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OSU / Chemistry / CHEM 1210 / What is it called when you have 8 valence electrons?

What is it called when you have 8 valence electrons?

What is it called when you have 8 valence electrons?


School: Ohio State University
Department: Chemistry
Course: Basics of Chemistry
Term: Fall 2016
Tags: General Chemistry
Cost: 50
Name: Chemistry 1210 midterm 3
Description: The study guide for the upcoming midterm
Uploaded: 11/07/2016
6 Pages 83 Views 3 Unlocks

Chemistry 1210 Midterm 3

What is it called when you have 8 valence electrons?

Definition=Green Important=yellow 

Chapter 8: Chemical Bonding 

Octet Rule and Lewis Structures 

∙ The octet rule states that atoms in compounds should have 8 valence electrons.  o This rule is important when drawing Lewis structures because it is the basic principle of  Lewis structures  

▪ All atoms in a Lewis structure must have a valence of 8 electrons  

∙ Some expectations are ClO2 , NO , & NO2which had an odd number of  

valence electrons

∙ Atoms in period 3 and beyond can have more as they can use their  

Why are period 3 elements exceed an octet?

empty d shell for the extra electrons

∙ Some atoms with lower than an octet are Be, B, and Al 

Lewis structures are drawn by having the atom surrounded by dots which represent the amount of  valence electrons it has. The number of valence electrons can be taken from the group that the electron  is in  

∙ For transition metals, the group restarts at the beginning of the period being 3B for most  periodic tables. Once the section is done. It continues with the first period of nonmetals o Transition metals also have a special way of writing electron configuration. For example  Ti = 4s23d2, Ti+2 = 3d2, and Ti+4 = [Ar] . These electron configurations happen because the  lowest energy electrons are removed first so that it can get closer to a Nobel gas  configuration as show with Ti+4.

How many valence electrons are in the pyridinium ion?

When drawing a Lewis structure, it may be necessary to use either a single, double, or triple bond.  If you want to learn more check out At what time period does adolescence take place?

∙ The stronger the bond, the shorter it becomes due to the increase in attractive forces between the atoms nucleus’

o Strength = triple > double > single If you want to learn more check out Which group is most impacted by regressive taxes?

▪ The strength increases with the number of bonds because those bonds will  include more pi bonds which are stronger than sigma bonds

o Bond length = single > double > triple  

When a compound has more than one Lewis structure, that is referred to as a resonance structure

∙ Can be seen with NO3-where the double bond switched places with each of the oxygens.  ∙ To determine the correct structure, the formal charge of the entire molecule must be closest to  zero with the most EN atom having the most negative number  

o Formal charge = (number of valance electrons for neutral atom) – ½(bonding electrons)  – (unshared electrons) 

Lattice Energy

∙ Def: the amount of energy used to pull electrons from an atom as well as the stability og an  ionic substance Don't forget about the age old question of What is m and n in a matrix?

∙ Two things are important when determining lattice energy of atoms

1. Charge: the higher the change means more protons so it will be harder to pull electrons  from the atom

▪ Remember: atomic radius increases down and to the right of a periodic talk  2. Size (tie breaker): the smaller, the more energy it takes to pull electrons rom the atom ∙ Can be represented using a Born Haber cycle

1. The order of a born Haber cycle is from left to right: enthalpy of formation, enthalpy of  sublimation, ionization, bond disassociation energy, electron affinity, and lattice energy ▪ When bond disassociation is given, be sure to divide by 2 to get disassociation of  one molecule of the diatomic molecule

2. An up arrow means endothermic; a down arrow means exothermic

3. The amount on the left should equal the right, so you can solve for whatever step by  setting each side to itself If you want to learn more check out What is mogadishu famous for?

Bond Polarity and Electronegativity 

Electronegativity(EN): relative ability for an atom to attract electrons to itself 

∙ EN is used to determine if a covalent bond is polar or nonpolar based on the dipole moment  o Dipole moments is shown by one atom having positive sigma (goes to less  

electronegative) and the other atom having a negative sigma(goes to most  

electronegative atom)

o A bond is nonpolar if there is no dipole moment or the dipole moment is very  insignificant

o A bond is polar if there is a huge dipole moment between the atoms.

▪ Remember: EN increase as we go up a table and from left to right.  

Therefore, Fluorine is the most electronegative  

Bond Enthalpy 

▪ Definition: the enthalpy required to break one mole of a particular substance in the gaseous state 

▪ This is determined by drawing out the molecule and looking at which bonds it has, then  using those bonds to get the bond enthalpy number. Then use the formula: (Summation  of bonds broken) – (summation of bonds formed). Be sure to take into account the  coefficients for each atom Don't forget about the age old question of What are the three waves of globalization?

o Not to be confused with Hess’s Law where it is reactants minus products, it is  actually the opposite in this case.  

o This value is always positive as breaking of bonds is always endothermic  

o The larger the bond enthalpy, the stronger the bond due to the atoms being closer  together so a harder time breaking the bond]

▪ Bond strength and bond enthalpies increase with the increasing number of  bonds between atoms. With increasing bond length, bond distance shortnens. 

Chapter 9: Molecular geometry and bonding theories

VSEPR Model  

▪ Explains the repulsions of electron domains amongst a central molecule being minimized in a  certain arrangement  

a. Electron domains is a region where electrons can be found. The two types are 1. Nonbonding electron domain: where the unshared electrons can be found 2. Binding electron domain is where the bonding electrons can be found

Electron domains are used to predict the electron and molecular geometry of a molecule. This can be  done in the following number of steps:  

1. Draw the lewis structure showing bonding and nonbonding electron domains 2. Count number of electron domains around central atom

1. Multiple bonds are considered as one electron domain. Example, a double and triple  bond is considered as one electron domain and not two or three We also discuss several other topics like What are amine hormones?

3. Use the following table to determine ED and molecular geometry using the amount of electron  domains  

a. Molecular domain geometry only considers the amount of bonding electrons b. Electron domain geometry considers bonding and nonbonding electron domains i. If molecular bonding has now lone pairs, that it is the same as the electron  domain geometry  

▪ This picture was obtained from Student Guide for Chemistry by James C. Hill, 12th edition.  The table looks a bit confusing but there’s a trick to understanding it all:

1. The molecular bonding is:

a. Linear

b. Trigonal planar

c. Tetrahedral

d. Trigonal bipyramid

e. Octahedral

The electron geometry is what happens when the molecular geometry is changed due to the addition of  nonbonding pairs:

a. Bent = 2 B, 1 NB Trigonal Planar , 3ED

b. Trigonal pyramid = 3 B, 1 NB Tetrahedral , 4 ED

c. Bent = 2B, 2NB

d. Seesaw = 4B, 1 NB

e. T shaped = 3B , 2NB Trigonal Bipyramid, 5ED

f. Linear 2B, 3NB

g. Square pyramidal =5B, 1 NB Octahedral , 6ED

h. Square planar = 4B,2NB  

a. The pattern is that the NB and B equal up to the molecular geometry and as you remove  EDs, it becomes easier to picture  

Molecular & electron orbital geometry 

∙ Hybridization is how molecules arrange itself in the lowest energy form and Is based on the  amount of electron domains; formed from the overlap of linear compound orbitals  o A tip: to determine hybridization, count the amount of bonding domains and that  corresponding number to the spppdddd in order to obtain hybridization. For example, a  molecule with four Electron domains would have a sp3 hybridization.  

▪ This can also be used to determine the amount of p orbitals left for bonding  A sigma bond is when electron density is concentrated along the internuclear axis 

▪ Is also considered a single bond and always the first bond in a double or triple bond ▪ The longest and weakest

▪ The formation for 2 s orbitals, an s and p orbital or p orbitals situated on the internuclear axis

A pi bond is when two parallel p orbitals along itself perpendicularity to one another and has the  electron density above and below the internuclear axis 

▪ Stronger than a sigma bond

▪ Has one pi bond in a s single bond and 2 pi bonds in a triple bond

▪ The overlap of p bonds causes a delocalized bond

o Resonance structure is one of the indications of delocalized electrons which make the  molecule more stable

Molecular orbitals happen when atomic orbitals combine

▪ Unpaired amount of electrons = paramagnetic and will react with an electric field ▪ All paired electrons will be diamagnetic and will not react with an electric field ▪ The amount of atomic orbitals is equivalent to the number of molecular orbitals.

▪ There are two different kinds of MO energy diagrams. After the second row, the sigma 2p  orbital is lower  

▪ When molecular bonding happens, bonding and nonbonding orbitals occur

o Bonding MO has electron density between two nuclei; can be thought of as sigma  bonding

o Nonbonding MOs are the representation of delocalized bonds in which the nuclei repel itself

Bond order is the net number of bonding electrons 

▪ Use the atomic number to determine amount of electrons that go on the orbital. Example N2  would have 14 electrons because it’s it has 14 protons therefore 14 electrons in its neutral state  ▪ (Number of bonding electrons – number of antibonding electrons) / 2

▪ Greater bond order = greater bond strength

o 1=single bond

o 2=double bond

o 3=triple bond

o 0=no bond and is completely unstable therefore won’t exist

Chapter 10: Gases (0nly up to 10.4)  

Gases have a particular set of properties:

a. They expand to fill the container they are in: no definite shape

b. They are compressible

c. The volume of the gas molecules is actually a small portion of the total volume of the gas in a  container  

d. Gas Is defined by F=P/A

Some instruments to measure pressure are;

▪ Barometer which is an apparatus filled with mercury where the height of the mercury is directly  proportional to the pressure of the gas 

▪ A manometer is used to measure the difference in height with the mercury, It is usually less than  atmospheric pressure as it is done inside a closed container

There are certain units of gas as well:

▪ Temperature in (K)

▪ Volume in (L)

▪ Quantity (N)  

▪ Pressure (In either atm, torr, pascals…)

o Important conversions: 1 atm=760mmHg=760 torr=101.3kPa 

Gas Laws 

Charles Law : volume is directly proportional; to temp at constant pressure

▪ V=T *P, n constant

Boyle’s Law: gas is inversely proportional to pressure at constant temp

▪ V=(1/p) *t,n constant  

Avogadro’s law: At constant pressure and temp, volume is directly proportional to moles of gas ▪ V=1*n *p,n constant  

All these equations lead to one overall equation called the ideal law equation:


o R=0.0821 L-atm/mol-k

o Pressure in atm

o Temp in K

o Volume in L

Other important equations:

a. Density of gas(g/L): P(Molar mass of gas)/RT

b. Molar mass (M): grams*R*T/PV

Daltons Law of Partial Pressures  

▪ States the total pressure is made up of all the pressures of the individual gasses  o Pt=nt*R*T/v **when you know the moles of all gases

Mole fraction: the ratio of moles in a component over the total moles of mixture 

▪ Xi=number of moles of component i/ total number of moles in mixture  ▪ Partial pressure can be calculates using the mole ratio like Pa=XiPt

Kinetic Molecular Theory of Gas  

▪ At same temp, gas molecule has the same average kinetic energy

▪ Lighter molecules move faster than heavier ones

▪ If kinetic speed changes so, does average speed

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