Chapter Four

The Structure of the Atom

4.1 Early Theories of Matter

Philosophers believed matter was made of earth, water, air, and fire.

Democritus (460-370 BC) was first to propose that matter was made up

of atomos, which could not be further divided Atoms have different sizes and shapes giving

them different properties

John Dalton (1766-1844)

Dalton’s Atomic Theory-• Referred to the atom as

a “hard sphere”

1. All matter is made up of atoms2. Atoms of the same element are

identical3. Atoms cannot be created,

divided, or destroyed4. Atoms combine in certain

reactions to form compounds5. In chemical reactions atoms are

separated, combined, or rearranged

4.2 Subatomic Particles and the Nuclear Atom1879-Crookes invented the

cathode rayCathode Ray- a ray of radiation

that originates from the cathode and travels to the anode of a cathode ray tube

Led to invention of television

• By end of the 1800s scientists concluded that cathode rays were a stream of charged particles

• Particle carried a negative charge

• Electron- negatively charged particle

Evolution of the Atom

J.J. Thomson (1856-1940) determined mass-to-charge ratio of the electron

Determined that charged particle mass was less than that of smallest element, hydrogen

Meant that atoms were made of smaller particles, disproving Dalton’s theory

Created Plum-pudding model of the atom

Robert Millikan (1868-1953) determined that an electron has a negative charge

Thomson’s Plum-pudding or chocolate-chip cookie dough model of the atom

Proposed that negatively charged electrons (chips) were distributed through a “dough” of positive charge.

The Nuclear Atom

Ernest Rutherford (1871-1937) used a gold foil experiment to discover existence of nucleus

Nucleus- dense region in center of atom which is positively charged and contains virtually all of its mass

Used a gold foil experiment to see if positive alpha particles would be deflected by the electrons in the atom.

Since the positive charge was thought to be spread out, he thought it would not alter the path of the alpha particles.

Amazingly some were deflected at large angles which meant there must be a concentrated positive area.

Completing the Atom- The Discovery of Protons and NeutronsRutherford refined

concept of nucleus to include protons and neutrons

Bohr (1913)

Electrons orbit the nucleus

Orbitals have a set size and energy level

Lowest energy level is the smallest orbit

Defining the Atom

Atom- the smallest particle of an element that retains the properties of the element

How big is an atom?consider this:

world population in 2000: 6 000 000 000

# of atoms in a single copper penny:

29 000 000 000 000 000 000 000

Proton- subatomic particle carrying a positive charge

Electron -subatomic particle carrying a negative charge

Neutron- has a mass nearly equal to a proton, but carries no charge (neutral)

Protons, Electrons and Neutrons

Protons +ve charge

Electrons -ve Charge

Neutrons 0 charge

Mass

Protons 1.007316

Electrons 0.000549

Neutrons 1.008701

4.3 How Atoms Differ

• Atomic Number- number of protons in an atom

• Atomic number= # protons = # electrons

Mass of Individual Atoms

• Atoms have extremely small masses which are hard to work with, so scientists use a standard for comparison

• Standard used is a Carbon-12 atom

• Carbon-12 atom has mass of 12 atomic mass units

• 1 atomic mass unit (amu) is nearly equal to mass of 1 proton or 1 neutron

• Atomic mass is average mass of isotopes of element

Symbolic Notation

Mass number

atomic number = number of protons = number of electrons

Symbol

Isotope identification

Uranium – 238

Element name mass #

• Mass Number= # protons + # neutrons

• Isotope- atoms with same number of protons but different numbers of neutrons

Percent Abundance

• The chance that it will be found in nature

Average Atomic Mass

• The average atomic mass is equivalent to the most abundant isotope

Isotopes Potassium-39

Potassium-40

Potassium-41

Protons 19 19 19

Electrons 19 19 19

Neutrons 20 21 22

Symbolic Notation

39 K 19

40 K 19

41 K 19

Calculating Atomic Mass

• Isotopes of elements exist in nature in varying amounts

• Atomic Mass = sum of % abundance x atomic mass for each isotope

• Calculate the atomic mass unit of chlorine, whose percent abundance is 75% of chlorine-35, and 25% of chlorine -37.

• Ex: Chlorine:– Isotopes

• Chlorine-35 x (75%) = 26.25 amu• Chlorine-37 (25%) = 9.25 amu

–Atomic Mass = 35.5 amu (26.25 + 9.25 amu)

STOP!!!