Chapters 4 & 5 Chemical Bonding Valence Electrons Outermost
electrons s and p electrons for main group elements Responsible for
chemical properties of atoms Participate in chemical reactions Core
Electrons Valence Electron Problems Write out the electron
configurations for the following elements and identify how many
core and valence electrons each has. 1)Mg 2)S 3)Br 4)Kr Lewis Dot
Structures LDS: a representation of an atom using its chemical
symbol surrounded by dots that signify valence electrons Problems
Write the Lewis Dot Structures for the following atoms Li Be Br C N
Ne Li: [He]1s 1 Na: [Ne]2s 1 K: [Ar]3s 1 Octet Rule Octet Rule: the
tendency for atoms to seek 8 electrons in their outer shells
Natural electron configuration of the Noble Gases Done by gaining,
losing, or sharing electrons Increases stability H and He seek a
Duet Ionic Bonding Ions: atoms that have a charge due to gain or
loss of electrons Anion: (-) charged atom Cation: (+) charged atom
Ionic Bond: a bond formed through the transfer of one or more
electrons from one atom or group of atoms to another atom or group
of atoms Formula Unit Ionic Compounds: compounds composed of
oppositely charged ions that are held together by their attraction
to each other Metal + Non-metal NaCl Metal + Polyatomic Ion NaNO 3
Polyatomic Ion + Non-metal NH 4 Cl Polyatomic Ion + Polyatomic Ion
NH 4 NO 3 Net charge on compound equal to zero Oxyanions SO 4 2-
Sulfate SO 3 2- Sulfite PO 4 3- Phosphate PO 3 3- Phosphite NO 3 -
Nitrate NO 2 - Nitrite ClO 4 - Perchlorate ClO 3 - Chlorate ClO 2 -
Chlorite ClO - Hypochlorite Rules For Naming Ionic Compounds 1)Name
the cation by its elemental/polyatomic name 2)If the metal is a
transition metal with a variable charge, indicate its charge with a
Roman Numeral in parentheses 3)Next, name the anion and change its
ending to -ide 4)If the anion is polyatomic, do not change the
ending to -ide 5)Do NOT use prefixes (mono, di, tri etc.) to
indicate how many of each atom are present Problems Write the name
for the following compounds: 1)KI 2)MgBr 2 3)Al 2 O 3 4)FeCl 2
5)CaSO 4 6)Ba(NO 2 ) 2 7)Cu(NO 3 ) 2 Write the Formula for the
following ionic compounds: 8)Sodium Fluoride 9)Calcium Sulfite
10)Calcium Chloride 11)Iron (III) Oxide 12)Cobalt (II) Hydroxide
13)Ammonium Bromide 14)Ammonium Carbonate 15)Aluminum Carbonate
Iron (II) Chloride Iron (III) Chloride Covalent Compounds Covalent
Compounds: compounds composed of atoms bonded to each other through
the sharing of electrons Electrons NOT transferred No + or charges
on atoms Non-metal + Non-metal Also called molecules Examples : H 2
O CO 2 Cl 2 CH 4 or H-H or Duet Naming Covalent Compounds 1)Name
the first non-metal by its elemental name 2)Add a prefix to
indicate how many 3)Name the 2 nd non-metal and change its ending
to -ide 4)Add a prefix to indicate how many Problems Write the name
of the following compounds: 1)CO 2)NI 3 3)N 2 O 4)SF 6 5)B 2 O 3
Write the formula for the following compounds: 6)Phosphorous
Pentachloride 7)Nitrogen Monoxide 8)Dinitrogen Tetroxide
9)Tetraphosphorous Decoxide Problems 1)KCl 2)Na 2 S 3)H 2 O 4)SO 2
5)K 3 PO 4 6)FeCl 3 7)(NH 4 ) 2 SO 4 8)SCl 2 9)Cu(OH) 2 10)P 2 O 5
8)Sodium Iodide 9)Aluminum Sulfate 10)Phosphorous Pentabromide
11)Magnesium Nitride Naming Acids Acids that do not contain oxygen
1)Begin the name with hydro 2)Name the anion, but change the ending
to -ic 3)Add acid on the end HCl HF Acids that contain oxygen 1)Do
not put hydro at the beginning 2)Begin the name with the anion 3)If
the anion has the ending -ate, change this to -ic acid 4)If the
anion has the ending -ite, change this to -ous acid HClO 4 HClO 3
HClO 2 HClO Problems Name the following 1)HBr(g) 2)HBr(aq) 3)HNO 2
(aq) 4)HNO 3 (aq) 5)HI (aq) 6)HI (g) 7)H 2 CO 3 (aq) 8)H 3 PO 4
(aq) 9)H 3 PO 3 (aq) 10)HCN (aq) Molecular Structures Water Ball
& Stick ModelsSpace-Filling Models Methane Ethanol Lewis Dot
Structures 1)Count the total number of valence electrons in the
molecule. Ex: PCl 3 2)Use atomic symbols to draw a proposed
structure with shared pairs of electrons. Atoms dont tend to bond
to other atoms of the same element when they can avoid it
Exception: Carbon 3)Place lone pair electrons around each (except
H) to satisfy the octet rule, beginning with the terminal atoms
4)Place any leftover electrons on the central atom 5)If the number
of electrons around the central atom is less than 8, change single
bonds to the central atom to multiple bonds (double or triple). Ex:
CH 2 O Problems Draw the LDSs for the following molecules: 1)Cl 2 O
2)C 2 H 4 3)C 2 H 6 O What Things Like To Do Halogens Like to be
terminal Like to have one single bond and 3 lone pairs (non-bonding
electrons) Carbon Likes to have 4 single bonds and no lone pairs A
double bond counts as two singles A triple bond counts as three
singles Likes to be central Likes to bond to other carbons Silicon
Likes to do what carbon does Oxygen Likes to have two single bonds
and 2 lone pairs Sulfur Likes to do what oxygen does Nitrogen Likes
to have 3 single bonds and one lone pair Phosphorous Likes to do
what nitrogen does Hydrogen Likes to be terminal with only one
single bond No lone pairs! Problems 1)SH 2 2)C 3 H 8 3)Si 2 H 6
4)PI 3 5)CH 3 OH 6)C 2 H 2 7)CCl 2 O 8)N 2 H 4 9)CH 2 OS 10)C 2 H 6
O 11)CO 12)BrHO Electronegativity The measure of the ability of an
atom to attract electrons to itself Increases across period (left
to right) and Decreases down group (top to bottom) fluorine is the
most electronegative element francium is the least electronegative
element Electronegativity Scale Types of Bonding 1)Non-Polar
Covalent Bond: Difference in electronegativity values of atoms is
0.0 0.4 Electrons in molecule are equally shared Examples: Cl 2, H
2, CH 4 EN Cl = = 0 Pure Covalent 2)Polar Covalent Bond: Difference
in electronegativity values of atoms is 0.4 2.0 Electrons in the
molecule are not equally shared The atom with the higher EN value
pulls the electron cloud towards itself Partial charges Examples:
HCl, ClF, NO EN Cl = 3.0 EN H = 2.1 = 0.9 Polar Covalent 3)Ionic
Bond: Difference in EN above 2.0 Complete transfer of electron(s)
Whole charges EN Cl = 3.0 EN Na = 0.9 = 2.1 Ionic Problems Predict
the type of bonding in the following compounds using differences in
EN values of the atoms. Indicate the direction of the dipole moment
if applicable 1)KBr 2)HF 3)BrI 4)FI Valence Shell Electron Pair
Repulsion Theory VSEPR theory: Electrons repel each other Electrons
arrange in a molecule themselves so as to be as far apart as
possible Minimize repulsion Determines molecular geometry Defining
Molecular Shape Electron pair geometry: the geometrical arrangement
of electron groups around a central atom Look at all bonding and
non-bonding e - s Molecular Geometry: the geometrical arrangement
of atoms around a central atom Ignore lone pair electrons 2 e -
groups surrounding the central atom e- pair geometry: linear MG:
linear AXE designation: AX 2 E 0 A: Central Atom X: Bonding pairs
E: Non-bonding pairs Example: BeCl 2 3 e - groups 3 Bonds, 0 Lone
Pairs e - PG: Trigonal Planar (Triangular planar) MG: Trigonal
Planar AX 3 E 0 BF 3 2 Bonds, 1 Lone Pair e - PG: Trigonal Planar
(Triangular planar) MG: Bent/angular AX 2 E 1 GeCl 2 4 e - groups 4
bonds, 0 Lone Pairs e - PG: Tetrahedral MG: Tetrahedral AX 4 E 0 CH
4 3 bonds, 1 Lone Pair e - PG: Tetrahedral MG: Triangular Pyramidal
AX 3 E 1 NH 3 2 bonds, 2 Lone Pairs e - PG: Tetrahedral MG:
Bent/Angular AX 2 E 2 H 2 O Drawing LDS With Correct Geometry
Molecular Polarity Problems 1)NF 3 2)CH 2 O 3)CBr 4 4)CHCl 3 5)CH 2
Cl 2 Draw the Lewis Dot Structures for the following molecules and
then identify the direction of polarity, if any.
