Write the electron configurations for the following

• S2-

• Ca

• Br

• O2-

Draw Box diagrams for the following

• Co

• Al

• B

What did you discover in the periodic properties lab?

• What data can you use to support your claim?

• Alkali metals are more reactive than Alkaline earth metals

• As you go down the period things become more reactive.

• What patterns did you discover in the graphing activity?

• Atomic Size – Increases as you go down the periodic table– Decreases as you move from left to right

across the periodic table

• Ionization Energy – Decreases as you move down the periodic

table– Increases from left to right– Increases when an orbital is full

How does this relate to reactivity?

• Elements are more reactive as you move down the periodic table

• Relates to the amount of Electrons in an orbital.

• Atoms that only “need a few” electrons or can give away “a few” electrons to have a full orbital are more reactive.

What did you discover in the periodic properties lab?

• Which substance was the most volatile?

• Which substances had the lowest melting point

• Which substances conducted electricity?

• Which substances dissolved in water? Hexane?

Lets go back to the penny challenge…

• Why does water hold so many drops compared to hexane?

• Why does it bend toward charge?

• How does this relate to it being so unique?

• How are substances held together?

• Why are we able to live on the earth?

• Why is water so “unique”

• Why can bugs run across the water?

• Why do metals conduct electricity?

Chemical Bonds

Definition:The force that holds two atoms together.

Why does a bond form? So that an atom:

1. becomes more stable

2. takes on a noble gas configuration

To determine the type of bond

Electronegativity:

• measure of how strongly an atom attracts the electrons that are shared in a bond

• NaCl

• FeNO3

• KCl

• CsSO4

All these substances contain Ionic Bonds

What rules could you determine about ionic bonds from examining these compounds?

Types of Bonds

1. Ionic

The attraction between oppositely charged ions

• Atoms become ions by adding or losing electrons

• They form these charges to reach a noble gas configuration in their outer energy level

• Usually a Metal and a Non Metal

These compounds have covalent bonds. What rules could you

determine about covalent bonds?• CO2

• H2O

• CH4

• SiO2

2. Covalent Bonds

• A sharing of a pair of electrons between two atoms

• Individual atoms attain a noble gas configuration with the shared electrons in their outer energy level

Lewis Dot DrawingsShow the sharing or transfer of electrons

Also called “electron dot” drawings.

• Involve only valence electrons (those in the outermost energy level)

• Show the type of bond formed (either ionic or covalent)

• All atoms will satisfy the “octet” rule (except for hydrogen (duet rule) and metals)

Element valence electrons Lewis dot

N

O

F

C

Lewis dot drawings for1. Ionic bonds• Show electrons being transferred

• Include brackets and charges on ionsexamples:

• H and F• Na and Cl• Na and OH-

Lewis Structure for Covalent bonds

Technique: Place the atom with the largest number of unpaired electrons in the middle.

(Never put H in the middle of a molecule!!)

Determine how the electrons will be shared so that all atoms are stable (Octet Rule)

H2O

CH4

SCl2

Try these..

Double and Triple BondsExample: HCNMake a table:

atomhave needH 1 2C 4 8N 5 8total 10 18

Difference: 18-10=8 divide by 2 = 4You need 4 bonds in this structure

Sharing 4 or 6 electrons (Double or Triple bonds allow this to happen)

Try These

Examples

C3H6

SO2

Electron dot drawings for polyatomic ions

Always include brackets and charges, but have covalent bonds inside the ion

Count the number of valence electrons for each and the add or subtract and electron to make the correct charge

NH4+

OH-

SO42-

Draw NH4OH

Exceptions to the octet rule1. Metals

MgH2

BH3

2. Molecules with an odd number of electronsNO

NO2

3. Some Nonmetal atoms

because of their size, they can have more than an octet of electrons (due to the presence of empty “d” orbitals which can be used for bonding).

SF6

PCl5

DON’T FOCUS ON THESE BUT KNOW THEY OCCUR!

Try these….

• Mg(OH)2 C3H6 O2

Note:

• Not all covalent bonds have equal sharing of electrons…

• There are electron hogs!!! Elements that hold on to the electrons more tightly than others

• You can determine if a bond is ionic,covalent and if there is an electron hogs, through looking at a characteristic property.

To determine the type of bond

Lets remind ourselves…

Electronegativity:

• measure of how strongly an atom attracts the electrons that are shared in a bond

• The difference of electronegativity will determine the type of bond

What would you predict are the trends in electronegativity?

in families?in periods?

What family has the highest electronegativity?What family has the lowest electronegativity?What period has the highest electronegativity?What family has the highest electronegativity

Electronegativity

• Allows you to predict the nature of the bond between two atoms

• To determine where the electrons tend to spend the most time in the molecule

To determine the type of bond

• When the difference in electronegativity (ΔE.N.) is 2.0 or greater, the bond is ionic

Examples:NaClKF

Where are these atoms on the periodic table in relation to one another?

When the ΔE.N. is less than 2.0, the bond is covalent

Examples: H2O NO2

• This means the electrons spend more time around one of the elements giving it a partial charge

• Draw a picture of how you think the electrons would be distributed for each of these molecules.

When the electrons are shared equally

ex: H2 NCl3the bond is pure covalent and has no partial charge

Why do you think there would not be a partial charge on these bonds?

Which covalent bond do you think is stronger? H2 or N Cl

These bonds are called intramolecular forces

• Have various strengths– Ionic (STRONGEST)– Polar Covalent (NEXT STRONGEST)– Covalent (STRENGHTH DEPENDS ON

ELECTRONEGATIVITY DIFFERENCE)

Draw Partial Charge distribution for the following bonds

• C-F Si-H P-H

Classify the typeof bonds in each

• CH4

• H2O

• Na3PO4

• F2

Why are molecules a particular shape?

Shapes of Molecules/Compounds

IN Molecules Shape is

Determined by• # of bonds• Lone pair Electrons

Ionic substances

Ions stack together, anions alternating with cations

The structure of Ionic solids

Possible Shapes

• Linear Ex:

• Trigonal Planar

• Tetrahedral

• Triganol bipyramidal

• octahedral

2. Covalent compounds

The shape of the molecule is determined by the repulsion between the electrons that the atoms share

Understanding the placement of electrons can help us determine the shape of a molecule

VSEPR Theory

Valence Shell Electron Pair Repulsion

The shape of a molecule is determined by the repulsion between the electron of the bonded atoms

Molecules containing a central atom

• Electrons want to be as far apart as possible

• # of bonds(from central atom) influences shape

• Lone pairs of electrons on a central atom influence shape– They are more repulsive than bonded

electrons because they flare (take up more space)

Different shapes1. Molecules with only two atoms will

always be linearEx: CO HCl

2. Molecules with two bonds can have two different shapes

Example: BeCl2

H2O

This molecule is linear.

The H2O molecule is “bent” or “angular”

Effect of the lone pairs on shape or H2O

Molecules with 3 bonds Can have two shapes

Ex: NH3

BF3

This one is trigonal pyramidal because of the lone pair of electrons

• This one is trigonal planar due to the absence of lone pairs on the central atom

Molecules with 4 bonds

If there are no lone pairs on the central atom:

Ex: CH4

This is called tetrahedral

This makes all bonds equidistant from each other

Molecules with 5 bonds (and no lone pairs on the middle atom)

Ex.: PCl5

This is called trigonal bipyramidal

Molecules with 6 bonds and no lone pairs on the central atom

Ex.: SF6

This is called octahedral

• Predict the shapes of these molecules

• PH3

• SO2

• CO2

Practice• Make electron dot drawings of the

following substances and predict the shape

• F2

• SiO2

• PF5

• BF3

Back to water

Water is so unusual because of the forces

“within” the molecule”

These are called “intra molecular forces”