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Atomic Structure Atoms: The Building Blocks of Matter Early Atomic Theory 400 BC to the early 1800’s The Ancient Greeks • Democritus – the • Aristotle – particle theory succeeded of matter, Democritus and nature’s basic did not believe particle is the in atoms. He atom from the thought all Greek word matter was atomos which continuous and means made of a indivisible substance called hyle. Antoine Lavoisier – the Father of Chemistry • Lavoisier used the balance and measured everything which lead to the Law of conservation of Mass • He named oxygen which means “acid former” and helped determine that air is a mixture not an element John Dalton’s Atomic Theory 1808 • All matter is composed of small particles called atoms and empty space • All atoms of an element are identical • Atoms of different elements are different • Atoms combine in simple whole number ratios to form compounds Modern Atomic Theory The late 1800’s to Present J. J. Thompson in 1898 discovered the electron • The atom is a mass of positive charge which contains regions of negative charge (electrons) like the plums in plum pudding • Plum Pudding Model Rutherford, Geiger, and Marsden – discovered the nucleus • The center of an atom is dense and postively charged • The positive “nucleus” of the atom is very small compared to the total volume of the atom • Rutherford suggested that the electrons surround the nucleus like planets around the sun 4.2 The Atomic Nucleus •Rutherford’s Gold-Foil Experiment 4.2 The Atomic Nucleus •Alpha particles scatter from the gold foil. 4.2 The Atomic Nucleus • The Rutherford Atomic Model •Rutherford concluded that the atom is mostly empty space. All the positive charge and almost all of the mass are concentrated in a small region called the nucleus. •The nucleus is the tiny central core of an atom and is composed of protons and neutrons. Rutherford-Bohr Model Electron Proton Neutron 6.1 The Periodic Law •In the modern periodic table, elements are arranged in order of increasing atomic number. Atomic Number (Z) • Is the number of protons in the nucleus of an atom • Is the whole number in the periodic table • Is used the calculate the number of neutrons in an atom by its subtraction from the Mass Number 6.2 Squares in the Periodic Table 4.3 Mass Number The total number of protons and neutrons in a single atom of an element •Au is the chemical symbol for gold. 79 Protons 79 Electrons 118 Neutrons for Sample Problem 4.1 4.3 Atomic Mass •It is useful to compare the relative masses of atoms to a standard reference isotope. Carbon-12 is the standard reference isotope. Cabon12 has a mass of exactly 12 atomic mass units. •An atomic mass unit (amu) is defined as one twelfth of the mass of a carbon-12 atom. 4.3 Distinguishing Among Atoms •Just as apples come in different varieties, a chemical element can come in different “varieties” called isotopes. The Nucleus of the Atom • All elements are identified by their number of protons • Atoms with the same number of protons but different numbers of neutrons are called isotopes ISOTOPES Average Atomic Mass Number • The average of the total number of protons and neutrons in the nuclei of all the isotopes of an element • The decimal number in the periodic table Relative Abundance • Tells how common a particular isotope of an element is in nature • It is used to calculate the average atomic mass of an element 4.3 Average Atomic Mass •Some Elements and Their Isotopes 4.3 Average Atomic Mass •To calculate the average atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products. 4.3 Average Atomic Mass •For example, carbon has two stable isotopes: − Carbon-12, which has a natural abundance of 98.89%, and − Carbon-13, which has a natural abundance of 1.11%. for Conceptual Problem 4.3 Conceptual Problem