Electrons in Principle Energy Levels

  1. Arrangement of electrons in principle energy levels.
  1. What keeps an electron from crashing into the nucleus?
  1. Electrons in an atom have their to certain energy levels. So instead of being all around the nucleus the electrons are located in specific or energy levels. An orbital is the region of space surrounding a nucleus in which there is a high probability of finding up to electrons. The of these levels increases as the distance between the electron and the nucleus . This is a , which means it is a property that can have only certain values, that is not all values are allowed.
  2. These energy levels are called electron shells in this book and in others are called and are designated by whole numbers 1 - 7. is the region of space around the nucleus that contains electrons that have approximately the same energy and that spend most of their time approximately the distance from the nucleus. You can think of energy levels like steps on stairs.

a)For an electron to go to a higher energy level, it must energy.

b)For an electron to go to a lower energy level, it must energy. This can be seen as .

c)The energy emitted is to the energy difference of the two levels.

d)The maximum number of electrons in an energy level is , where n = the principle energy level (i.e. 1 – 7).

N=5 ______Bigger gap equals smaller  (wavelength)

N=4 ______The bigger the gap the more energy.

N=3 ______Energy is inversely proportional to wavelength.

N=2 ______

N=1 ______

  1. Arrangement of electrons in sub-levels or sub-shells.
  1. Principle energy levels are divided into sublevels also called .
  1. These sub-levels are labeled , and
  1. Each orbital has its own shape.
  1. s orbital =
  2. p orbital = on an axis – px is on the x axis, ….

px holds 2 e-, py …  6 electrons total for the p orbital.

  1. d orbital = double dumb-bell shape.
  2. f orbital = very strange.
  1. They have a maximum number of electrons that they can contain.

s = electrons

p = electrons

d = electrons

f = electrons

Filling the sublevels

  1. Lowest energy sublevel is filled first. principle states that electrons normally occupy electron sub-shells in an atom in order of increasing energy.

1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p<6s<4f<5d<6p<7s<5f<6d.

The principle number being smaller does mean that it fills first.

  1. Each sublevel can only hold so many electrons (identified earlier: s = 2; p = 6; d = 10; f = 14).
  2. exclusion principle. This says that each orbital can have only 2 electrons.
  3. rule. This says that you fill orbitals so that you maximize the number of electrons.
  4. Fill each sublevel before proceeding to the next until you run out of electrons.

Bonding

Ionic bond

  1. Ionic bond – the force of attraction between ions of charge. This holds ions together in an ionic compound.
  2. These oppositely charged ions are formed by theof electrons from one atom to another.

Covalent bond

Chemical bond formed by of electrons of the atoms

The smallest unit of a covalent bond is a compound

Covalent vs. Ionic compounds

Covalent compounds have melting points than ionic compounds. O2 is a covalent compound.

Covalent compounds conduct electric current in a liquid solution. Ionic compounds do conduct electricity.

All of the are covalent compounds H2, N2, O2, F2, Cl2, Br2, I2

Bond length: the distance between two of atoms that are bonded together.

Energy is given off in the formation of bonds. You must add energy to break bonds.

Electronegativity of covalent bonds.

The the radius of the atom the greater the attraction of the nucleus to the outer electrons.

Coulomb’s Law: between the electron and the proton increases as the distance decreases

As the distance decreases the energy decreases until a bond is formed. Then if you try to make the distance smaller you must add a large amount of energy.

Atoms with fewer electrons have less shielding of the outer electrons. Shielding would block the charge of the nucleus from the electron

a.As you get more electrons you get more distance between the outer electrons and the nucleus.

When filling the same period the electronegativity increases as the nuclear charge increases.

Polar Bonds

Differences in electronegativity cause of electrons in covalent bonds.

Electronegativity differences

nonpolar covalent bond

moderately polar covalent

very polar covalent

ionic bond

C.ExampleHCl

ENofH=2.1

ENofCl=3.0

The difference is .9 this is a polar bond

H2OEN of H=2.1

EN of O=3.5

The difference is 1.4 polar bond

N2EN of N=3.0

EN of N=3.0

The difference is 0 covalent bond

LiFEN of Li=1.0

EN of F = 4.0

Ionic bond

Metallic Bond

The metallic bond is formed between two or more metals where the valance electrons are shared between the metals. This sharing of electrons makes a or sea of electrons around the metals and this is why metals are good conductors of electricity.

The Covalent bond.

A type of a chemical bond formed when supplies both electrons of the electron pair that make a chemical bond and the other atom only has an empty orbital.

ExampleNH4+

EN of N=3.0

EN of H=2.1

.9 Polar

Why do we care about all this bonding?

The internal bonding of the molecules determines the way the atoms come together and the structures that are formed.

Why do we do the modeling lab that shows that the structure of is and not linear? It is because of the bent shape of water that water has the properties that it does. The bent shape allows for between the different molecules of water. Figure 8.26 p.241 in the book. See the animation

This hydrogen bonding holds water together and gives it a much boiling point than you would expect if the there was no hydrogen bonding. If water did not have a bent shape, it would not be a polar compound, it would not hydrogen bond and we would not be here.