Unit 2 The Atom

2-1 Models of the Atom Over the years, the theories of the structure of the atom have changed, as new ideas and different experiments necessitated modifications to the model of the atom. A brief history of some of the major contributors follows:

Dalton an English scientist and teacher , proposed an early atomic theory: Elements are made of tiny particles called atoms. All atoms of a given element are identical. Atoms of one element combine with atoms of other elements to form compounds. A given compound always has the same relative # and types of atoms: Law of Constant Composition. Example: water always has 2 H atoms bonded to 1 O atom. Any other ratio yields a different chemical. In chemical reactions, atoms are neither created nor destroyed: Law of Conservation of Matter. Example: Determine the mass of Carbon in the following reaction:

16.0 g ? g 22.0 g O2 + C → CO2 16.0 + 6.0 = 22.0

Thomson: envisioned the atom as similar to an English dessert, plum pudding, with a uniform positive pudding containing negatively charged electrons (plums) spread throughout. This model was based on experiments with cathode ray tubes (pictured below).

Rutherford: : Ernest Rutherford and colleagues conducted a classic experiment that bombarded positively charged alpha particles at thin, gold foil to observe the deflections (pictured above right). Their results led to a nuclear atom: a dense, positive nucleus with a negative, surrounding, electron cloud.  

Bohr: To address the question regarding the Rutherford Model of the atom (why don’t the electrons get drawn into the nucleus), Niels Bohr proposed a model of quantized energy levels for the electrons. He thought that the electrons orbited the nucleus, much like the planets around the sun, and when energy was absorbed by the atom, the electron would make a “quantum leap” to higher energy levels. Modern quantum theory accepts quantization, but rejects the “orbits.”

Schrӧdinger: By treating the electron as a standing wave around the nucleus (below left), Erwin Schrӧdinger developed a mathematical equation that produces probability maps (“orbitals”) for the most likely regions of space where the electron is located [(h2/2m) 2 + v ]  = ih (/t)

A Summary Graphic

2-2 Parts of the Atom  

Atoms are made of: Protons Neutrons Electrons   Location: nucleus nucleus outside nucleus Charge: + neutral or zero -1 Mass: 1 amu (= 1.66 x 10-24 g) 1 amu ~0 amu (1/1947 amu)  

Atomic #: the number of protons; identifies the element !!!!!!!!!!

Isotopes different # of neutrons therefore: different masses!!!!!!!! atoms of same element (same # of protons) but with different # of neutrons therefore: different masses!!!!!!!!

Ions an atom that has gained or lost electron(s), resulting in a net + (positive) or – (negative) charged particles

Size atom size in on the order of 1 to 5 angstroms, where an angstrom (A) = 10-8 cm. 1 A is to 1 cm as 1 cm is to 600 miles! The nucleus is very small and very dense: it’s size is that of a grain of sand to a football field If the nucleus were the size of a pea, it would weight 250 million tons.