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Chapter 7

Chapter 7 . Ionic & Covalent Bonds. Ionic Compounds Covalent Compounds. EN difference and bond character >1.7 = ionic 0.4 – 1.7 = polar covalent <0.4 = nonpolar covalent. Ionic >1.7. Electrons not shared at all

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Chapter 7

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  1. Chapter 7 Ionic & Covalent Bonds

  2. Ionic Compounds • Covalent Compounds

  3. EN difference and bond character >1.7 = ionic 0.4 – 1.7 = polar covalent <0.4 = nonpolar covalent

  4. Ionic >1.7 • Electrons not shared at all • Atom with higher electronegativity takes electrons from atom with lower electronegativity • NaCl • K2O

  5. Polar Covalent 0.4 – 1.7 • Electrons are shared by not equally • More electronegative atom pulls electrons closer • HCl

  6. Nonpolar Covalent <0.4 • Electrons equally shared • Directly in the middle of the two atoms • H2 • O2

  7. 7.1 Ion Formation • Ions are formed when atoms gain or lose valence electrons to achieve a stable octet electron configuration

  8. Valence Electrons and Chemical Bonds • Recall: • All elements in the same group have the same number of valence electrons and therefore…

  9. Valence electrons are involved in the formation of chemical bonds • The force that holds two atoms together • Attraction between positive and negative ions

  10. Recall: Dot structures • Show only valence electrons • Carbon: 1s2 2s2 2p2 • Bromine: [Ar] 4s2 3d10 4p5

  11. Recall: octet rule • Atoms will gain, lose, or share electrons to acquire the stable electron configuration of a noble gas • Metals – • Nonmetals -

  12. Positive Ion formation • Name: • Formed by: • Group 1: • Group 2: • Group13:

  13. Transition metal ions • Form cations only • Charges may vary in some atoms • Fe can lose 2 or 3 electrons • Fe2+ or Fe3+ • Periodic table

  14. Negative Ion Formation • Name: • Formed by: • Group 15 • Group 16 • Group 17

  15. 7.2 Ionic Bonds & Ionic Compounds • Oppositely charged ions attract each other, forming electrically neutral compounds

  16. Formation of an Ionic Bond • Ionic bond – the force of attraction that holds oppositely charged ions together

  17. Ionic compounds – compounds that contain ionic bonds • Cation + anion • Metal + nonmetal • Also called salts

  18. Binary ionic compounds – contain two different elements (a metal + a nonmetal) • NaCl • MgO • K2S • CaI2

  19. Compound formation and charge • Ionic compounds are electrically neutral • Total positive charge must = total negative charge • Net charge of all ionic compounds = 0

  20. Formation of Sodium Chloride • Na: [Ne]3s1 + Cl: [Ne]3s23p5 • Dot Structure: Na + Cl  [Na]+ +[ Cl ]-

  21. Na2O

  22. Al2O3

  23. How would an ionic compound form from each of the following: • Na + N • Li + O • Sr + F • Group 1 + group 15

  24. Formulas for ionic compounds • Formula unit = the chemical formula for an ionic compound • Simplest ratio of ions involved • Mg6Cl12 • MgCl2 • Overall charge = 0

  25. Monatomic ions – one atom ions • Ex: • Oxidation number - the charge of a monatomic ion • Most transition metals have more than one oxidation number

  26. What is the oxidation number of the ions in the following compounds? • FeO • MgCl2 • Cu3N • Cu3N2 • Fe2O3

  27. Properties of ionic compounds • Physical structure – ions are packed into a regular repeating pattern

  28. Crystal lattice – 3D geometric arrangement of particles in an ionic compound • Formed by the strong attractions among positive and negative ions • Each positive ion is surrounded by negative ions & each negative ion is surrounded by positive ions

  29. Physical properties – ionic bonds are very strong, take a lot of energy to be broken apart • High melting point • High boiling point • Hard, rigid, brittle solids

  30. More physical properties • Brilliant colors – due to transition metals in crystal lattices • Electrolytes when dissolved or melted • Conducts electricity • IONIC SOLIDS DO NOT CONDUCT ELECTRICITY

  31. Energy and the Ionic Bond • Exothermic reactions • Endothermic reactions • Formation of ionic compounds always releases energy & therefore is…

  32. Lattice energy – the amount of energy required to separate 1 mol of ions in an ionic compound • Greater lattice energy = stronger force of attraction

  33. Lattice energy is directly related to size of ions bonded • Smaller ions = stronger bond • Which is stronger KCl or LiCl?

  34. Lattice energy is also related to the charge of the ions • Bond formed from attraction of ions with larger charges = stronger • Which is stronger MgO or NaF?

  35. What is a covalent bond? • Bond in which atoms share electrons • Always 2 nonmetals • Called a molecule

  36. Diatomic molecules H2 N2 O2 F2 Cl2 Br2 I2 • Exist because they are more stable than individual atoms

  37. Single Covalent Bonds • One pair of electrons is shared • Cl – Cl • H - F

  38. Sigma bonds (σ) • Pair of shared electrons is in an area centered between the two atoms • Valence orbitals overlap which concentrates electrons in a bonding orbital between the two atoms • s overlaps with s or p or two p overlap

  39. Double Covalent Bond • Two pairs of electrons are shared • Typically C, N, O, S • O = O

  40. Triple covalent bond • Three pairs of electrons are shared • Typically C or N • N N

  41. The pi Bond • Forms when parallel orbitals overlap and share electrons • A double bond = 1 sigma + 1 pi bond • A triple bond = 1 sigma + 2 pi bonds

  42. If atoms need 1 electron, it will usually form 1 covalent bond. • H and Halogens typically only form one bond • If atoms need 2 electrons, it will usually form 2 covalent bonds. • If atoms needs 3 electrons, it will usually form 3 covalent bonds.

  43. Strength of Covalent Bonds • Bond length – the distance between bonded nuclei • Determined by size of bonding atoms & how many electrons are shared • As more electrons are shared, bond length decreases • Cl2 = 1.43 x 10-10 m • O2 = 1.21 x 10-10 m • N2 = 1.10 x 10-10 m

  44. Shorter bond length = stronger bond • Which bond type is weakest? • Which bond type is strongest?

  45. Bonds and energy • Energy released when bonds form • Energy absorbed to break bond • Bond dissociation energy • Smaller bond length = larger bond dissociation energy

  46. 7.3 Molecular Structures • Molecular formula tells the type and number of atoms in a molecule • PH3 • Lewis structure uses electron dot structures to show how electrons are arranged in molecules

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