Intramolecular Forces Intermolecular Forces Bonding Intramolecular Forces Intermolecular Forces

Chemical Bond A force that holds groups of 2 or more atoms together to make them function as a unit

Why? Atoms form bonds to obtain a “noble gas configuration” Octet Rule- Atoms need to have 8 valence electrons

Types of Bonds Metallic Bonds: specific only to metals Ionic: results from transfer of electrons Covalent: results from sharing of electrons Coordinate covalent bond: when one atom provides both electrons of the shared pair

Metallic Bonding Metallic characteristics Occurs between like atoms of a metal in the free state Valence e- are mobile (move freely among all metal atoms) Positive ions in a sea of electrons Metallic characteristics High mp temps, ductile, malleable, shiny Hard substances Good conductors of heat and electricity as (s) and (l)

It is the mobile electrons that allow the metal to conduct electricity

Ionic Bonds Force of attraction between closely packed opposite ions Atoms that lose electrons easily (metals) Atoms that gain electrons easily (nonmetals) Electrons are transferred from metal to nonmetal Oppositely charged ions form and they attract each other.

Ionic Compounds The ions combine in quantities that provide the stable octet for each element involved They are electrically neutral, so the charges must balance out to zero, as we learned in our formula writing topic.

Properties of Ionic Compounds Hard Poor electrical conductors (solids) Good electrical conductors (in molten and solution) High melting and boiling points Dissolve in water Form crystalline solids

Ionic Lewis Diagrams Step 1 after checking that it is IONIC Step 2 Determine which atom will be the +ion Determine which atom will be the - ion Step 2 Write the symbol for the + ion first. NO DOTS Draw the e- dot diagram for the – ion COMPLETE outer shell Step 3 Enclose both in brackets and show each charge

Covalent In covalent bonding electrons are SHARED by the nuclei of each of the atoms present WHY? http://www.youtube.com/watch?v=1wpDicW_MQQ https://www.youtube.com/watch?v=LkAykOv1foc#t=29.851

Characteristics of Covalent Compounds Soft Poor conductors of electricity Poor heat conductors Low melting point High vapor pressures

Lewis Structures for Covalent

Covalent bonding: two types Nonpolar Results from EQUAL sharing Electronegativity difference 0 Polar Results from UNEQUAL sharing Note electronegativity difference between 0 and 1.7

Using Electronegativity to determine bond polarity H-H H-S H-Cl H-O H-F

Choose the More Polar Bond A) H-P or H-C B) O-F or O-I C) N-O or S-O D) N-H or Si-H

Polar Covalent Bonds Covalent bond, sharing electrons, But electron sharing not always equal. Fluorine pulls harder on the shared electrons than hydrogen does. Therefore, the fluorine end has more electron density than the hydrogen end.

Polar Covalent Bonds The greater the difference in electronegativity, the more polar is the bond.

Coordinate Covalent Bonding A situation in which one atom provides BOTH electrons that are in the bond.

Lewis Structures Diagrams for bonding in molecules Lines correspond to 2 electrons in bond Lewis structures are representations of molecules showing all valence electrons, bonding and nonbonding.

Lewis Structures Keep the following in mind: Must include all of the valence electrons from all of the atoms. The total number of electrons available in the structure is the sum of all the valence electrons from all of the atoms in a molecule Atoms that are bonded to each other share one or more pairs of electrons The electrons are arranged to fill the valence orbitals of that atom. This means 2 electrons for H, 8 for other elements

Lewis Structures A way to keep track of those valence electrons Find the sum of valence electrons of all atoms in the substance PCl3 5 + 3(7) = 26

Lewis Structures A way to keep track of those valence electrons The central atom is the least electronegative element that isn’t hydrogen (why?). Connect the outer atoms to it by single bonds. Keep track of the electrons: 26  6 = 20

Writing Lewis Structures Put eight electrons around the outer atoms (“fill their octet”) Keep track of the electrons: 26  6 = 20  18 = 2

Writing Lewis Structures Fill the octet of the central atom. Keep track of the electrons: 26  6 = 20  18 = 2  2 = 0

Writing Lewis Structures If you run out of electrons before the central atom has an octet… …form multiple bonds until it does.

Lewis Structures Try drawing the Lewis structure for each molecule: NF3 N2 CO2 SO3 ‐2 CCl4 O2 H2

Octet Rule Exceptions Hydrogen only requires 2 electrons BCl3 , B only has three electrons to share and therefore will only form three bonds Some circumstances provide P, S, Cl, Br and I to form molecules with more of their valence electrons PCl5, SF6

Resonance The bonding in some molecules or ions cannot be described by a single Lewis structure due to electron delocalization within multiple bonds. Resonance structure: an equivalent Lewis dot diagram placing a multiple bond between a different bond site in the molecule. Draw the Lewis structure for ozone, O3.

- Resonance Draw the Lewis structure for ozone, O3. But why should one O be different from the other? + -

Resonance One Lewis structure cannot accurately depict a molecule such as ozone. We use multiple structures, resonance structures, to describe the molecule. - + + -

Resonance It is at odds with the true, observed structure of ozone, …both O—O bonds are the same length.

Resonance Structures Draw Resonance Structures for CO32- Draw Resonance Structures for NO3-

Fewer Than Eight Electrons The lesson is: If filling the octet of the central atom results in a negative charge on the central atom and a positive charge on the more electronegative outer atom, don’t fill the octet of the central atom. + - - - + +

VSEPR Theory and Shapes Valence Shell Electron Pair Repulsion Theory Three dimensional shapes of molecules can be predicted using Lewis structures Shapes are determined by considering groups of electrons around central atom, both bonded and unbonded (lone pairs) Electron groups are arranged to minimize the repulsion between their negative charges Focus on central atom, number of bonds and number of lone pairs

VSEPR Theory and Shapes Central atoms with two electron groups: minimal repulsion when electron groups are on opposite sides of central atom LINEAR

VSEPR Theory and Shapes Central atoms with three electron groups Example: H2CO Minimal repulsion occurs when three electron groups are as far apart as possible- TRIGONAL PLANAR

VSEPR Theory and Shapes SO2 there are three electron groups, but one group is a lone pair, so the shape is bent

VSEPR Theory and Shapes Four electron groups When all of the electron groups are bonded, the shape is tetrahedral

VSEPR Theory and Shapes Four electron groups with one lone pair- trigonal planar

VSEPR Theory and Shapes Four electron groups, two lone pairs

VSEPR Theory and Shapes

VSEPR Shapes

Summary https://www.youtube.com/watch?v=KjoQHqgzda8 https://www.youtube.com/watch?v=wYZg1j7o2x4&ebc=ANyPxKryVMXWuhz6cg_UxEfwZHCa7yILX5elCuThuVg2h45P6bPR-UTe_bw3tO35eWdvZu7JSpJf7TmC3o129mvZ33r3FQ5wCA