4th Quarter Science Notes

January 5

Electron structure of atoms

*Neils Bohr - student of Ernest Rutherford



*Max Plank - thought that energy produced by atoms is quantized (has certain values)


*Note:
-the electrons in the atom can occupy a specific orbit
-energy increases as its distance from nucleus increases


*Modern model of the atom

SchrÖdinger - Austrian scientist who studied the modern model of the atom

orbitals - three-dimensional region in space where electrons are likely to be found






January 6

Principal Energy Level
-Orbitals of similar size can be found here


Principal Energy Level
|
|
Sublevels
/ \
| Orbitals
s, p, d, f
s - sharp lines
p - principle lines
d - diffuse lines
f - fundamental lines


*As the principal energy level increases, the orbital extends further away from the nucleus

*There are 3 p orbitals in a sublevel. They lie perpendicular to one another.
So that's 1s, 3p, 5d, and 7f for each orbital in a sublevel.

*Each orbital can hold two electrons, so s can hold a total of two each level, p can hold a total of 6, d can hold a total of 10, and f can hold a total of 14.

Principal Level@@@ Sublevel@@@ Orbitals
1@@@@@@@@@ 1s@@@@@@@@@ s1
2 @@@@@@@@@ 2s 2p@@@@@@@@@p3
3@@@@@@@@@ 3s 3p 3d@@@@@@@@@d5
4@@@@@@@@@4s 4p 4d 4f@@@@@@@@@f7

And so on. Level 4 has s, p, d, and f. So, 4s, 4p, 4d, and 4f pertain to the naming of the sublevels.







January 7

Orbital diagrams

-shows the sublevels and orbitals that can exist at each principal energy level.
-each box represents an orbital. Each group of boxes represents a sublevel.




*Orbital diagram rules
-two principles and one rule determine how the electrons are filled in the principle energy levels and sublevels.

*Aufbau Principle
-electrons fill orbitals starting with the lowest-energy orbitals

*Pauli exclusion principle
- a maximum of two electrons can occupy each orbital, and they must have opposite spins.

*Hund's rule
-electrons are distributed into orbitals of identical energy (same sublevel) in such a way as to give the maximum number of unpaired electrons.

*electrons are always filled in their ground state, or lowest energy state.






January 8

Electron Configuration
-shorter way of conveying an orbital diagram

E.C. of Na: 1s^2 2s^2 2p^6 3s^1
Electrons in Na: 11

E.C. dissected:
1s^2
1s is the orbital box. ^2 means "to the power of two", since electrons are expressed as superscripts here. So, 1s^2 would mean 2 electrons in box 1s.







January 11

Abbreviated Electron Configurations

EVEN SHORTER way of showing an orbital diagram


The image is pretty self explanatory.


Valence electrons for the Main-Group Elements

*The last filled principal energy level is called the valence level, or valence shell
*The valence level contains electrons that are highest in energy and occupy orbitals that extend further from the nucleus than those in the lower levels.
*Valence electrons - the number of electrons in the valence shell
*Elements in the same family have the same valence electrons

For example:
E.C. of Na: 1s^2 2s^2 2p^6 3s^1

When looking for the valence shell, look for the HIGHEST sublevel, not the LAST sublevel in order.

The valence shell of Na is 3.
How many electrons are there in the 3 sublevel? only 1.
So Na has 1 valence electron.








January 12

Electron Configuration for Ions

*In atoms the number of Electrons is equal to the numbers of protons, or the atomic number

*In ions, we must add or subtract electrons depending on whether the Ion is a Cation or an Anion.

So, you must always subtract electrons when it is a Cation (positive sign )
And add when it is an Anion (negative sign)
So, in adding or subtracting,m you do the opposite of each sign.

Let's take sodium for an example.

Na
Electrons: 11
Na^2-
That's an Na anion. Why? Its superscript denotes its charge: two minus.

So, in effect...

Na^2- has 13 electrons (add)
Na^2+ has 9 electrons (subtract)



January 18

*Electronic Configurations are related to the following properties:
-Relative Reactivity
-Ionization Energy (tendency to lose electrons)
-Atomic Radii (atomic size)
-The most reactive metals are the alkaline and alkaline-earth metals.
-The more reactive, the more easily an atom loses electrons to form ions.

*Ionization Energy - measure of the energy required to remove a valence electron from a gaseous atom.
*In general, atoms with low ionization energy do not bond their electrons very tightly, and, therefore, are very reactive.
*Ionization energy increases going right and up because the valence electrons become increasingly closer to the nucleus.


*Ionization energy increases going right and up. This means that, since reactivity is negatively proportional to ionization energy, (and vice versa) reactivity increases going left and down. (according to the periodic table)



More ionization energy, less reactivity. And vice versa.



To explain further, let us make an analogy.
The atom is comparable to two blocks of wood glued together.
One block of wood is the nucleus, the other block of wood is an electron.
The glue holding the wood together is ionization energy, and the probability
of the wood breaking apart is reactivity.

If the wood has plenty of glue, it has a low chance of falling apart.
If the wood does not have that much glue, however, it will most likely fall apart.






January 19

*Atomic Size is usually reported as atomic radius, but they have the same meaning: half the distance between the centers of two bonded atoms.
*Atomic size increases down a group because valence electrons are in orbitals that extend further from the nucleus.
*Atomic size decreases from left to right


*Sizes of ions
-Cations are smaller than their original atoms
-Anions are larger than their original atoms

Cations:
The higher the charge, the less the atomic size.


Anions:
The higher the charge, the greater the atomic size.







January 26

Types of chemical bonds
-Ionic - electrostatic forces: forces between oppositely charged ions
-Covalent covalent bonds in which electrons are shared

*Ionic Bonds
-Bonds created by electrostatic attraction between oppositely charged ions
-between metals and nonmetals
-Electrons transferred between the cation and anion
-extremely strong bonds

*Covalent Bonds
-bonds created by the sharing of electrons between atoms
-occurs between two nonmetals (neutral overall charge)
-electrons are shared, not transferred
-often shared in pairs
-bonds are generally weaker


*Polar bonds and nonpolar bonds
*Two types of covalent bonds:


*Polar Covalent
-uneven sharing of electrons (some are transferred)
-occurs when different elements bond to each other
-why? Because different elements have different electronegativity



*Nonpolar Covalent
-equal sharing
-occurs only when all of the atoms in a molecule belong to the same element

*Polar Covalent Bonds are:
-shorter (distance)
-stronger bonds

*Nonpolar covalent bonds are:
-Typically longer bonds
-weaker bonds

*Polarity
-occurs in polar covalent molecules
-polarity is the degree of transfer of electrons in a covalent bonded molecule composed of different elements atoms



January 28
*Lewis Dot Symbol
*Electron dot symbol
-Dots planned around an element's symbol represent valence electrons
*Pair electrons as needed

*Octet rule
-tendency of an atom to achieve an electron configuration having 8 valence electrons
-Same as the electron configuration of a noble gas
-the 8 electrons exist in 4 pairs
-Ions achieve 8 electrons by losing or gaining electrons

@@!..
@@:N:
@@!'!'

Maximize unpaired electrons!
Fill up in this order: right left up down


February 1

Behavior of gases at the molecular level

*Gases consist of particles (atoms or molecules) that are relatively far apart.
*Gas particles move about rapidly.
*An average oxygen molecule moves at a velocity of 980 miles per hour at room temperature.
*Gas particles have little effect on one another unless they collide. When they collide, they do not stick to one another.
*Gases expand to fill their containers.
*All gases expand when heated.
*Pressure (P) is the amount of force applied per unit area:
Pressure is equal to force divided by area
*we can describe the pressure exerted by gas particles by:
Pressure is equal to force of gas particles divided by the area of container

*psi
-pounds per square inch (pounds per inch squared)

*Atmospheric pressure is commonly measured using a barometer
*Inches of Hg
*mmHg
*torr
*atm


February 3

atm= atmosphere
pa=pascal

1 atm= 760 mmHg
1 atm= 760 torr
1 atm= 101,325 pa


Properties that affect gas:
*Volume
-Measured in Liters

*Pressure
-measured in atmosphere

*Temperature
-Measured in Kelvins (K)

*Particles
-Moles

-An Ideal gas is a gas that behaves according to predicted linear relations

*Volume & Pressure
are inversely proportional

Volume is equal to 1 divided by pressure


Week of February 15
Random facts:
Scientist who discovered principle assoc. with barometer: Evangelista Toricelli
English chemist who discovered relationship between pressure and volume in a gas: Robert Boyle
Gas that deviates from ideal gas: Real gas
French Scientist who discovered the relationship between temp and volume: Joseph Gay-Lussac


Transitions between states

*Also called phase changes
*6 main phase changes
*Evaporation- occurs because liquid molecules have high kinetic energy, requires heat.
*Condensation- gas particles cannot escape the container and come into contact with liquid
*Freezing- when temperature is decreased, particles lose kinetic energy and eventually stop moving
*Melting- also called fusion, same temperature as freezing point. (also called fusion because the energy required to melt is called the latent heat of fusion)
*Sublimation- solid to gas
*Deposition- gas to solid
*Endothermic- reaction requiring heat
*Exothermic- reaction that releases energy
*Equilibrium iss a state in which opposing proccesses occur at equal rates (example: in a closed cook pot, evaporation and condensation occur at same rates)
*Boiling point: when the vapor pressure is equal to the external pressure of the atmosphere.
*If the external pressure is 1atm, then that is called the "normal boiling point".



Composition of Solutions
*Solution - homogeneous mixture
*Solute- being dissolved, usually less that solvent
*Solvent- doing the dissolving
*Aqueous solution- solution with water as the solvent
*Solubility- how readily a substance dissolves in a solvent
*A soluble ionic compound dissociates into ions when dissolved in water
*insoluble compounds remain mostly undissociated in their ionic crystalline structure.
*Solubilites for ionic compounds cannot be readily predicted through a pattern
*Temperature - affects solubility of most compounds in water
*Most ionic solids are more soluble in water at higher temperatures
*Gases are less soluble as temperature increases
*solubility of a gas in liquid is directly proportional to the pressure of gas above the liquid
*more pressure = more solubility
*Concentration - relative amounts of solute and solvent
*Saturated solution- maximum amount of solute in a given amount of solvent
*Unsaturated solution -less solute than in a saturated solution
*Solubility describes the concentration of a saturated solution. (gsolute/100gsolvent)
*Supersaturated solution- more solute than saturated solutions

I LOVE the illustrations Lance.

Cassandra wrote:I LOVE the illustrations Lance.

Well... what can I say.