Tests are not graded yetTurn in your project up front and work on

warm up:

Write the molecular formula for:Trinitrogen hexoxideAluminum nitrideCopper (II) sulfate

Write the names for:NO2

PCl3CaI2

Chapter 8

Covalent compounds consist of what? Only nonmetals

When naming, we use … Prefixes: mono, di, tri, tetra… Prefix = number of atoms (subscript)

N2O7

SF6

Why are there no charges (like in ionic compounds)?

In ionic compounds, electrons are _______________, so atoms gain or lose charge

In covalent compounds, electrons are _____________, so no charges are formed

What does the octet rule state? In order to be stable, an atom wants a full outer

shell (which generally means 8 valence electrons)

Which nonmetal is the exception to this rule? Which group do all elements want to be like?

When neither atom wants to give up their electrons, they will just share Electronegativity

When 2 electrons are shared between atoms, they form a single bond

When 2 or more atoms bondcovalently, this is called

a molecule

Consider ionization energy and electronegativity– when 2 elements are near each other on the periodic table, these values will be very near each other

Ionization energy Energy required to remove an electron

Electronegativity How well an element attracts electrons in

a bond

If both atoms have very similar strengths (for holding on to their electrons) then….

Neither one will be strong enough to take electrons away from the other

Lewis structures – using electron dot diagrams, shows the arrangement of the atoms in a molecule

How many valence electrons does carbon have? How many more electrons does it need to be “happy”? How many times do you think carbon will bond? How about hydrogen? Oxygen? Generally, the # of “missing” electrons will equal how

many times an element will bond

CH4

CCl4

Calculate the number of valence electrons Arrange the atoms in the molecule

○ Generally, the atom you have one of will go in the middle

○ Hydrogen only bonds once, bonds on the outside○ How many times will carbon bond? Oxygen? (look

at their valence electrons)

Put pairs of electrons between the central atom and all of the outer atoms

Put electrons to fill the central atom Put remaining electrons around outer

atoms Check to see that every atom is “happy”

PH3

H2S

SiH4

When 2 electrons are shared between atoms, you draw a line to show the bond

All other electrons that are not shared are called lone pairs and are included in the structure

Single covalent bonds are also called sigma bonds

Orbitals – the area where you will most likely find an electron How many electrons per orbital?

When these orbitals overlap, they form a sigma bond (σ)

Let’s try carbon dioxide…

Sometimes, atoms may share more than 2 electrons

If 4 electrons are shared, how many bonds would there be?

This is called a double bond How many electrons would a triple bond share?

Double or triple bonds consist of sigma and pi bonds (π)

Draw: O2 N2 F2

What do you notice about the bonds? Bond length : the distance between two

bonding nuclei Which of these 3 do you think would have

the shortest bond length?

Warm up:Draw the Lewis structures for the following:C2H6 C2H4 C2H2

Keep in mind how many times each element wants to bond

As the number of bonds increases, the bond length becomes shorter

Which bond would be the strongest?

Bond dissociation energy : energy required to break a bond in a molecule

What is the relationship between bond length and bond dissociation energy? Shorter bonds = more energy

In chemical reactions, bonds are broken and formed

Breaking bonds _____________ energy Requires (breaking a stick)

Forming bonds _____________ energy Gives off (Aladdin)

If more energy goes in, then it is _______________ Endothermic

If more energy is given off, then it is ___________ Exothermic

PO43- what is this called?

When an ion has a charge, that means it has lost or gained ______________

What has phosphate done?

Start the lewis structure like we did for the others – add up all valence electrons

Now we have 3 extra electrons

ClO4-

NH4+

CO32-

H3O+

sulfite

H2SO4

CH3OH

HCN

Warm up:

Name and draw the Lewis structures for the following compounds

H3P

CS2

N2H2

H – 1 time O – 2 times N – 3 times C – 4 times

Lowest electronegativity element goes in the center

Look at the word… Molecules that contain how many atoms?

My fish’s name will help you know these

In nature, when these elements are not bonded to another element, they like to exist with 2 of themselves. They are more stable that way.

What does it mean when something resonates? To vibrate or sound, especially in

response to another vibration

Resonance structures are different ways to draw Lewis structures for a molecule or ion

Only the arrangement of the electrons is changed

Let’s draw the structure for NO3-

How many resonance structures do each of these have?

O3

NO2-

SO2

CCl2O

Sometimes an atom may not obey the octet rule

Odd number of valence electrons (NO2) Fulfill the octet of the “outer” atoms

Less than 8 electrons present around an atom (BH3) Compounds with Be or B Tend to be very reactive Coordinate covalent bond – when one atom

donates both electrons in a shared pair (BH3 + NH3)

Draw the Lewis structure for SO3 and draw its resonance structures

Draw the Lewis structure for ClF3

Expanded octet: happens with elements in period 3 and below – d orbital electrons can hold more than 8

Generally, the central atom gets the extra electrons

PCl5 SF6

Let’s look at H2SO4 again The S-O bonds have been experimentally

determined shorter than single bonds

ClF5

More than an octet on chlorine

ICl4 -1

More than an octet on iodine

BeH2

Less than an octet - Beryllium and boron generally follow the less than 8 exception

NO Odd number of valence - Nitrogen generally

takes the odd number of electrons

Draw the Lewis structures for ammonia (NH3) and the ammonium ion

The hypothetical charge on an atom in a covalently bonded molecule

Helps to determine the best Lewis structure Want to keep the formal charge low – most

stable structure

FC = (# valence e-) – [(# of bonds) + (# of unshared e-)]

In a molecule, the sum of the formal charges (for every atom in the molecule) is zero

In a polyatomic ion, the sum is equal to the charge

Use the structures for NH3 and NH4+ from

the warm up

Determine the FC for each nitrogen and hydrogen in both structures Write the value next to the atom; if there is no

number, it is understood to be zero

Draw the structure for NOCl There are 2 possibilities, one is more preferred

Draw the structure for sulfate

Draw the structures and determine the FC for each atom

Cl2O

SO2

AsF3

Valence Shell Electron Pair Repulsion – used to determine the shape of a molecule

What determines how a molecule will arrange itself? What part of the atom are we generally

concerned about?... ELECTRONS

Something to keep in mind: lone pair electrons occupy more space than bonded electrons

On a separate sheet, draw the Lewis structures for each of the compounds on the handout

Let’s see how many bonded pairs there are, and how many lone pairs on the central atom there are

Don’t fill in the picture column or angle column yet

Linear

Bent

Trigonal planar

Tetrahedral

Trigonal pyramidal

Trigonal bypramidal

Octahedral

107.3o

104.5o

120o

109.5o

90o/ 120o

180o

90o

If the bond is not lying in the plane, then you use either dashes or wedges

When electrons are bonded, think of them as “trapped” between the 2 atoms, therefore occupying less space

Lone pairs occupy more space, therefore causing the bonded electrons to repel (and bend the molecule)

NCl3 OCl2 HOF NHF2

CO2

H2Se CH2O NH4

+1

Pick one of the VSEPR shapes and build a molecule

Include: label the type, an example of a specific molecule (none that are on the table), the angle between the atoms, represent lone pairs (if there are any)

Use anything you would like to build this – no drawings, and the model must be an accurate representation of the shape

Due next Wedn. Feb 10th

Hybrid – when 2 things combine and have properties of both

When atoms bond, they want to arrange their orbitals to have lowest energy possible

Hybridization – describes the arrangement of the orbitals

Hybrid orbitals – combined orbitals; intermediates between orbitals between s and p lies the hybrid orbital sp

Draw the orbital diagram for Carbon

From this, it looks as if there are only 2 places for electrons from another atom to pair up (in the p orbital), but how many times does carbon like to bond?

sp3

Write the formulas for the following compounds:

Aluminum sulfate Iron (III) phosphide Hydronitric acid Nitrous acid Dicarbon trisulfide

Regions of high e- density

VSEPR shape Hybridization

2 Linear sp

3 Trigonal planar sp2

4 Tetrahedral sp3

5 Trigonal bipyramidal sp3d

6 octahedral sp3d2

• When giving the hybridization, you are generally talking about the hybridization for the central atom

Generally, the # of things you are bonded to = the number of hybrid orbitals Bonded to 2 things = sp

Lone pairs(on the central atom) occupy hybrid orbitals as well Ex: draw the Lewis structure for water

Those 2 lone pairs count towards the hybrid orbitals, so water is sp3

NCl3 OCl2 HOF NHF2

CO2

H2Se CH2O NH4

+1

If something is polar, it means it has opposing ends

Need to know electronegativity and shapes

Influenced by the electronegativities of atoms in a molecule What is electronegativity? An atom’s attraction for electrons when

in a bond What is the trend for electronegativity?

(remember shielding and nuclear strength)

Who has the highest electronegativity value?

Ionic: Look at the electronegativities of Na and Cl – who has more attraction for the electrons?

Covalent: look at the values for the nonmetals Polar covalent – unequal sharing of the

electrons in a bond Nonpolar covalent – equal sharing of

electrons in a bond

Electronegativity Difference Bond Type

Less than 0.4 Nonpolar covalent

0.5 to 1.9 Polar covalent

Greater than 2.0 Ionic

What kind of bond would carbon and oxygen form?

Phosphorus and fluorine? Chlorine and chlorine?

Draw the Lewis structure, determine the shape and hybridization for the following:

BF3

SF4

PF6-

Draw the Lewis structure for water What is water’s shape? Who is stronger? Who will the electrons be closer to?

This makes partial charges.

Draw carbon tetrachloride and label the partial charges

Compare carbon tetrachloride’s structure to water’s Polar molecules are asymmetric, while

nonpolar are symmetrical Which one of these would you consider

symmetrical? You have to look at the polarity of each

bond, and look at the overall molecule to determine if it is polar

Determine if the following molecules/ion are polar:

NCl3H2S

CS2

SF6

If the bonds are polar, it could be polar or nonpolar, check the structure

Solubility (what is this?) is determined by polarity

What is the universal solvent?

Are most substances polar or nonpolar?

Determine the more polar molecule in each pair:

methyl chloride (CH3Cl) or methyl bromide (CH3Br)

water or hydrogen sulfide (H2S)

hydrochloric acid or hydroiodic acid

boron trihydride OR ammonia

silicon tetrabromide OR HCN

What were the properties of ionic compounds in terms of conductivity, melting point and solubility? High melting point, conducts (when dissociated),

and soluble in water (meaning ionic compounds are what?)

What are properties of covalent? Many covalent compounds exist as liquid or gas

Which type is more strongly held together?

What are intermolecular forces? (interstate) Forces that hold one molecule to another

3 Types:Hydrogen bondingDipole-dipoleDispersion/London forces

In the solid/liquid state (not concerned with gaseous state – why?)

Dipole – contains oppositely charged regions (partial charges)

Results from the attraction between the partial positive end of one molecule and partial negative end of another molecule

Also known as induced dipole forces animation

Occur between nonpolar molecules with no permanent dipoles

Result from a temporary shift of electrons, and dipoles are instantaneously created Ex. 2 chlorine molecules

Occur between hydrogen and O, N or F Due to their high electronegativities it makes

H more partially positive Causes these compounds to have higher

boiling points

What is the strongest intermolecular force present for each of the following compounds?

1) water 2) carbon tetrachloride3) ammonia4) carbon dioxide5) phosphorus trichloride6) nitrogen7) ethane (C2H6)

8) acetone (CH2O)

9) methanol (CH3OH)

10) borane (BH3)

1) water hydrogen bonding2) carbon tetrachloride London dispersion forces3) ammonia hydrogen bonding 4)carbon dioxide London dispersion forces 5)phosphorus trichloride dipole-dipole forces 6)nitrogen London dispersion forces 7)ethane (C2H6) London dispersion forces

 8)acetone (CH2O) dipole-dipole forces

 9)methanol (CH3OH) hydrogen bonding

 10) borane (BH3) dipole-dipole forces

Grab a chemistry book, and work on the following questions –

p. 274 83, 85, 89, 96, 98, 101, 108, 112, 114, 120, 127

Be sure to look through my powerpoints and study guide on my website

Name the following compounds: ZnCl2

KNO3

H2S

NF3

Name and draw the Lewis structures for the following compounds: CS2

PH3

CCl4

Write the formulas for the following compounds:

Aluminum sulfate Iron (III) phosphide Hydronitric acid Nitrous acid Dicarbon trisulfide

Go ahead and take out the worksheet with the PT with electronegativities from yesterday

Name the following acids:

H3N

H3SO3

H2Se