Atomic Structure Chapter 4
Section 1 Defining the Atom
Section 1 – Learning Targets 4.1.1 – I can describe Democritus’ ideas about atoms. 4.1.2 – I can explain Dalton’s atomic theory. 4.1.3 – I can identify what instrument is used to observe individual atoms.
Early Models of the Atom • Atom – smallest particle of an element that retains its identity in a chemical reaction.
Democritus’ Atomic Philosophy • Greek philosopher –first to suggest the atom. • Democritus believed that atoms were indivisible and indestructible. • He was not believed because he had no way to explain chemical behavior.
Dalton’s Atomic Theory • John Dalton (English 1766-1844) • Using experimental methods, Dalton transformed Democritus’ ideas of atoms into a scientific theory.
Dalton’s Atomic Theory • All elements are composed of tiny indivisible particles called atoms.
Dalton’s Atomic Theory • Atoms of the same element are identical. The atoms of any one element are different from those of any other element.
Dalton’s Atomic Theory • Atoms of different elements can physically mix together or can chemically combine in simple whole-number ratios to form compounds.
Dalton’s Atomic Theory • Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction.
Sizing Up the Atom • Despite their small size, individual atoms are observable with instruments such as a scanning tunneling microscope.
Section 2 Structure of the Nuclear Atom
Section 2 – Learning Targets 4.2.1 – I can identify three types of subatomic particles. 4.2.2 – I can describe the structure of atoms according to the Rutherford atomic model.
Subatomic Particles • Most of Dalton’s theory holds true today. • One change though – atoms are divisible. • Three kinds of subatomic particles are: electrons, protons, and neutrons.
Electrons • 1897 JJ Thomson (English 1856-1940) discovered the electron. • Electrons – negatively charged subatomic particle. • Cathode ray – glowing beam that travels from the cathode to anode.
The ray was deflected by negative charges and attracted to positive. • Thomson hypothesized the ray to be negative.
Robert A. Millikan (US 1868-1953) found the charge and mass of the electron. • In 1916 Millikan reported the negative charge and the mass of the electron to be 1/1840 the mass of a hydrogen atom.
Protons and Neutrons • Atoms have no net charge so positives and negatives equal out. • Eugen Goldstein (1850-1930 found canal rays (opposite of cathode rays) and concluded they were positive. • Protons – positively charged subatomic particle.
Protons are about 1840 times heavier than an electron. • 1932 James Chadwick (English 1891-1974) confirmed the neutron.
Neutron – subatomic particles with no charge but a mass nearly equal to that of the proton. • Protons and neutrons can be separated into smaller pieces called quarks.
The Atomic Nucleus • Most scientists thought that the protons and electrons were evenly distributed in the atom. • Ernest Rutherford (1871-1937) a student of Thomson proved otherwise.
Rutherford’s Gold-Foil Experiment • 1911- Rutherford and coworkers tested popular thought. • They bombarded a very thin sheet of gold foil with alpha particles (helium nucleus minus the electrons). • Thinking was that the particles would all go through with little deflection.
Actually almost all went through but some deflected (even straight back).
The Rutherford Atomic Model • Based on his results he modified the atomic theory. • Proposed that the atom is mostly empty space and most of the mass is in the center. • Nucleus – tiny central core of the atom and is composed of protons and neutrons.
In the nuclear atom, the protons and neutrons are located in the nucleus. The electrons are distributed around the nucleus and occupy almost all the volume of the atom.
Section 3 Distinguishing Among Atoms
Section 3 – Learning Targets 4.3.1 – I can explain what makes elements and isotopes different from each other. 4.3.2 – I can calculate the number of neutrons in an atom. 4.3.3 – I can calculate the atomic mass of an element. 4.3.4 – I can explain why chemists use the periodic table.
Atomic Number • Elements are different because they contain different numbers of protons. • Atomic number – number of protons in the nucleus of that element. • Remember atoms are neutral so protons equal electrons.
Mass Number • Mass number – the total number of protons and neutrons in an atom. • The number of neutrons in an atom is the difference between the mass number and atomic number.
Shorthand: gold-197 Mass # Atomic #
Isotopes • Isotope – atoms of the same element that have the same number of protons but different numbers of neutrons. • Because isotopes of an element have different numbers of neutrons, they also have different mass numbers.
Atomic Mass • Because the masses of atoms are so small a new unit was created. • Atomic mass unit (amu) – one twelfth the mass of a carbon-12 atom. • Or, the mass of a proton or neutron.
So why are the masses on the periodic table decimals? • Because not every isotope has the same abundance.
Atomic mass – a weighted average mass of the atoms in a naturally occurring sample of the element. • To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, then add the products.
Example: • Calculate the atomic mass of bromine. The two isotopes of bromine have atomic masses and relative abundance of 78.92 amu (50.69%) and 80.92 amu (49.31%).
The Periodic Table – A Preview • Periodic table – an arrangement of elements in which the elements are separated into groups based on a set of repeating patterns. • A periodic table allows you to easily compare the properties of one element (or group of elements) to another. • Elements in the same group have similar properties.