Atomic Theories of Matter

2.1 The Atomic Theory of Matter

• Greek philosophers: Can matter be subdivided into fundamental particles?
• Democritus (460 – 370 BC): All matter can be divided into indivisible atomos.
• Dalton: Proposed atomic theory with the following postulates:
 Elements are composed of atoms.
 All atoms of an element are identical.
 In chemical reactions atoms are not changed into different types of atoms. Atoms are neither created not destroyed.
 Compounds are formed when atoms of elements combine.
• Atoms are the building blocks of matter.
• Law of constant composition: the relative kinds and numbers of atoms are constant for a given compound.
• Law of conservation of mass: during a chemical reaction, the total mass before reaction is equal to the total mass after reaction.
 Conservation means something can neither be created nor destroyed. Here it applies to matter (mass). Later we will apply it to energy (Chapter 5).

• Law of multiple proportions: if two elements A and B combine to form more than one compound, then the mass of B that combines with the mass of A is a ratio of small whole numbers.
• Dalton’s theory predicted the law of multiple proportions.


2.2 The Discovery of Atomic Structure

• By 1850 scientists knew that atoms were composed of charged particles.
• Subatomic particles: those particles that make up the atom,
• Recall the law of electrostatic attraction: like charges repel and opposite charges attract.

Cathode Rays and Electrons
• Cathode rays were first discovered in the mid- 1800s from studies of electrical discharge through partially evacuated tubes (cathode ray tubes or CRTs).
 Computer terminals were once popularly referred to as CRTs (cathode ray tubes).
 They are now commonly called VDTs (video display terminals)
• Cathode rays = radiation produced when high voltage is applied across the tube.
• The voltage causes negative particles to move from the negative electrode (cathode) to the positive electrode (anode).
• The path of the electrons can be altered by the presence of a magnetic field.
• Consider cathode rays leaving the positive electrode through a small hole.
 If they interact with a magnetic field perpendicular to an applied electric field, then the cathode rays can be deflected by different amounts.
 The amount of deflection also depends on the applied magnetic and electric fields.
 The amount of deflection also depends on the charge-to-mass ratio of the electron.
 In 1897 Thomson determined the charge-to-mass ratio of an electron.
• Charge-to-mass ratio: 1.76 x 108 C / g
• C is a symbol for coulomb
o SI unit of electric charge
• Millikan Oil-Drop Experiment
 Goal: find the charge on the electron to determine its mass.
 Oil drops were sprayed above a positively charged plate containing a small hole.
 As the oil drops fall through the hole they acquire a negative charge.
 Gravity forces the drops downward. The applied electric field forces the drops upward.
 When a drop is perfectly balanced, then the weight of the drop is equal to the electrostatic force of attraction between the drop and the positive plate.
 Millikan carried out the above experiment and determined the charges on the oil drops to be multiples of 1.60 x 10-19 C.
 He concluded that the charge on the electron must be 1.60 x 10-19 C.
• Knowing the charge-to-mass ratio of the electron, we can calculate the mass of the electron:

Radioactivity
• Radioactivity is the spontaneous emission of radiation.
• Consider the following experiment:
 A radioactive substance is placed in a lead shield containing a small hole so that a beam of radiation is emitted from the shield.
 The radiation is passed between two electrically charged plates and detected.
 Three spots are observed on the detector:
• A spot deflected in the direction of the positive plate.
• A spot that is not affected by the electric field.
• A spot deflected in the direction of the negative plate.
 A large deflection toward the positive plate corresponds to radiation that is negatively charged and of low mass. This is called -radiation (consists of electrons).
 No deflection corresponds to neutral radiation. This is called -radiation (similar to X-rays).
 A small deflection toward the negatively charged plate corresponds to high mass, positively charged radiation. This is called -radiation (positively charged core of a helium atom).
 X-rays and  radiation are true electromagnetic radiation, whereas - and -radiation are actually streams of particles – helium nuclei and electrons, respectively.

The Nuclear Atom
• The ‘plum pudding’ model: an early picture of the atom.
• The Thomson model pictures the atom as a sphere with small electrons embedded in a positively charged mass.
• Rutherford carried out the following experiment:
 A source of -particles was placed at the mouth of a circular detector.
• The -particles were shot through a piece of gold foil.
• Both the gold nucleus and the a-particle are positively charged, so they repel each other.
• Most of the -particles went straight through the foil without deflection.
• If the Thomson model of the atom was correct, then Rutherford’s results was impossible.
• Rutherford modified Thomson’s model as follows:
 Assume that the atom is spherical, but the positive charge must be located at the center.
 In order for the majority of -particles shot through a piece of foil to be undeflected, the majority of the atom must consist of empty space where the electrons can be found.
 To account for the small number of large deflections of the -particles, the center or nucleus of the atom must consist of a dense positive charge.



2.3 The Modern View of Atomic Structure

• The atom consists of positive, negative, and neutral entities (protons, electrons, and neutrons).
• Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the atom is due to the nucleus.
• Electrons are located outside of the nucleus. Most of the volume of the atom is the space where electrons are found.
• The quantity 1.602 x 10-19 C is called electronic charge. The charge on an electron is –1.602 x 10-19 C; the charge on a proton is +1.602 x 10-19 C.
• Masses are so small that we define the atomic mass unit, amu.
 1 amu = 1.66054 x 10-24 g
 The mass of a proton is 1.0073 amu, a neutron is 1.0087 amu, and an electron is 5.486 x 10-4 amu.
 The angstrom is a convenient non-SI unit of length used to denote atomic dimensions.
• Since most atoms have radii around 1 x 10-10 m, we define 1  = 1 x 10-10 m.

Isotopes, Atomic Numbers, and Mass Numbers
• Atomic number (Z) = number of protons in the nucleus
• Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons).
• By convention, for element X, we write
• Isotopes have the same Z but different A.
 There can be a variable number of neutrons for the same number of protons. Isotopes have the same number of protons but different numbers of neutrons.
 An atom of a specific isotope is called a nuclide.
 Examples: Nuclides of hydrogen include:
• H-1 (protium); H-2 (deuterium); H-3 (tritium): tritium is radioactive

CHEMISTRY The Central Science 8th Edition Brown, LeMay, Bursten Ch 2: Atoms, Molecules, and Ions