Chemical Bonds: The Formation of Compounds from Atoms Dr. Bixler-Zalesinsky

PERIODIC TRENDS IN ATOMIC PROPERTIES 11.1

Metals and Nonmetals (review)

Ionization energy is the energy required to remove an electron; corresponds to their charge

Atomic Radii increase going down a group and decrease across a period

LEWIS DOT DIAGRAMS Valence Electrons

Valence Electrons & Per. Table

Lewis Structures of an atom shows the valence electrons (ones involved in bonding)

Octet Rule Every atom aspires to have eight electrons in its outermost shell (2 s electrons and 6 p electrons just like the noble gases) They must borrow (covalent molecules), release or accept (ionic compounds) electrons to get to the eight.

Types of bonding

BONDING Ions and

Ionic bonding occurs between metals (cations) and nonmetals (anions)

The nonmetal accepts the electron(s) and the metal donates the electron(s) ionic bond is the attractive between oppositely charged ions

Form large crystals; our formulas are the smallest whole number ratios not the true number of atoms

The charges must cancel out and equal zero to form stable compounds If you have a +2 ion then you need either two -1 ions or one -2 ion

Link to Video clip on Ionic Bonding (1:39)

COVALENT V. IONIC Bonding

11.5 Covalent Bonding: Sharing Electrons Covalent bonding occurs between two nonmetal atoms Electrons are shared between two atoms

11.7 Lewis Structures of Molecules 1.Find the number of valence electrons for each element in the structure 2.Multiply the number of valence electrons times the number of atoms you have of that element 3.Determine which element can make the most bonds and put it in the center and attach the other elements to it 4.Make each atom have 8 valence electrons around it. 5.Add up the number of electrons you used in the structure. This number must match the total number of electrons you started with

H 2.1 Li 1.0 Na 1.5 He B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Ne Na 0.9 Mg 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 Ar K 08 Ca 1.0 Br 2.8 Kr Rb 0.8 Sr 1.0 Se 2.4 Cs 0.7 Ba 0.9 I 2.5 Fr 0.7 Ra 0.9 Xe Rn Number of Valence Electrons: (number of bonds each can make) S S/D/T S/T S/D S 0 (types of bonds s=single, D= double, T= triple) Page 151 in textbook

Ex. Write the Lewis Dot Diagrams for the following molecules I 2 H 2 O FCl CF 4 NBr 3

Molecule (Covalent) Nomenclature Naming: These binary inorganic molecules are named using prefixes like mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, deca-, The first element only gets a prefix if there is more than one of it; otherwise, the element name remains the same. The second element ALWAYS gets a prefix and the ending changes to –ide. Ex. CO is carbon monoxide (two words not capitalized)

HW p. 164 # 2 a – f Write the question and answer!

Multiple Bonds Double and Triple Bonds Knowing when NOT to use them is as important as understanding when to use them!

Multiple Bonds Some times using the correct number of electrons will not give you a full octet. When this happens: 1 st double check your math and counting 2 nd see if the atoms involved can make a double or triple bond

Double Bonds O, S, and C can make double bonds with each other than themselves but no others! Double bond is 4 electrons in a bond Symbolized by an = sign Take a look at CO 2

Triple Bonds P, N, and C can make a triple bond with each other or themselves A triple bond is 6 electrons in a bond The symbol for a triple bond is = Let’s try N 2

Molecular Geometry VSEPR Theory and Application

Structural Formula Shows how elements of a molecule are connected to each other

VSEPR V = valence S = shell E = electron P = pair R = repulsion Electrons will arrange themselves as far apart from one another as possible Unbonded pairs take up more room than bonded ones

3-D Hybridized orbitals, shapes, and decision tree

Linear Shape

Bent Shape

Trigonal planar

Pyramidal

Tetrahedral

VSEPR Video Review (3:21)

sp hybridization

Path to hybridization

Sp3 hybridization

HYBRIDIZATION Video Review (1:36)

VSEPR Theory of Molecular Geometry # of atoms Central Atom ShapeBond angle Example

VSEPR Theory of Molecular Geometry # of atoms Central atom Shape Bond Angle (in degrees) Example 2NoneLinear180HF 3AnyLinear180CO 2 3S or OBent105H2OH2O 4BTrig. Planar 120BCl 3 4 (3-D)P or N Pyramidal 107NH 3 5 (3-D)C or Si Tetrahedral 109.5CH 4

Shape Decision Tree How many atoms?2 = Linear 3 No unshared pairs = Linear 2 unbonded pairs = bent 4 No unshared pairs = Trigonal Planar 1 unbonded pair = Pyramidal 5 = Tetrahedral

Polar Covalent v. Nonpolar Covalent

Polar and Nonpolar Covalent Bonds If they are shared equally they are said to be nonpolar bonds if they are not equally shared they are said to be polar bonds Sharing of electrons has to do with the pull of one element compared to the other element sharing the electron pair. This pulling is called electronegativity (eneg) which increases across the period and up the group The larger the electronegativity the greater the time the electrons spend with the more electronegative atom giving it a slightly positive charge and because of this imbalance we call this a polar molecule If the eneg difference lies between 0.5 to 1.6 it is a polar bond

Polar or Nonpolar Molecules 1.Determine the shape of the molecule 2. Determine how many polar bonds there are in the molecule 3.If there are NO polar bonds the molecule must be NONpolar. 4.If there is exactly one polar bond, the molecule is polar. 5.If there is more than one polar bond, follow the chart below.

Molecules with more than one polar bond (assuming polarity is equal) Shape# of polar bondsMolecular polarity Linear2Nonpolar Bent2Polar Trig Planar2Polar Trig Planar3Nonpolar Pyramidal2Polar Pyramidal3Polar Tetrahedral2Nonpolar Tetrahedral3Polar Tetrahedral4Nonpolar

VSEPR Theory of Molecular Geometry # of atoms Central Atom ShapeBond angle Example

Polarity of Molecules with more than one polar bond Shape# of polar bondsMolecular Polarity