TRENDS ON THE PERIODIC TABLE: October 28, 2010

Elements close to eachother on the periodic table display similar characteristics

There are 7 important periodic trends:

1) Reactivity

2) Ion Charge

3) Melting Point

4) Atomic Radius

5) Ionization Energy

6) Electronegativity

7) Density

(There are more than just these 7, but these are the most important ones for us to learn at this time...)

1) REACTIVITY

- metals and non-metals shoe different trends

- the most reactive metal is Francium; the most reactive non-metal is Fluorine

-reactivity increases as you go down for metals and up for non-metals

-Noble gases are very unreactive

2) ION CHARGES

- elements ion charges depend on their group (column)

3) MELTING POINT

- elements in the centre of the table have the highest melting point

- noble gases have the lowest melting points

- starting from the left and moving right, melting point increases (until the middle of the table, it then starts to decrease)

-an exception to this rule is Carbon. Carbon has a high melting point!

4) ATOMIC RADIUS

- radius decrease to the up and the right

- Helium has the smallest atomic radius

- Francium has the largest atomic radius

5) IONIZATION ENERGY
- ionization energy is the energy needed to completely remove an electron from an atom

- it increases going up and to the right

- all noble gases have high ionization energy

- Helium has the highest I.E.

- Francium has the lowest I.E.

- opposite trend from atomic radius
-Quick note: think about why Ionization energy has the opposite trend of the Atomic radius. Since the atomic radius gets smaller when ionization energy increases, it tells us that when the shells are smaller, the energy needed to completely remove an electron is easier, therefore, electrons can leave their small atomic radius which is an increase in ionization energy.

6) ELECTRONEGATIVITY

- refers to how much atoms want to gain electrons
- same trend as I.E.

7) DENSITY

...yet to be learned!

and our own pictures!

Post by: Adrienne Ross (with pictures from Ren Flores)

ISOTOPES AND ATOMS: October 26, 2010

Today we learned about Isotopes! We have already learned that ions are atoms that are either missing or have extra electrons. Let's say an atom is missing a neutron or has an extra neutron. That type of atom is called an isotope. An atom is still the same element if it is missing an electron. The same goes for isotopes. They are still the same element. They are just a little different from every other atom of the same element.

*Note: the most common ion charge is listed on top (if theres 2 or more charges)

* Atomic mass - Atomic number = # of neutrons

*If you are given the atomic number and number of neutrons, add them to get the mass number

A decimal found in the atomic mass is an average of the isotopes

Going good so far? Great.

Mass spectrometry (MS) is an analytical technique that measures the mass-to-charge ratio of charged particles.^[1] It is used for determining masses of particles, for determining the elemental composition of a sample or molecule, and for elucidating the chemical structures of molecules, such as peptides and other chemical compounds. The MS principle consists of ionizing chemical compounds to generate charged molecules or molecule fragments and measurement of their mass-to-charge ratios.^[1] In a typical MS procedure:

1. A sample is loaded onto the MS instrument, and undergoes vaporization
2. The components of the sample are ionized by one of a variety of methods (e.g., by impacting them with an electron beam), which results in the formation of charged particles (ions)
3. The ions are separated according to their mass-to-charge ratio in an analyzer by electromagnetic fields
4. The ions are detected, usually by a quantitative method
5. The ion signal is processed into mass spectra

-Source: Wikipedia

Now its your turn! Fill out the chart on isotopes!

Answers: Row 1: 28 neutrons. Row 2: Mn(Manganese) Protons:25 Netruons: 31 Row 3: Mass # 12, Atomic #:6, # of protons: 6, # of neutrons 6

Post by Ren Flores

BOHR'S MODEL: October 15, 2010

Now that you know the history of atomic theory, let's move on to the more modern stuff, such as Bohr's atomic model!

BOHR (1920S)

• Rutherford's model was inherently usable (protons and electrons should attract to eachother, no?)
• Matter emits light when it is heated (black body radiation)
• Light travels as photons
• The engerdy photons carry depend on their wave length
• Bohr based his model on the engery (light) emitted by different atoms
• - Each atom has a specific spectra of light
• To explain this emission spectra, Bohr suggested that electrons occupy shells or orbitals

Summerizing Bohr's Theory

• Electrons exist in orbitals
• When they absorb energy, they move to a higher orbital
• As they fall from a higher orbital to a lower one, they release energy as a photon of light

Hmm...

When electrons move from level to level (depending on the energy), they don't travel there, they actually just appear there, you could call it "transporting".

This led to us learning (by questioning Mr.Doktor) that scientists have tried harnessing this power to transport more than just an electron. Too bad the most they've ever managed to transport was about 5 atoms...

Still, we never know what lies in the future! Maybe one day we will all quit walking and physical activity and just transport everywhere we go!

Borh-d? (Get it?...Haha.)

Check out

http://phet.colorado.edu/en/simulation/hydrogen-atom

to play around with a hydrogen atom and see many of the different atomic models!

Post by Adrienne Ross

ATOMIC THEORY: October 13, 2010

A timeline, and history, on the modern atomic theory...

ARISTOTLE

• Created the four elements theory
• Lasted about 2000 years
• Believed all matter was formed from 4 basic elements: fire, water, air, earth, which created/connected to make: hot, cold, wet, dry
• It is not a scientific theory because it could not be tested against observation

  

DEMOCRITUS (400-300 B.C.)

• Said atoms were invisible particles
• First mentioned atoms
• Not a testable theory, only a conceptual model
• No mention of any atomic nucleus or its constituents
• Cannot be used to explain chemical reactions

LAVOISIER (LATE 1700S)

• Law of conservation of mass
• Law of definited proportions; example: in a molecule of water (H2O), it will always be 11% H and 89% O, no matter how much or how little water there is

PROUST (1799)

• If a compound is broken down into its constituents, the product exists in the same ratio as the compound
• Experimentally proved Lavoisiers' Laws

DALTON (EARLY 1800S)

• Atoms are solid, indestructable spheres (like billiard balls)
• Provides for different elements (those would be different spheres)
• Based on law of conservation of mass
• Having a molecule (atoms combined in simple whole # ratios) explains the law of constant composition
• If the atoms are not destroyed then the mass does not change

J.J. THOMSON (1850S) 

• Raisin bun model
• Solid, positive spheres with negative particles embedded in them
• First atomic theory to have positive (protons) and negative (electrons) charges
• Demomstrated the existence of electrons using a Cathode Ray Tube

RUTHERFORD (1905)

• Showed that atoms have a positive, dense center with electrons outside it
• Resulted in planetary model
• Explains why electrons spin around nucleus

Post by Adrienne Ross

to be continued...

TEST DAY!: October 8, 2010

Today we had our Chemistry Unit test. Hopefully we did alright!

REVIEW: October 6, 2010

Today we were suppose to have a Chemistry test but Mr. Doktor is so cool that he postponed it to Friday! Instead of taking a difficult, back breaking chemistry test, we reviewed how to balance chemical equations, ionic formulas, physical and chemical changes, measurement, uncertainty, significant digits, and dimension conversions.

Mr. Doktor also told us to memorize the basic Common Metric Prefixes Used in Chemistry (i.e mega, kilo, deci)

In conclusion, we are ready for our test.........!

SODIUM CHLORIDE (LAB): October 4, 2010

Today, we went over the homework from last class which was page 52 in our text book.


Then we did our Sodium Chloride (a.k.a. Table Salt) lab. We had to figure out what amount of sodium chloride would dissolve in different amounts of water. 
We went into groups or pairs, Adrienne was my partner.We gathered our designated materials on the tables in the back of the classroom, including a beaker and a graduated cylinder, and prepared our table.

 The materials we used in this experiement included :
- sodium chloride (salt)
- 150mL of distilled water
- 100mL graduated cylinder
- weight paper
- electronic scale
- lab coat
- safety glasses


Our Chemistry Lab consisted of:

1. A Graph (Mass of Salt vs Volume of Water)
2. We included a title, axis points, data points, an appropriate scale and a line of best fit
3. The following class, we finished the “conclusion” portion of the lab where we calculated our percent error and most of us were way off. But thats okay, Chemistry is all a learning experience and one day our percent error will be very low.

Steps from our Lab:

 1. We put our coats and glasses on and made sure we listened to directions carefully from Mr.Doktor.
2. We safely plugged the scale in from sockets in the wall.
3. Next, we put the weight paper on the scale and Mr. Doktor put the sodium chloride on the weight paper for us.
4. We set the scales up so it would tell us how many grams we were taking off from it. Taking small amounts of salt, we dissolved it in water and did this until the salt wouldnt dissolve anymore. We knew that the salt would not dissolve anymore when we saw the crystals forming on the bottom of the beaker.
5. We repeated this 3 times, using 3 different amounts of water with the same temperature. The first amount, we used 10mL. The second amount, we used 20mL. The third amount, we used 40 mL.
6. Lastly, we recorded our data on the table provided. For all 3 different measurements of water we concluded that the more water, the more salt could be dissolved in it. Thus, the number of dissolved salt increased as we went up the mL scale.


However, we all know that error is a fundemental part of Science. Some variables in our experiment that could've played a role in the outcome are:
-The temperature of the water (hot,cold, or mild)
-The purity of the water (distilled or not)
-The incorrect unit measurement from the graduated cylinders and beakers
-Measuring the amount of water in the graduated cylinder incorrectly (not paying attention to the meniscus)

Ivys interpretation and drawing of her experience doing her first chemistry lab! Yay!

Post by Ivy Gloria/Edited by Ren Flores