Chapter 2 Atomic Structure

Why do we study what matter is made of?
—better understanding, control, and use of matter

Law of conservation of mass
—In a physical or chemical change, mass is neither created nor destroyed. (However, this is not the case in nuclear reactions, which will be discussed in the next lecture.)

e.g., Charcoal + oxygen ? carbon dioxide + water + ash
The mass of charcoal + oxygen (consumed) = the mass of carbon dioxide + water + ash.

Law of constant composition
—The elements present in a compound have fixed proportions (ratios), regardless of the source

e.g.,
(a) 1.00 g H2O contains 0.11 g H and 0.89 g O;
thus, 18.0 g water contains 2.0 g H and 16 g O.

(b) 4.200 g vitamin C contains 1.720 g C, 0.190 g H, and 2.290 g O.
C% = 1.720 g/4.200 g = 0.4092 (= 40.92%)
H% = 0.190/4.200 = 0.0452 (= 4.52%)
O% = 2.290/4.200 = 0.5452 (54.52%)

Matter is composed of basic building blocks—atoms.
—by John Dalton, with experimental support which assists the understanding of chemical reactions at the atomic level.

Atomic theory on the basis of what proposed by Dalton:

1. An atom is the smallest particle that can be divided (by chemical means).
2. Atoms of the same element (e.g. isotopes) have virtually the same chemical properties (may have different physico-chemical and biochemical properties, such as reaction rates and bond strengths), and are different from those of other elements. (Different atoms have different "sizes" and different "weights".)
3. Atoms of different elements can combine (react) only in certain ratios in forming compounds, e.g., 1 O (1 O2) and 2 H (2 H2) to give 1 H2O (2 H2O); 6C, 12 H, and 6 O to give 1 glucose (C6H12O6).
4. Chemical reactions-union, separation, or/and rearrangement of atoms.
5. Atoms are not changed into other atoms by chemical reaction-remain their identities.

Atomic theory (http://antoine.fsu.umd.edu/chem/senese/101/atoms/dalton.shtml)
A New System of Chemical Philosophy (http://maple.lemoyne.edu/~giunta/dalton.html)
John Dalton (http://www.woodrow.org/teachers/chemistry/institutes/1992/Dalton.html)

How small is an atom? Can we "see" atoms?

Example: 1 g lead (Pb) contains 2.9 x 1021 Pb atoms

i.e., (1/207.6) x (6.02 x 1023)

We can "see" images of atoms by the use of scanning-tunneling microscope (STM), as shown below the image of gold atoms, and atomic force microscopes (AFMs).

Some other micro-images of chemical and biological interest using modern equipment can be seen at this link (http://www.di.com/Theater/NTMain.html)

Atomic masses

—Hydrogen was defined by Dalton to have a relative mass of 1.0. Then, Cl as a relative mass of 35.2 since the mass ratio of Cl/H in HCl is 97.24/2.76.

—A more accurate relative mass is used nowadays, in which 1/12 of the mass of the isotope C-12 is defined to be 1.00. Which is defined as 1.00 atomic mass unit (amu). Then, the relative mass of Cl is 35.45.

What is the exact masses of an atom? (Answer 1)

The structure of atoms
The subatomic particles of chemical significance
—electron, proton, neutron

electron-small, light (1/1840 lighter than hydrogen atom H), and negatively charged (–1), which controls the overall charge on and the chemical properties of an atom or a compounds (WHY? Answer 2).

proton-positively charged, 1840 times heavier than electron, and is the "acid matter" we are familiar with (i.e., H+—The hydrogen atom is composed of a proton and an electron.).

neutron-no charge with a mass equal to that of the proton (proton + electron)

nucleus of the atom-composed of protons + neutrons, and thus has a positive charge. The nuclear charge is neutralized by the electrons of the atom.

Atomic number-the number of protons of an atom (and is the "position" of an atom on the Periodic Table). Why not "the number of electron"… (answer 3)

mass number: The total number of protons and neutrons in the nucleus; thus, is the number of neutrons plus the atomic number.

Isotopes: The same element (thus the same atomic number) with the same (actually "nearly identical" but not exactly the same) chemical properties, but slightly different masses (different mass numbers or different number of neutrons).

e.g., (1) The "stable" H-1, the stable H-2 (in heavy water), and the radio-active H-3 (as a tracer in biochemical studies) are all hydrogen; (2) the stable P-31 and the radio-active P-32 (as a tracer/label/probe in biochemical research); (3) Co-59 and the radio-active Co-60 (used in cancer therapy) are the same element-cobalt; (4) the stable C-12, the NMR-active C-13, and the radio-active C-14 (for geological dating) are all carbon. These different atoms of a same element are called "isotopes".

Exercise

 Atomic number Mass number No. protons No. neutrons No. electrons Symbol 4 ___ ___ 5 ___ ___ ___ ___ 11 12 ___ ___ ___ 52 24 ___ ___ ___ 40 18 ___ ___ ___

Representing the composition of isotopes

(mass number)
element symbol
(atomic number)

For example:

1                 14             60
H                 C         Co
1                  6             27

for hydrogen-1, carbon-14, cobalt-60, respectively

Atomic mass of an element

—the "average" mass of the mixture of isotopes (based on the population of the isotopes).

e.g.,
atomic mass of H
= mass of H-1 (amu) x natural abundance (%) + mass of H-2 (amu) x natural abundance (%)
= 1.0078 x 0.99985 + 2.0141 x 0.00015 = 1.0079 (amu)

atomic mass of O
= mass O-16 (15.995 amu) x 99.759% + mass O-17 (16.995) x 0.037% + mass O-18 (17.999) x 0.204%
= 15.999 amu

How about an element which has three isotopes of masses X, Y, and Z amu and natural abundance of a, b, and c%? (answer 4)

(Continue to the discussion about electron configuration and the Periodic Table)