How do you think H and O in water are bonded together? Draw a picture to help your explanation.

6.1 Introduction to Chemical Bonding Ch. 6 Bonding 6.1 Introduction to Chemical Bonding

Chemical Bonds atoms rarely exist alone when atoms are bonded together, they have less potential energy and are more stable What is potential energy? chemical bond – mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together

Ionic Bonds results from electrical attraction between large numbers of cations and anions atoms donate or accept electrons from each other

Covalent Bonds results from sharing of electron pairs between two atoms the electrons shared belong to both atoms

Ionic vs. Covalent

Ionic vs. Covalent bonding usually does not fall in one category or the other, but somewhere in between type of bond depends on the elements differences in electronegativities 0.3

Practice Determine whether each of the following bonds will be: Ionic or covalent

Patterns What kind of patterns do you see? metals + nonmetals = ionic nonmetals + nonmetals = covalent

Ch. 6 Bonding 6.2 Covalent Bonding

What is the difference between covalent and ionic bond? How do you determine which a compound contains?

Molecular Compounds molecule: neutral group of atoms held together by covalent bonds molecular compound: compound whose simplest unit is a molecule

Formulas chemical formula: tells the number of each type of atom in a compound molecular formula: tells the number of each type of atom in a molecular compound ex. H2O, Cl2, C6H12O2

Molecular Compounds diatomic molecule: a molecules containing only 2 atoms usually refers to 2 of the same atoms ex: O2, Br2, F2, etc. 7+1 rule

Formation of Covalent Bond

Formation of Covalent Bond two nuclei and two electron clouds repel each other creating an increase in PE approaching nuclei and electron clouds are attracted to each other to create a decrease in PE

Formation of Covalent Bond a distance between the nuclei is reached in which repulsion and attraction forces are equal potential energy is at the lowest point possible at the bottom of the curve on PE graph

Covalent Bonds Bond Length distance between two bonded atoms at their lowest PE average distance since there are some vibrations measured in pm (1012 pm = 1 m) stronger the bond, shorter the bond

Covalent Bonds Bond Energy energy is released when atoms become because they have lower PE the same amount of energy must be used to break the bond and form neutral isolated atoms stronger bond, higher bond energy average since varies a small amount based on atoms in entire molecule in kJ/mol

Which elements naturally exist as diatomic molecules? Remember, the 7 + 1 rule How many valence electrons do each of the halogens have?

Octet Rule representative elements can “fill” their outer energy level by sharing electrons in covalent bonds Octet Rule- a compound tends to form so that each atom has an octet (8) of electrons in its highest energy level by gaining, losing or sharing electrons Duet Rule- applies to H and He

Octet Rule Less than 8: More than 8: Boron: 6 in outer energy level anything in 3rd period or heavier because may use the empty d orbital ex: S, P, I

Electron Dot Diagrams a way to show electron configuration identifies the number and pairing of valence electrons to show how bonding will occur write the noble gas notation identify the number of valence identify how many are paired and how many are alone do not go by Figure 6-10

N Example Nitrogen Sulfur 1s2 2s2 2p3 5 valence 2 are paired 3 are alone Sulfur 1s2 2s2 2p6 3s2 3p4 6 valence 4 paired (2 pairs) 2 are alone N

Lewis Structures like dot diagrams but for entire molecules atomic symbols represent nucleus and core electrons and dots or dashes represent valence electrons unshared electrons: (lone pairs) pair of electrons not involved in bonding written around only one symbol bonding electrons: written in between 2 atoms as a dash

Types of Bonds single- sharing of one pair of electrons weakest, longest double- sharing of 2 pairs of electrons stronger and shorter triple- sharing of 3 pairs of electrons strongest and shortest multiple bonds include double and triple bonds

Drawing Lewis Structures find the number of valence electrons in each atom and add them up draw the atoms next to each other in the way they will bond add one bonding pair between each connected atoms add the rest of the electrons until all have 8 (consider exceptions to octet rule)

H H C Cl CH3Cl Example 1 methyl chloride C: 4 x 1 = 4 H: 1 x 3 = 3 Cl: 7 x 1 = 7 total = 14 electrons carbon is central H H C Cl duet octet duet octet duet H H C Cl

H N H H NH3 Example 2 ammonia N: 5 x 1 = 5 H: 1 x 3 = 3 total = 8 N is central Example 2 H N H H

Example 3 N2 nitrogen gas N: 5 x 2 = 10 10 electrons N N N N

H C H O Example 4 CH2O formaldehyde C: 4 x 1 = 4 H: 1 x 2 = 2 O: 1 x 6 = 6 total = 12 C is central H C H O

O O O O O O Example 5 O3 ozone O: 6 x 3 = 18 two completely equal arrangements the real structure is an average of these two where each bond is sharing 3 electrons instead of 4 or 2 O O O O O O

O O O O O O Resonance Structures resonance – bonding between atoms that cannot be represented in on Lewis structure show all possible structures with double-ended arrow in between to show that electrons are delocalized O O O O O O

NO31- N: 5 x 1 = 5 O: 6 x 3 = 18 total = 23 + 1 = 24 Example 6

Covalent Network Bonding a different type of covalent bonding not specific molecules lots of nonmetal atoms covalently bonded together in a network in all directions example: diamond silicon dioxide graphite

Ch. 6 Bonding 6.3 Ionic Bonding

Ionic Compounds ionic bonds do NOT form molecules chemical formulas for ionic compounds represent the simplest ratio of ion types made of anions and cations

Ionic Compounds combined so that amount of positive and negative charge is equal usually crystalline solid formula of ionic compound depends of the charges of the ions combined

Formation attractive forces: repulsive forces: oppositely charged ions nuclei and electron clouds of adjacent ions repulsive forces: like-charged ions electrons of adjacent ions

Formation distance between the ions creates a balance between those forces ions minimize their PE by combining in an orderly arrangement called a crystal lattice

Formation specific lattice pattern created depends on: charges of ions size of ions Calcium Bromide: each Ca2+ is surrounded by 8 F- each F- is surrounded by 4 Ca2+ Sodium Chloride each Na+ is surrounded by 6 Cl- each Cl- is surrounded by 6 Na+

Ionic vs. Molecular ionic bonds and molecular bonds are both strong ionic bonds connect all ions together molecules are more easily pulled apart because intermolecular forces are weak

Ionic vs. Molecular Molecular Compounds: low melting and boiling points many are gases at room temperature Because the intermolecular forces of the molecules are weak so they are easily separated

Ionic vs. Molecular Ionic Compounds: higher melting and boiling points all are solid at room temperature hard: Because of the strong forces, it is difficult for one layer of ions to move past another brittle: if one layer is moved, the layers come apart completely

Ionic vs. Molecular Ionic Compounds: good conductors in liquid state Because ions are free to move and carry charge poor conductor in solid state Because ions are fixed in place

Polyatomic Ions charged group of covalently bonded atoms Example: CN-

NH4+ : ammonium ion

O O S O O O H SO42- : sulfate ion OH- : hydroxide ion 5 x 6 = 30 total = 30 + 2 = 32 OH- : hydroxide ion 6 + 1 + 1 = 8 total O S O O O H

Draw the Lewis Structures for NO21- and PO43-

Ch. 6 Bonding Metallic Bonding

Bonding of Metals the highest energy level for most metal atoms does not contain many electrons usually have empty p and d block these vacant overlapping orbitals allow outer electrons to roam freely around the entire metal the electrons are delocalized – are not with one specific atom

Bonding of Metals these roaming electrons form a sea of electrons around the metal atoms metal atoms are packed in a crystal lattice metallic bonding – bonding that results from the attraction between metals atoms and sea of electrons

Properties of Metals conductivity luster (shininess) from the freedom of electrons to move around the atoms luster (shininess) contain many orbitals with only small differences in energy many amounts of energy can be absorbed and emitted

Properties of Metals malleability and ductility bonding is the same in every direction one layer of atoms can slide past another without friction

Bond Strength depends on the nuclear charge (Z) or the number of protons depends on the number of electrons in the “sea” heat of sublimination – amount of heat required to turn solid, bonded metal atoms into gaseous individual atoms

Metallic vs. Ionic