Thursday, September 16, 2010

where I left off.......preparing for exam I

So anyway, let me continue.......

Dalton's atomic theory was tweaked over time (as we like to say today). He originally said that atoms were the most basic unit of nature and that it was the smallest unit of matter. This is both true and not quite true today. It is true that in a chemical reaction, the smallest unit  is an atom. We look at whole number of atoms per formula unit. You cannot have a formula unit with half of an atom of any element.  However, it is not quite true in that subatomic particles exist. The smallest unit is, I think, yet to be discovered. However, the subatomic particles we study in this class are the proton, neutron and electron which I will talk about in a minute.

The other part of Dalton's theory that needed to be tweaked was his idea that all atoms of the same element have identical mass and chemical properties. This is not true because of the notion of neutrons (of which Dalton was not aware). Neutrons alter the mass of the nucleus and the overall atom. We now treat the mass of any given element on the periodic table as a representation of the relative abundance of that element in nature. 

Protons, neutrons and electrons and why they matter.....
Protons are the crux of all chemistry because they define the identity of an element. The number of protons is like someone's last name. It identifies their characteristics. If something has 6 protons it is carbon regardless of whether or not it contains neutrons or how many neutrons it contains. The number of protons defines the type element you are dealing with: other subatomic particles dictate the variety of that particular element.

Electrons are the subatomic particle that gives flavor to the element. Is it positively or negatively charged? This will depend on the ration of protons to electrons. Naturally something with 11 protons (Na) and 10 electrons will have a 1+ charge. This is very useful to understand in terms of formulating compounds that have an overall neutral charge.  We talked a bit about how JJ Thompson discovered the electron with a cathode ray tube and then Millikin came along and measured the mass and charge of the electron with the oil drop experiment. Ultimately, it became clear that the electron was exponentially smaller than the proton/neutron.  For this reason we discount its mass when calculating mass values on the periodic table.

Neutrons are their own category of subatomic particle that provide neither overall charge nor chemical identity to an element. Neutrons influence mass value only. Remember that we performed several relative mass value calculations in class to find relative atomic mass values reported on the periodic table.

The plum pudding model of the atom was eventually debunked by Rutherford who performed the gold-foil experiment to show that the nucleus of the atom contained most of the mass of the atom surrounding by a cloud of tiny electrons.

We talked about different notation systems to show number of protons, neutrons and mass value for various elements. Know the various notation systems and be able to use them on a test.

What are isotopes? Isotopes are different forms of an element that differ only in number of neutrons. Typically radioactive isotops are found in nature in small amounts and used as tools for radiometric dating and other scientific tools.

Once you understand that the periodic table is arranged by atomic number (# of protons) and you understand where the mass value comes from you are ready to start understanding how different columns/families tend to form different ions. Groups 1,2,3 tend to form 1+, 2+, 3+ charges respectively. The anions on the right side tend to form minus charges relative to how far they are located from group 8 on the periodic table. Why? Each group wants to gain the number of electrons that is identical to the noble gases. Once that magical "octet" of electrons is formed the ion is in its most stable state.

Familiarize yourself with the names of the groups in the periodic table. Know the halogens, alkali metals, alkaline earth metals and the noble gases. Know how to determine charge of an ion based on position in the periodic table.

Know your periodic table trends. Where are the metals? Where are the nonmetals? Can you identify chemical properties just by looking at position on the periodic table?

Chapter 3:
The key to chapter 3 is knowing how to classify a compound by its general appearance. For example, can you classify the following as molecular, ionic or acid?

HBr      HCl03,      NaCl,            CuO,                     Fe2O3,            NO2
Forgive me, I haven't figured out how to get a subscript to appear as a subscript in blogger. Perhaps that element of sophistication will evolve over time.

The answers are as follows:  binary acid,   polyatomic acid,    ionic,   ionic (variable charge),   ionic (variable charge), molecular

Now you use the different categories to use rules relevant to naming in that category.
Binary Acids: Start with "hydro" and add name of anion (bromine) with "ic" on the end: hydrobromic acid
polyatomic acids: Use name of polyatomic (chlorate) with "ic" and acid. chloric acid
Ionic compound (one charge): Sodium is group I and these always form +1 charge. Name all ionic compounds with "ide" ending: sodium chloride
Ionic Compound (variable charge for transition metals): Copper (name metal) (II) oxide (name charge with roman numeral and anion with "ide")
ionic compound (variable charge): Iron (name metal) (III) [charge of metal] and then "oxide" always an "ide" ending for ionic compounds
NO2: two nonmetals make a molecular compound. Use prefixes for second element (not necessary for first) nitrogen dioxide  (It would be nitrogen monoxide with only one oxygen)

That, in a nutshell is chapter 1-3 of the text!