Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109
General Chemistry
Spring 2009

Lecture Notes:: 30 January
© R. Paselk 2002

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*Humboldt State University* ® *Department of Chemistry*

Richard A. Paselk
Chem 109	General Chemistry	Spring 2009
Lecture Notes:: 30 January	© R. Paselk 2002

PREVIOUS		*NEXT*

Chapter 2: Atoms Molecules & Ions, cont.

Dalton's Atomic Theory

These various laws imply that matter is made up of small discreet units, which we call atoms. The earliest "successful" theory of atoms is that of John Dalton (1803). Dalton's atomic theory states:

All matter is composed of ultimately small particles, called atoms.
Atoms are permanent and indivisible - they can neither be created nor destroyed.
Elements are characterized by their atoms. All atoms of a given element are identical in all respects. Atoms of different elements have different properties.
Chemical change consists of a combination, separation, or rearrangement of atoms.
Chemical compounds are composed of atoms of two or more elements in fixed ratios.

All of these statements are close to reality, and nearly describe chemical behavior. But there are exceptions. Thus atoms can be created and destroyed via nuclear processes. They consist of different forms called isotopes. Atoms are not the smallest particles, etc.

Chemical Periodicity

Look at the Periodic Chart.

Periods: The rows are referred to as periods. The pattern arises due to a repetition or periodicity of chemical properties.
- Combining ratios with hydrogen. 1 - 4 - 1: LiH, BeH₂, B₃, C₄, H₃N, H₂O, HF.
- Ions:
  1. Group IA always +1
  2. Group IIA always +2
  3. Group IIIA commonly +3, Al always+3
  4. Group IV +4 or -4 (usually covalent)
  5. Group VA commonly -3
  6. Group VIA commonly -2
  7. Group VIIA commonly -1
- Metal oxides basic combined with water (e.g. NaOH, KOH), non-metal oxides acidic combined with water (e.g. SO₂ and water gives H₂SO₃, one of the important acids in acid rain).
- IA & IIA oxides are basic, combining with water to give compounds such as NaOH and Mg(OH)₂, while VIA & VIIA oxides form acids with water such as H₂SO₄ and HClO₄. Intermediate groups show transitional behavior, e.g. carbon froms a weak acid, nitrogen strong acid, aluminum a weak base, phosphorous a weak acid, sulfur a strong acid.
- Properties go from metallic in Group IA to non-metallic in VIIIA.
Groups: The vertical columns of elements on the table sharing a family resemblance of properties (e.g. Li - Fr) are called groups. For example:
- IA (1) = alkali metals;
- IIA (2) = Alkaline earth metals;
- VIIA (17) = Halogens (note the generic symbol of X standing for any halogen);
- Group VIII (18) is known as the Noble Gases, or sometimes the Inert Gases because until the 1960's they had no known compounds. Very unreactive and only known compounds are with very reactive elements like F and O, and even they don't form compounds with smaller Noble gases such as He and Ne.
Note the numbering of the groups. The numbers from 1 - 18 are the internationally accepted numbers. We will also use the I - VIII "American" numbering system for the representative elements.

Terms etc.:

Representative elements: the elements of the s-block and p-block ("tallest" columns, blue and green on the table below).
Transition metal elements: the elements of the d-block (yellow in the table below).
Inner-transition metal elements: The Lanthanides and Actinides (not shown on the table below)

Periodic Table of the Elements

IA IIA IIIA IVA VA VIA VIIA VIIIA

H He

Li Be B C N O F Ne

Na Mg IIIB IVB VB VI VIIB VIIIB IB IIB Al Si P S Cl Ar

K Ca Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Ag Cd Sn I Xe

Cs Ba W Pt Au Hg Pb

**Periodic Table of the Elements**
IA	IIA		IIIA	IVA	VA	VIA	VIIA	VIIIA
H	He
Li	Be		B	C	N	O	F	Ne
Na	Mg	IIIB	IVB	VB	VI	VIIB	VIIIB	IB	IIB	Al	Si	P	S	Cl	Ar
K	Ca		Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn	Ga	Ge	As	Se	Br	Kr
Rb	Sr									Ag	Cd		Sn			I	Xe
Cs	Ba				W				Pt	Au	Hg		Pb

Chemical Nomenclature

Nomenclature is covered in section 2.8 (pp 57-67) in your textbook-you should be able to do the examples and exercises in the assigned problems. Note also the Discussion Module.

First, let's look at the the elements that you should learn the names of, as listed on the web:

http://www.humboldt.edu/~chem_dpt/resources/Supplements/element.html

The common ions and acids and bases are summarized on the handout and the web:

http://www.humboldt.edu/~chem_dpt/resources/Supplements/IonTable.html & http://www.humboldt.edu/~chem_dpt/resources/Supplements/AcidTable.html

Note that formulae are more or less written with the elements ordered by electronegativity (elements on the right side precede those on the left).

The Nuclear Atom

Atoms are now known to consist of three different types of particles: electrons, protons and neutrons (the common form of one very important atom, hydrogen, has only two kinds: a proton and an electron). The protons and neutrons reside in a small inner portion called the nucleus while the electrons reside in a relatively large cloud centered on the nucleus. Important properties of these particles are listed in the table below:

Particle	Charge	Relative Mass	Mass
Electron (e^-)	-1	1/1840	9.11 x 10^-28g
Proton (p or H⁺)	+1	ª1	1.67 x 10^-24g
Neutron (n)	0	ª1	1.67 x 10^-24g

Some important terms which you must know are:

Atomic number (Z) - the number of protons in the nucleus. This number is characteristic of a given element.
Atomic mass number (A) - the sum of the protons and neutrons in a given atom (p + n).
Atomic mass - the actual mass of an average atom in a sample. The characteristic atomic masses for Earth are shown on periodic tables.
Atomic Mass Unit: the atomic mass unit = amu is a unit of mass for atoms. It is defined as 1/12 the mass of one atom of ¹²C, where the mass of ¹²C is defined as 12 exactly.

Isotopes

Isotopes are forms of elements which differ only in the number of neutrons. This means different isotopes of the same element have essentially the same chemical properties but slightly different physical properties. They can also differ substantially in terms of their nuclear stability. Let's look at some examples of isotopes:

Symbol Z A p n e^-

¹⁴C 6 14 6 8 6

²³⁸U⁶⁺ 92 238 92 146 86

³⁵Cl^- 17 35 17 18 18

¹⁸O^2- 8 18 8 10 10

You should be able to fill in the blanks in a table like this with, the aid of a periodic table, on a quiz.

FYI - Classical Characterization and Descriptions of Atoms

During the latter half of the 19th century we began to actually characterize what these "atoms" might be like. There were two main lines of experimental work:

Spectroscopy: pure substances turned out to have a very complex spectra (which we will look at later), implying a complex substructure for atoms. Even the simplest atom, hydrogen has a spectra implying underlying structure as seen in the image of a hydrogen atom line spectra compared to a continuous spectra ("rainbow")
"Particle" experiments: a variety of phenomena could be investigated using electricity and radioactivity which indicated what type of substructure atoms might have. We will look at some of these experiments and their results now.

When two electrodes are placed in an evacuated tube and a high voltage placed across them a stream of negatively charged particles flows between the electrodes (from the negative cathode to the positive anode). Since they come off of the cathode they are called cathode rays, and the evacuated tube a cathode ray tube (the ancestor of todays TV or computer monitor cathode ray tube or "CRT"). These rays can be diverted with magnets or charged plates outside the tube. In fact by careful manipulation 19th century physicists were able to determine the ratio of the mass to the charge of the cathode rays. (The cathode ray is also known as the electron.)

In the early 20th century Milliken determined the charge on the electron by watching the rate of fall of tiny oil drops between two charged plates. The droplets had been ionized (charged) by exposure to x-rays. He noted that the droplets behaved like they had a multiple of some smallest charge on them, and determined this charge to be that of a single electron. Thus Milliken determined the charge on a single electron by making the assumption that the stepwise charge on the oil drops was due to single electron differences = -1.6 x 10^-19 Coulomb. With Thompson's determination of the charge/mass ratio for cathode rays (electrons), the mass of the electron could be calculated as = 9.11 x 10^-28g.

Goldstein used a modified Crook's tube with the cathode in the middle with a hole in the center and found positive rays, which he called "canal rays." These particles had positive charges in multiples of +1.6 x 10^-19 C, but variable charge to mass ratios! Thus matter consists of electrons and positively charged particles to make neutral matter.

Thompson proposed an atom based on this information (c. 1890) called the "plum pudding model" in which the atom is a positively charged mass with negative electrons distributed through it like raisins in a pudding.

However, this model was destroyed by the next great bit of experimental evidence. Rutherford used a source of a-rays and aimed them at a thin foil of metal (1911). By the plum pudding model he expected that these particles would be slowed or deflected. He was very surprised to find that nearly all of the particles passed through the foil unobstructed, while some were deflected a lot, even being essentially reflected back to the source!

So what does this mean for the description of the atom? It must consist of mostly empty space, with the positively charged portion confined to a very small space, a nucleus, in the center.

Syllabus / Schedule