2.5-2.7 Quiz Flashcards
What elements may have less than an octet in a Lewis diagram?
H, Be, and B
What elements may have an expanded octet in a Lewis diagram?
Any nonmetal in periods 3 or greater, as a central atom only.
Meaning of the HONC acronym and how to use it
Full acronym is HONC-1234
H- hydrogen(1)
O- oxygen(2)
N- nitrogen(3)
C- carbon(4)
The numbers associated with each one is how many bonds that those elements need to make. So ensure that your Lewis diagram is correct by using HONC-1234
What element is typically in the middle of a Lewis diagram?
The least electronegative
Practice drawing Lewis diagrams
2.5 in chapter 2 packet
What is a rule with halogens and Lewis diagrams?
Never double bond halogens
Why do bonds end up sharing electrons?
Because they do not have enough to complete their octet so they bond with another element to get a full valence shell(8 valence electrons)
Ex: H and C; Carbon has 4 valence electrons and they all need someone to pair up with so their compound is CH4 because every valence electron needs to make a bond to get to that full octet.
How can you confirm that a double bond(or more) has formed?
Through bond energy
Ex: F2 has a bond energy of around 155, but O2 has a bond energy of around 495 because it created a double bond and that bond is harder to break which causes its increase in bond energy
How does the number of valence electrons change if you have a -1 charge? a +1 charge?
You add an electron; You subtract an electron
What is different about Lewis diagrams for ions?
They have brackets around them and their charge written on the outside
Formal charge equation
How do you check that your formal charge calculation is accurate?
(# valence electrons)- (# lone electrons(count each dot as one) + # bonds connected(one dash counts as 1))
The sum of the formal charges of each atom should add up to zero for a molecule or add up to the charge for a polyatomic ion.
What is the point of formal charge? How do you know which structure is most valid?
It helps you decide which structure is most valid. The most valid structure will have formal charges closest to zero for all atoms(best structure will have no formal charge(aka zero)). Also, negative formal charges will be placed on the most electronegative atom.
How do you draw resonance structures?
Draw the possible structures with brackets around them and arrows between them
Bond order equation
of bonds/ Total bonding domains
What is the rule of thumb when drawing resonance structures with elements capable of expanded octets?
Try to follow the rule of octet and only use the expanded octet if it reduces formal charge
If you were to perform the reaction: KCl yields K+ Cl would energy be released?
No, the separation of ions requires energy
Ionizing an H2 molecule to H2+ changes the strength of the bond. Based on the description of covalent bonding given previously, do you expect the H—H bond in H2+ to be weaker or stronger than the H—H bond in H2?
Weaker, because a H—H covalent bond in H2+ has one less electron than in H2.
Suppose a Lewis structure for a neutral fluorine-containing molecule results in a formal charge of +1 on the fluorine atom. What conclusion would you draw?
There must be a better Lewis structure, since F is the most electronegative element and it should carry a negative formal charge.
Why do we draw resonance structures?
To show all possible structures and that the real structure is an average of all of those possibilities
How do you draw a resonance hybrid structure?
You draw one structure instead of all of the possibilities and you draw a dotted line instead of a full line on all of the outside elements. With the hybrid structure you use bond order.
Ex: CO3 2-; you draw a dotted line to each oxygen instead of drawing three different resonance structures where a different oxygen has a double bond. Bond order will show that this polyatomic ion has a bond order of 1 1/3
Practice drawing resonance structures
Chapter 2 packet- 2.6
2.6 daily video 1
VSEPR Theory
It uses the Coulombic repulsion between electrons as a basis for predicting the arrangement of electron pairs around a central atom
Hybridization/ hybrid atomic orbital
Used to describe the arrangement of electrons around a central atom
Hybridization of molecule with bond angles around 180 degrees
sp
Hybridization of molecule with bond angles around 120 degrees
sp^2
Hybridization of molecule with bond angles around 109.5 degrees
sp^3
How is the shape or geometry determined?
By the lone pairs and bonds on its central atoms because these areas will minimize electron-electron repulsion by positioning themselves as far apart as possible. Lone pairs of electrons repel more than bonds and tend to compress the angles between bonding atoms.
How do you determine the molecular geometry if there is more than one central atom?
If there is more than one central atom we determine the geometric shape of each central atom individually
by following the normal steps for each central atom.
Molecular vs Electron Geometries
Molecular geometry goes into more detail and includes options for lone pairs but electron geometry is more basic and just goes over number of bonding domains
Molecular, Electron Geometry, and Bond Angle of 2 bonds and 0 lone pairs
Linear; 180
Molecular and Electron Geometry and Bond Angle of 3 bonds and 0 lone pairs
Trigonal Planar; 120
Molecular and Electron Geometry and Bond Angle of 2 bonds and 1 lone pairs
Bent; trigonal planar; <120
Molecular and Electron Geometry and Bond Angle of 4 bonds and 0 lone pairs
Tetrahedral; 109.5
Molecular and Electron Geometry and Bond Angle of 3 bonds and 1 lone pairs
Trigonal pyramidal; tetrahedral; <109.5
Molecular and Electron Geometry and Bond Angle of 2 bonds and 2 lone pairs
Bent; tetrahedral; <109.5
Molecular and Electron Geometry and Bond Angle of 5 bonds and 0 lone pairs
Trigonal bipyramidal; 90 and 120
Molecular and Electron Geometry and Bond Angle of 4 bonds and 1 lone pairs
Seesaw; trigonal bipyramidal; <90 and <120
Molecular and Electron Geometry and Bond Angle of 3 bonds and 2 lone pairs
T- shaped; trigonal bipyramidal; <90
Molecular and Electron Geometry and Bond Angle of 2 bonds and 3 lone pairs
Linear; trigonal bipyramidal; 180
Molecular and Electron Geometry and Bond Angle of 6 bonds and 0 lone pairs
octahedral; 90
Molecular and Electron Geometry and Bond Angle of 5 bonds and 1 lone pairs
square pyramidal; octahedral; <90
Molecular and Electron Geometry and Bond Angle of 4 bonds and 2 lone pairs
Square planar; octahedral; 90
Sigma bonds
Very strong bonds that are seen as single bonds
Pi bonds
It pulls the atoms closer together. There is:
One in a double bond
Two in a triple bond
In sigma and pi, what are double and triple bonds made up of?
Double- 1 sigma and 1 pi
Triple- 1 sigma and 2 pi
What happens to bonds as the number of bonds between two atoms increases?
The bond increases in strength and energy but decreases in length. Triple bonds are the strongest and shortest bonds. While single bonds are the weakest and longest bonds.
Bond order
Number of bonds between two atoms
Explain how bond order works.
For a molecule that exhibits resonance, we see that the bonds that have resonance are experimentally determined to be the same length. In ozone, O3, we would expect the single bond to be shorter than the double bond, but that is not the case. They both have a bond length that is halfway between a single and double bond, so the bond order is halfway between a single and double bond order, and its bond order is 1.5. Another way to think about this is that there are three bonds shared between two atoms, and 3/2 = 1.5.
How to determine the polarity of a molecule?
If it only has nonpolar bonds in it, then it will definitely be nonpolar.
If symmetrically placed atoms surround it then it is nonpolar even if those bonds are polar.(linear, trigonal planar, tetrahedral, trigonal bipyramidal, square planar, octahedral)
If the molecule has an asymmetrical shape then it is polar.
