Name _________________________
Block _________
Unit 4 Bonding Notes 
Atoms are generally found in nature in combination held together by _______________ ____________ . 
A chemical bond is a _______________________________ between the nuclei and valence electrons of different atoms that binds the atoms together 
There are three main types of bonding: __________, ___________ and __________. 
Ionic Bonding occurs between a ______________ and a ______________. 
Metallic bonding occurs between two ____________________. 
Covalent bonding occurs between a ______________ and a ______________. 
A positive ion is called a _____________________________. 
A negative ion is called an_____________________________. What determines the type of bond that forms? 
The ______________________ of the two atoms involved are redistributed to the ___________________________________________. 
The interaction and rearrangement of the _______________________ determines which type of bond that forms. 
Before bonding the atoms are at their highest possible _______________________. There are __________ philosophies of atom to atom interaction. 
One understanding of the formation of a chemical bond deals with balancing the opposing forces of _____________________________________. ‐

Repulsion occurs between the _________________ clouds of each atom. 
Attraction occurs between the __________________ and the negative electron clouds. 
When two atoms approach each other closely enough for their electron clouds to begin to overlap 
The electrons of one atom begin to _______________ the electrons of the other atom 1

And repulsion occurs between the _______________ of the two atoms. 
As the optimum distance is achieved that balances these forces, there is a release of potential energy 
The atoms _________________ within the window of maximum attraction/minimum repulsion 
The _______________ released the stronger the _______________ bond between the atoms 
Another understanding of the formation of a chemical bond between two atoms centers on achieving the most __________________________ of the atoms’ valence electrons 
By rearranging the electrons so that each atom achieves a ______________________ ________________ ________________ creates a pair of stable atoms (____________ ___________________________________) 
Sometimes to establish this arrangement one or more valence electrons are ________ ______________________________________________. 
Basis for _________________ bonding 
Sometimes valence electrons are _______________________________________ 
Basis for _____________________ bonding 
A good predictor for which type of bonding will develop between a set of atoms is the difference in their __________________________________________.. 
Remember, electro‐negativity is a measure of the ______________________ an atom has for e‐s after developing a bond 2

The _________________________ the difference between the two atoms, __________ ___________________________________________________ of electrons 
Let’s consider the compound Cesium Fluoride, _______________. 
The electro‐negativity value (EV) for Cs is ____________; the EV for F is ___________. 
The difference between the two is _____________, which falls within the scale of ionic character. 
When the electro‐negativity difference between two atoms is greater than ______________________________________. The take home lesson on electro‐negativity and bonding is this: 3
‐

The ________________________ are on the P.T., the more evenly their e‐ interact, and are therefore more likely to form a _______________________ ‐

The ________________________________ are on the P.T., the less evenly their e interact, and are therefore more likely to form an ionic bond. 
Forming ionic bonds can be represented as _____________________________. 
Hint - metal w/nonmetal = ionic
nonmetal w/nonmetal = covalent
Covalent Bonding 

In a Covalent bond: The ___________________________ between the atoms involved is not extreme 
So the interaction between the involved electrons is ____________________________ 
It may not be an equal sharing relationship, but at least the electrons are being “___________________”. 
Covalent bonds are between _____________________________________________. 
Covalent bonds are formed when electrons are _____________________ between two atoms. If two atoms share 2 electrons, they form a _____________________________. If two 
atoms share 4 electrons, they form a _____________________________. If two atoms share 6 electrons, they form a ____________________________. 
There are two types of covalent bonds: polar and non‐polar. Polar bonds usually involve __________________________________. Non‐polar bonds usually involve ___________________________________. 
In polar bonds, the electrons are shared ________________________________. In non‐
polar bonds, the electrons are shared ________________________________. 4

Covalent compounds can exist in any state (solid, liquid or gas). They have _____________ melting and boiling points. 
A ____________________ is a diagram showing the arrangement of valence electrons among the atoms in a molecule. 
We use ______________ to represent valence electrons involved in bonding. 
______________________ are electrons present in the outermost energy level of an atom. Examples of Lewis Dot Symbol 5

Recall that the valence electrons for the elements can be determined based on the elements position on the ________________. 
Atoms can _________________ more than one electron pair. 
They may double or triple up pairs of electrons to satisfy the _________________. 
A _______________________ is the sharing of one pair of electrons between two atoms. 
A __________________________ is the sharing of ______ pairs of electrons between two atoms. 
A ____________________ is the sharing of _______ pairs of electrons between two atoms. 
_____________ element can form depends on the number of valence electrons. Rules for Writing Lewis Structures 
Determine whether the compound is __________________________. If covalent, treat the entire molecule. If ionic, treat each ion separately. 
Determine the __________ number of _______________ _________________ available to the molecule or ion by summing the valence electrons of all the atoms in the unit. 
Organize the atoms so there is _______________ ____________________ (usually the least electronegative) surrounded by ligand (outer) atoms. Hydrogen is never the central atom. 
So what is the bottom line? To be stable the two atoms involved in the covalent bond share their electrons in order to achieve the arrangement of a ________________. 6
Ionic Bonds 
In an Ionic bond: 
The electro‐negativity difference is ____________________, 
So the atom with the _______________________ pull doesn’t really share the electron 
Instead the electron is essentially ___________________________________________ _______________________________________________________________________ 
When a metal bonds with a nonmetal, an bond is formed. 
An ionic bond always involves the TRANSFER of electrons from the to the . 
The cation and anion are held together by ___________________________________. 
An ionic compound does not consist of individual molecules. Instead, there is a huge network of positive and negative ions that are packed together in a 
. Because their bonds are so strong, ionic compounds tend to have very 
_________________ points. Ionic compounds are _____________________________, which means they can conduct electricity. 
When forming ionic compounds the positive and negative charges must _____________. 
_______________________________________________________________________ _______________________________________________________________________ . 
An electron is _______________________ from the sodium atom to the chlorine atom. 
The bottom line is, to be stable the two atoms involved in the ionic bond will either ___________________ or _________________ their valence electrons in order to achieve a stable __________________ arrangement of electrons. 7
Metallic Bonding 
Metallic bonds consist of positively charged metallic cations that donate electrons to the __________________. 
The “sea” of electrons are shared by all atoms and can move throughout the structure. 
Properties: o
o
o
o
Thermal Conductivity Electrical Conductivity Malleability‐ the ability to be hammered down into thin sheets. Ductility‐ the ability to be drawn into a wire. 
In a metallic bond:

The resulting bond is a cross between ____________ and _____________ bonding

Valence electrons are transferred from one metal atom to the __________________ metal
atoms

But none of the involved metal atoms want the electrons from the original atom, nor their
own so they ___________________________________

What results is a ________________________________ of valence electrons that none of the atoms in the collection ________________ the valence electrons 
It resembles collection of positive ions floating around in a sea of electrons Determining the Type of Bond: Determine the type of bond (Ionic, Covalent or Metallic) in the following compounds: Compound NaCl CO FeNi SiS2 Bond Type Compound NCl3 PF3 CaCl2 Fe2O3 Bond Type End of PPT 1
8
Draw the Lewis structures for the following compounds:
1) PBr3
4)
NO2-1
2) N2H2
5)
C2H4
3) CH3OH
5)
HBr
Determine if the elements in the following compounds are metals or non-metals. Describe
the type of bonding that occurs in the compound.
Compound
NO2
Element 1
(metal or non-metal?)
N = non-metal
Element 2
(metal or non-metal?)
O = non-metal
Bond Type
covalent
NaCl
SO2
PO43MgBr2
CaO
H2O
Cu-Zn alloy
9
Rules of writing formulas:




· positive ion is written first … this is usually a metal
· negative ion is written second … this is usually a nonmetal
· subscripts are used to show how many ions of each part are in the compound. They
are used to balance the charge of the ions.
· criss-cross method:
Examples:
1.
sodium oxide
sodium is the positive ion = +1
oxide is the negative ion = -2
therefore … it takes 2 sodium ions to balance the charge of the oxide
Formula = Na2O
2.
calcium nitrate
calcium is the positive ion = +2
nitrate is the negative ion = -1
therefore … it takes 2 nitrates to balance the charge of calcium
Formula = Ca(NO3)2
3.
aluminum sulfide
aluminum is the positive ion = +3
sulfide is the negative ion = -2
therefore … it takes 2 aluminum ions and 3 sulfide to balance the charge
Formula = Al2S3
10
Covalent Naming  Binary covalent compounds are characterized by having two nonmetals. Naming these compounds involves the use of numerical prefixes: Prefix Number Prefix Number 1 6 2 7 3 8 4 9 5 10  If there is only ONE atom of the first element, you DON’T need a prefix. The FIRST element is named as a normal element. The SECOND element has an –IDE ending.  o N2O4 o XeF4 Diarsenic pentoxide o Phosphorous pentabromide o Carbon tetraiodide o N2O5 o Trisilicon tetranitride o CO o Tetraphosphorous decoxide o CBr4
11
Write the correct formulas for each covalent compound: Compound Name Oxidation States water O (‐2) H (+1) Carbon Dioxide C (+4) O (‐2) Chlorine (Diatomic Element) Cl (‐1)
Methane (5 total atoms) C (‐4) H (+1) N (‐3) H (+1) Ammonia (4 total atoms) Carbon tetrabromide (5 total atoms) C (+4) Br (‐1) Phosphorous trichloride (4 total atoms) P (‐3) Cl (‐1) Diphosphorous trioxide (5 total atoms) P (‐3) O (‐2) Covalent Formula Balancing Charges Criss‐Cross rule 1. Write out symbols and charge of elements 2. Criss‐Cross charges as subscripts (Swap and Drop) 3. Combine as a formula unit Equation Form of Balancing Charges (Number of Cations)x(Cation Charge) + (Number of Anions)x(Anion Charge) = 0 (1)(+3) + (X)(‐1) = 0, x = 3 o EX: Aluminum and Oxygen EX: Barium and Oxygen 12

Balancing Charges Practice: o Lithium Iodide o Strontium Chloride o Sodium Sulfide In each box, write the formula of the ionic compound consisting of the positive ion to the left of the box and the negative ion above the box. -2
-3
-3
Cl
S
F
N
O
P
+2
Mg
Cs+
Cr+3
Na
Zn+2
Al+3
K
I have provided you with charges on some but not all the elements above. 13
Valence Shell Electron Pair Repulsion Theory (VSEPR Theory). This is the way that we predict the geometry shape of molecules, A model was developed a qualitative model known as Valence Shell Electron Pair Repulsion Theory (__________________ Theory). The basic assumptions of this theory are summarized below. 1) The electron pairs in the valence shell around the central atom of a molecule repel each other and tend to orient in space so as to minimize the repulsions and maximize the distance between them. 2) There are two types of valence shell electron pairs: ______________ pairs and ________________ pairs Bond pairs are ______________ _________________ by two atoms and are attracted by two nuclei. Hence they occupy less space and cause less repulsion. Lone pairs ________________ ______________ involved in bond formation and are in attraction with only one nucleus. Hence they occupy more space. As a result, the lone pairs cause more repulsion. Note: The bond pairs are usually represented by a ___________ _______________, whereas the lone pairs are represented by a lobe with two electrons. 3) In VSEPR theory, the ___________________ bonds are treated as if they were single bonds. The electron pairs in multiple bonds are treated collectively as a single super pair. 4) The shape of a molecule can be predicted from the number and type of valence shell electron pairs around the central atom. When the valence shell of central atom contains only bond pairs, the molecule assumes symmetrical geometry due to even repulsions between them. 14
Bonds VSPER‐ Valence Shell Electron Pair Repulsion Lone Pairs Shape Linear Bent Trigonal Planar Trigonal Pyramidal Tetrahedral Steric number is the total number of ______________________________ AND ______________________. 15
VSEPR Practice Complete the table with the requested information. Molecule Structural Diagram Oxidation State of each element CClF3 SF2 BF3 SiBr4 NH3 Molecular Geometry 16
Bonding Notes Activity:
Bond Type
Formula
(Ionic/Covalent)
Lewis
Structure
Shared
Lone
Total
Electron Electron Electron Shape
Pairs
Pairs
Pairs
Structural
Formula
NaCl
NH3
SCl2
CH4
PCl3
MgCl2
CO2
SiBr4
H2O
BCl3
17
Draw the structural formula for each molecule, and determine if it is polar or non‐polar. Formula Lewis Dot Structure Structural Formula Geometric Shape NH3 SCl2 CF4 PCl3 H2S C2H2 18
Unit 4 ‐ Bonding and Shape Test Review Determine the formula for the compound formed by the two atoms and indicate if it is an ionic or covalent compound 1. Calcium and Oxygen 2. Nitrogen and Fluorine 3. Sodium and Chlorine 4. Carbon and Oxygen Draw the dot diagram for each of the IONIC compounds below 5. CaO 6. Na2S 7. SrF2 10. KI Review; http://www.youtube.com/watch?v=9DKId82_rUc Complete the table below. Formula Electron Dot Diagram Bonding Orbitals Shape Structural Formula Polar? NCl3 CO2 H2O CH3F 19
VSPER Worksheet:
1)
What is the main idea behind VSEPR theory?
2)
For each of the following compounds, draw the Lewis diagram, molecular
shapes, and structural diagram for all atoms:
a)
carbon tetrachloride
b)
BH3
c)
silicon disulfide
d)
PF3
e) carbon dioxide
f) SF2
20
Unit 4 Test Review 1. Calcium and Oxygen 2. Carbon and Fluorine Complete the table below Formula Electron Dot Diagram Shape Structural Formula Polar? CH4 H2O BH3 NH3 CO2 Draw the Lewis dot diagram for the compounds below 6. AlN 8. SCl2
10. BF3 Write the formula for the following compounds 12. Aluminum Bromide 14. Ammonium Fluoride 16. Copper (I) Chloride 7. HgCl2 9. NaF 11. SiCl4 13. Dinitrogen Tetroxide 15. Iron (II) Sulfate 17. Carbon Dioxide 21