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Transcript of Tests are not graded yet Turn in your project up front and work on warm up:
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