Unit 6: Ionic & Covalent Compounds, Lewis Structures Chapter 5 & 6

Key Concepts Ionic, Covalent, Metallic Compounds Polar/Nonpolar Covalent Bonds Bond Length, Bond Strength Resonance Structures Lewis Dot Structures VSEPR Theory Molecular Shapes Molecular Polarity

Compounds Ionic (M/NM): Covalent (NM/NM): Metallic (M/M): e- are transferred from a metal to a nonmetal forming charged ions that stick together Covalent (NM/NM): e- are shared in overlapping molecular orbitals Metallic (M/M): e- move freely between atoms

Octet Rule Octet Rule: atoms gain or lose e- to attain an e- configuration of the nearest noble gas Duet Rule: 2 val e- for hydrogen Octet Rule: 8 val e- for all others

Ions Ion: atom/molecule with a charge +1 +2 ±4 -3 +3 -2 -1 TM form ions with various charges Elements always gain or lose the same number of e- when they form ionic compounds (except TM) When an atom or molecule gains a negative electron, it becomes a negative ion When an atom or molecule loses a negative electron, it becomes a positive ion

Ionic Compounds (M/NM)

Ionic Compounds (M/NM) Ionic (M/NM): e- are transferred from a metal to a nonmetal forming charged ions that stick together Cation: ion with a positive charge Anion: ion with a negative charge All ionic compounds are salts Soome common salts: potassium chloride, potassium iodide Cation Anion [Overall charge is neutral] Ionic compounds (salts) form a crystal lattice structure

Covalent Compounds (NM/NM) Covalent (NM/NM): e- are shared in overlapping molecular orbitals

Metallic Compounds Metallic (M/M): Delocalized e- move freely between atoms Metals are conductive (they conduct electricity) because their e- move. Metals are malleable/not rigid because their e- move.

Ionic Compounds (M/NM) Ionic (M/NM): e- are transferred from a metal to a nonmetal forming charged ions that stick together Cation: ion with a positive charge Anion: ion with a negative charge All ionic compounds are salts Soome common salts: potassium chloride, potassium iodide Cation Anion [Overall charge is neutral] Ionic compounds (salts) form a crystal lattice structure

Unknown Ionic Compounds Lab Identify the metal cation and nonmetal anion components of an unknown ionic compound M NM *Every group will receive a different set of unknown compounds to identify

Unknown Ionic Compounds Lab NM

Unknown Ionic Compounds Lab NM Flame Test

Unknown Ionic Compounds Lab NM Flame Test AgNO3 Test

Unknown Ionic Compounds Lab NM Flame Test HCl Test AgNO3 Test

Covalent Compounds (NM/NM) Covalent (NM/NM): e- are shared in overlapping molecular orbitals

Covalent Bonds Nonpolar Covalent Bond: e- are shared equally Polar Covalent Bond: e- are attracted to the atom with higher electronegativity Electronegativity increases as you move up and to the right on the periodic table.

Covalent Bonds Ex. Determine polar or nonpolar. Draw a dipole arrow towards the more negative side H-Cl O-H C-C O-F S-O + - Dipole arrow points towards the more electronegative atom. Electronegativity increases as you move up and to the right on the periodic table. H-C + O-H + C-C = O-F + S-O +

Polarity C - C C - Cl Na - Cl Difference in Electronegativity None High Nonpolar Covalent Polar Covalent Ionic C - C C - Cl Na - Cl

Bond Length & Bond Strength Atoms form single, double, or triple bonds depending on what’s needed to make an octet Longer/Weaker Shorter/Stronger

Resonance Resonance: more than one Lewis structure can be drawn because there’s an extra bond that moves between two or more locations The bonds have the same length and strength! They’re not like one single bond and one double bond, but more like two 1 ½ bonds Ex. SO2 CO32- Laboratory measurements show that all the bonds are equal in length, and between single and double bond length. None of these Lewis structures drawn are actually correct. The actual structure has equal bonds between single and double bond length and strength. The structures are changing so quickly that all we can measure is an average blur instead of being able to detect individual structures.

Lewis Structures Lewis Structures: drawing valence e- using dots Count the valence e- in the highest energy level (the outermost orbital) Covalent bonds are drawn using lines 2 e- in each bond! Both need one more e- to fulfill the octet rule They share each other’s e- Shared e- form a covalent bond Both chlorines have 7 val e-. They share electrons so they both get 8 val e- (octet rule). If there’s two chlorine atoms and they both have eight val e-, how many total val e- are there? 14, not 16 because they share two of them.

Lewis Structures Drawing Lewis structures: Count the total number of valence e- Draw the SKELETON Lewis structure Arrange the atoms with carbon in the middle Draw single bonds between atoms (2 e- in each bond) Satisfy octet/duet rule for each atom Draw 8 e- around each atom Confirm the total number of valence e- Compare to the total number of valence e- you counted to the total number you should have (from step 1) If there are too many e-, add double/triple bonds until they are equal If the molecule is an ion, enclose the structure in parentheses and put the charge at the top right on the outside of the parentheses.

Lewis Structures Ex: Draw the Lewis structure for O2

Lewis Structures Ex: Draw the Lewis structure for O2 6 + 6 = 12 valence e-

Lewis Structures Ex. Draw the Lewis structure for CH3I

.. Lewis Structures Ex. Draw the Lewis structure for CH3I 1 + 1 + 1 + 4 + 7 = 14 valence e-

Electron Pairs Bond Pairs: e- pairs in a bond (line) Lone Pairs: e- pairs not in a bond (dots) Lone Pairs Bonded Pairs

VSEPR Theory Valence Shell Electron Pair Repulsion Theory: the shape of a molecule is determined by the valence e- around the central atom Valence e- pairs (bond & lone) repel each other and spread out as far as possible Balloons: visual demo of orbitals repelling each other (spreading out).

Hint: atoms always want 4 pairs (8 e-) according to the octet rule! Molecular Shapes .. Linear, trigonal planar, and tetrahedral were the shapes of the balloons. Trigonal pyramidal is just like tetrahedral except one of the bonded pairs is invisible because it’s a lone pair. Bent is just like tetrahedral except two of the bonded pairs is invisible because they’re lone pairs. 2 bonded pairs (two double bonds) 2 bonded pairs 1 or 2 lone pairs 3 bonded pairs 1 lone pair 3 bonded pairs (one double bond) 4 bonded pairs Hint: atoms always want 4 pairs (8 e-) according to the octet rule!

Molecular Models Activity Clay = atoms Toothpicks = bonded pairs (2 e-) Thumbtack = lone pair (2 e-) Draw the lewis structures and write the name of the molecular shape. Then make the molecules in their correct shapes with clay, toothpicks, and thumbtacks.

NONPOLAR: there are no poles POLAR: there are poles Cut the molecule in half: If it has different sides (poles), it’s polar. If it’s the same all around, it’s nonpolar. NONPOLAR: there are no poles

Dipoles reinforce one another Molecular Polarity Molecular Polarity: determine whether bond dipoles cancel out or make + and - poles Dipoles reinforce one another Polar Molecule Dipoles cancel out Non-Polar Molecule Hint: if all the atoms around the central atom are the same, the molecule is non-polar

Lewis Structure, Shape, Polarity Molecule Lewis Structure Molecular Shape Molecular Polarity CH4 HCN SO2 NO3-

Lewis Structure, Shape, Polarity Molecule Lewis Structure Molecular Shape Molecular Polarity CH4 HCN SO2 NO3- 8 e- 10 e- 18 e- 24 e-

Lewis Structure, Shape, Polarity Molecule Lewis Structure Molecular Shape Molecular Polarity CH4 HCN SO2 NO3- 8 e- 10 e- 18 e- 24 e-

Lewis Structure, Shape, Polarity Molecule Lewis Structure Molecular Shape Molecular Polarity CH4 Tetrahedral HCN Linear SO2 Bent NO3- Trigonal Planar 8 e- 10 e- 18 e- 24 e-

Lewis Structure, Shape, Polarity Molecule Lewis Structure Molecular Shape Molecular Polarity CH4 Tetrahedral Non-Polar HCN Linear Polar SO2 Bent NO3- Trigonal Planar N/A 8 e- 10 e- 18 e- 24 e-