIB IVB VIB VB VIIB VIIIB IB IIA IIIA IVA VA VIA VIIA VIIIA 1 2 3 4 5 6 7 84 59 59 63.5 65 70 73 75 79 80 38 39 40 37 41 42 43 44 46 45 47 48 49 50 51 52 53 54 La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Ac Th Pa U Np Pu Am Cm Bk Cy Es Fm Md No 55 56 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 87 89 88 90 91 92 93 94 95 96 97 98 99 101 100 102 103 104 105 106 107 108 109 110 111 112 113 114 116 85.5 88 89 91 93 96 (98) 101 103 106 108 112 115 119 122 131 133 137 128 127 175 178.5 181 184 186 190 192 195 197 201 204 207 209 (209) (210) (222) (223) 226 (260) (261) (262) (263) (262) (265) (266) (269) (272) (277) (289) (289) 138 140 141 144 (145) 150 152 157 159 162.5 165 167 169 173 227 232 231 238 237 (244) (243) (247) (247) (251) (252) (257) (258) (259) NON METAL C H Hg SOLID GAS LIQUID METAL METALLOID PERIODIC TABLE OF ELEMENTS
  • 24. ………… . . . 75 -436 0 Distance between hydrogen nuclei (pm) Potential energy (k/mol) 2.Change in Potential Energy Formation of ionic compound must be exothermic i.e. P.E of ionic compound (product) is less than reactants atoms. Why Ionic Bond are Formed? + ………………………… .. (a) (b) (c) (d) 75pm
  • 25. Simple Binary (double) Compounds of Metals and a Nonmetals
    • One element is metal and the other is a non-metal
    • The metal (+ve) is name first , followed by
    • the nonmetal (-ve)
    • Only two different elements are present
    • Number of each is given as a subscript
    • Overall formula must have no Charge
    • Examples:
    • metal nonmetal
    • NaCl Sodium chloride
    • SrO Strontium oxide
    • Al 2 S 3 Aluminum sulfide
    • Binary compounds = Compound composed of two different elements
    • In ionic compounds one is metal and the other is non-metal.
  • 26. Naming ionic compounds
    • IUPAC Methods of Naming
    • Firstly, Name the cation (metal)
    • When an element has only one charged state
    • Secondly, Name the anion (nonmetal) using the ending (-ide)
    • Examples
    • CrCl 3 Chromic chloride,
    • MgBr 2 Magnesium bromide,
    • Al 2 O 3 Aluminum oxide,
    • K 3 N Potassium nitride,
    MgS Magnesium sulfide Ca 2 C Calcium carbide SrO Strontium oxide
  • 27. Naming of Ionic Compounds Formula for some ionic compounds NaCl MgCl 2 AlCl 3 Na 2 O MgO Al 2 O 3 Na 3 N Mg 3 N 2 AlN Some simple ions Cations Na + Mg 2+ Al 3+ Anions Cl - O 2- N 3- Simple short cut can be used to obtain the formula of an ionic compound by exchanging superscript for subscript
  • 28. Naming ionic compounds B group metals may have more than one possible charge ( due to more than one oxidation state). Use roman numerals (I,II,III) in the name to indicate the charge on the metal. . Cu 1+ + O 2- = Cu 2 O copper (I) oxide copper (I) oxide Cu 2+ + O 2- = CuO copper (II) oxide copper (II) oxide
  • 29. Naming ionic compounds FeBr 3 Each bromide is 1 - so iron 3 + FeBr 3 Iron (iii) bromide FeCl 2 iron (ii) chloride FeCl 3 iron (iii) chloride SnBr 2 tin (ii) bromide SnBr 4 tin (iv) bromide AgCl silver chloride CdS cadmium sulfide There are some exception Some B metals only have a single state so the roman numeral may be omitted Some A metals have more than one State so numbers must be used
  • 30. Some ions have characteristic colors they impart to their compounds as seen below Solution of potassium chromate, which contains the yellow chromate ion,CrO 4 2- . Polyatomic ions Solution of potassium permanganate, which contains the violet permanganate Ion, MnO 4 - Solution of potassium dichromate ion, Cr 2 O 7 2- . A poly atomic ion is a group of nonmetal atoms (which are covalently bonded together), that carries electrical charge
  • 31. Some common polyatomic ions Cation Name of cation Anion Name of anion NH4+ ammonium ClO 2 - chlorite H3O+ hydronium ClO - or OCl - hypochlorite PO 4 3- phosphate HPO 4 2- hydrogen phosphate H 2 PO 4- dihydrogen phosphate CrO 4 2- chromate Cr 2 O 7 2- dichromate MnO 4 - permanganate CH 3 CO 2 - acetate NO 3 - nitrate OH- hydroxide O 2 2- peroxide SO 4 2- sulphate
  • 32. AgNO 3 silver nitrate H 2 O 2 hydrogen peroxide Fe 2 (SO 4 ) 3 iron (III) sulphate Naming compounds containing polyatomic ions Most polyatomic ions have names that end with -ate or- ide. NH 4 Cl ammonium chloride NaOH sodium hydroxide KMnO 4 potassium permanganate (NH4) 2 SO 4 ammonium sulfate Write the positive ion, first, and then write the name of the polyatomic ion
  • 33. Chemical Formulas of complex compounds Example: Copper sulfate CuSO 4 .5H 2 O CuSO 4 .5H 2 O. Blue crystal CuSO 4 White powder heat
  • 34. I-INTRAMOLECULAR BONDING 2- COVALENT BONDS
  • 35. Bonds are formed between two non metals by sharing of one or more pairs of electrons. CH 4 , CCl 4, O 2 , H 2 , H 2 O, HCl, HF , NH 3 , CO 2 COVALENT BONDS The attractive forces between two atoms dominate (take over) the repulsive forces. Total potential energy decreases as a result bond is formed at a short distance, called bond distance.
  • 36. P P P P P P P P e - e - e - e - e - e - e - e - D Combination of forces A No interaction B Attraction begins C Covalent bond + + + + + + + + Formation of covalent bond Both nuclei repel each other, as both electron clouds do The nucleus of one atom attracts the electron clouds of the other atom , vice versa electron share
  • 37. Formation of Covalent Bond Example: Formation of ICl Molecule Iodine Chlorine
  • 38. 1. H . + . H H :H or H­H 2. H . + . F H : F or H-F Or methane H 2 HF CH 4 Examples Electron-Dots Structures Molecular Models
  • 39. Cl 2 molecule a shared electron pair forms a single covalent bond Other ways of showing a chlorine molecule: Other ways showing a hydrogen chloride molecule: Covalent Compounds Cl-Cl
  • 40. An oxygen molecule two shared electrons pairs form a double covalent bond double bond A nitrogen molecule Other ways of showing a nitrogen molecule three shared electrons pairs form a triple covalent bond triple bond Other ways of showing an oxygen molecule Multiple Covalent Compounds
  • 41. Lewis Structure When two hydrogen atoms share a pair of electron the covalent bond is formed. 1. H• + H •  H••H 2.
  • 42. More Examples of Lewis Structure Multiple Covalent Bonds Covalent Bonds
  • 43. COORDINATE COVALENT BONDS
    • The type of bond, in which a pair of electrons from one atom is shared by two atoms, is called
    • a coordinate covalent bond .
    • Example
    • The reaction of boron tri-chloride, BCl 3 , and ammonia, NH 3 .
    Coordinate covalent bond.
  • 44.
    • PROPERTIES OF COVALENT COMPOUNDS
    • Low melting and boiling points because of weak
    • intermolecular binding force
    • b. They are generally volatile because of their low M.P. and B.P.
    • c. They are mostly insoluble in water
    • d. Even if some covalent compounds dissolve in water they do not form ion and remain almost undissociated.
    • e. They are generally soluble in organic solvents.
    • f. They are non-conductors of electricity in solid state .
  • 45. Polarity of molecules Electrons in a covalent bond rarely get shared equally. Unequal sharing results in polar bonds Slight positive side Slight negative side Smaller electro negativity Larger electro negativity Polar covalent bonds Electrons may not be shared equally . This is based on the electro-negativity difference between the two elements forming the bond. If: 1.Electronegativity difference > 1.7 (ionic bond) 2.Electronegativity difference = 0 (covalent non-polar bond) 3.Values between 0 and 1.7 are polar covalent bond – In polar covalent electrons are not shared equally
  • 46. Electro-negativity increases across the period & decreases down the group Increase Decreases
  • 47. Shapes of Molecules Electron Dot Bonded Molecular Molecule Structure Atoms Shape (angle) CH 4 4 tetrahedral ( 109 0 ) NH 3 3 pyramidal (109 0 ) H 2 O 2 bent (109 0 ) BCl 3 3 planar-trigonal
  • 48. Molecular Geometry Molecules have specific shapes which is determined by the number of electron pairs around the central species. An electron group can be the electron pairs bonded to atoms or a lone pair . Geometry affects factors like polarity and Solubility. Multiple bonds are treated as a single bond for geometry
  • 49. Characteristic physical properties Boiling point The temperature at which a liquid is converted to a gas at atmospheric pressure . Melting point The temperature at which a solid is converted to a liquid. Compounds Bond Mp Bp N 2 nonpolar -210 - 196 O 2 nonpolar -219 -183 NH 3 polar -78 -33 H 2 O polar 0 100 NaCl ionic 804 ?
  • 50. II- Intermolecular Attractions
  • 51. Forces between Molecules that determine Physical properties Intermolecular Attractions Types of Intermolecular Attractions 1 Dipole – dipole attractions 2. Hydrogen Bonding 3. London Forces
  • 52.
    • These attractions are weaker than ionic or covalent bonds (i.e. only 1% strong).
    • These attraction are due to the electrons which make the bond but
    are not equally shared between atoms because of a difference in electro-negativity of these atoms . δ+ and δ- ends are attracted to each other . δ+ δ- δ+ δ- δ+ δ- 1- Dipole–Dipole Attractions
  • 53.
    • 2. Hydrogen Bonding
    • (dipole-dipole interaction )
    • The attraction between a hydrogen atom in one molecule and an electronegative atom (F,O,N) in another molecule is hydrogen bonding
    • H δ+ —F δ-  H δ+ —F δ-  H δ+ —F δ-
    • This attraction is responsible for high boiling point of water
  • 54. Polar water molecule Hydrogen bonding Hydrogen bonding produces strong attractions between water molecules Hydrogen bonding in water The water molecule is very polar
  • 55. Water Ice The density of ice is 0.931 gm/cubic cm. This compares with a density of 1.00 gm/cubic cm. for water so ice floats on water In liquid water each molecule is hydrogen bonded to approximately 3- 4 other water molecules In ice each molecule is tetrahedrally hydrogen bonded to 4 other molecules. which controls the orientation such that ice has empty spaces which makes ice lighter Tetrahedral shape Density of Ice
  • 56. 3.London Dispersion Forces Van der Waal forces The attraction of positively charged nucleus of one atom for electron cloud of an atom in nearby molecule. These are relatively weak forces and exist in symmetrical non-polar molecules like CH 4 , SO 3 , CO 2 , O 2 , H 2 , He Random motion of electrons for instant causes electrons more on one side & instantaneous dipole is produced that induces a dipole in its neighbor.
  • 57. London forces As the dipole forms in atom A, it induces a dipole in atom B London forces are momentary and called instantaneous dipole δ- δ+ δ+ δ-
  • 58. The electron density fluctuates in two neighboring atoms, giving rise to fleeting attractions between the momentary dipoles. Flickering On and Off of the Dipoles No Polarization Instantaneous Dipoles on atom A induces a dipole on atom B A B A B
  • 59. Exercise CHEMICAL BONDING M.C.Q: Choose the best answer .
    • Force of attraction that holds atoms in a molecule together is:
    • a) covalent bond b) ionic bond
      • c) chemical bond d) coordinate bond
      • 2. If atoms gain or lose electron, they will form:
    • a) ions b) cations
    • c) anions d) Polar covalent bond
    • 3. Indicate the members of nonmetals:
    • a) Lithium b) Oxygen
    • c) Chlorine d) both b & c
  • 60. M.C.Q
    • 4 . The formation of ions from neutral atoms is nearly always:
      • a) endothermic b) exothermic
      • c) both a & b. d) none of the above
    • 5. In making bonds atoms obey octet rule, so number of electrons in their outer shell may be:
      • a) 2 or 8 electrons b) 6 or 8 electrons
      • c) 8 or 10 electrons d)18-electrons
    • 6. Choose the polar compound:
    • a) CH4 b) NaCl c) HCl d) Both a & c
    • 7. The bond formed in CH 4 molecule is by:
    • a) Loss of electrons b) Gain of electrons
    • c) Sharing of electrons d) Both a & b
  • 61.
    • 8. The loss of electron/electrons from outer shell of metals make them:
      • a) positive ion and smaller in size than corresponding atom
      • b) negative ion and smaller than atom
      • c) positive ion and bigger than atom
    • 9. Gain of electron/electrons by non metals changes them into:
      • a) positive ions b) negative ions
      • c) neutral atoms d) both a & b
    • 10. Example of anion:
    • a) K + b) O - 2 c) Mg +2 d) Na
    • 11. Ionic compounds:
    • a). HF b) Ca O C) H 2 O d). Both a & b
  • 62.
    • 12. If difference between electro negativity values of the atoms is > 1.7 they will form:
    • a) Ionic bond b) Covalent bond
    • c) Polar covalent bond d) Non polar covalent bond
    • 13. The force of attraction among water, H 2 O ,molecules
    • is:
    • a) Dipole-dipole interaction b) London forces
    • c) Hydrogen bonding d) None of the above.
    • 14The molecule that shows London dispersion forces is:
    • a) NaCl b) CO 2 c) H 2 O d) MgO
    • 15. What are Intermolecular Forces?
    • a) Hydrogen bonding b) London Forces
    • c) Dipole-dipole interaction d) All of these
  • 63.
    • 16. The atoms complete their outer shells by obeying octet rule and
    • attain electronic configuration of:
      • a) Alkali metals b) Halogens c) noble gases d) Transition metals
    • 17. Energy required to break a chemical bond & form a neutral
    • isolated atom is called:
    • a) Bond energy b) Bond order
    • c) Bond distance d) both a & b
    • 18. Find out the chemical formula of organic solvent named carbon
    • tetra chloride:
    • a) :CCl 2 b) CH 4 c) CCl 4 d) CHCl 3
    • 19. Identify the ionic molecule:
    • a) NH 3 b) N 2 c) H 2 O d) NaCl