The Structure of the Atom 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”

All matter is made up of atoms Atoms of the same element are identical Atoms cannot be created, divided, or destroyed Atoms combine in certain reactions to form compounds In chemical reactions atoms are separated, combined, or rearranged

4.2 Subatomic Particles and the Nuclear Atom 1879-Crookes invented the cathode ray Cathode 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 Neutrons Rutherford 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 +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 Mass number

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 Electrons Neutrons 20 21 22 Symbolic Notation 39 K 40 K 41 K

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)

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