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Introduction

Learn how isotopes of an element differ from one another by their atomic and mass numbers. Explore the periodic table to compare properties, and understand the concept of atomic mass.

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Introduction

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Presentation Transcript

  1. Section 4.3 – Distinguishing Among Atoms Introduction • Just as apples come in different varieties, a chemical element can come in different “varieties” called isotopes. • In this section you will learn how one isotope of an element differs from another

  2. 4. Atomic Number • Elements are different because they contain different numbers of protons. • The atomic number of an element is the number of protons in the nucleus of an atom of that element.

  3. 4.3 5. Mass Number • The total number of protons and neutrons in an atom is called the mass number. • The number of neutrons in an atom is the difference between the mass number and atomic number. Au is the chemical symbol for gold.

  4. 6. The Periodic Table—A Preview • A periodic table is an arrangement of elements in which the elements are separated into groups based on a set of repeating properties. • A periodic table allows you to easily compare the properties of one element (or a group of elements) to another element (or group of elements). The Periodic Table

  5. The Periodic Table—A Preview (cont.) • Each horizontal row of the periodic table is called aperiod. • Within a given period, the properties of the elements vary as you move across it from element to element.

  6. The Periodic Table—A Preview (cont.) • Each vertical column of the periodic table is called a group, or family. • Elements within a group have similar chemical and physical properties.

  7. 4.3 7. Isotopes • Isotopes are atoms 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. • Despite these differences, isotopes are chemically alike because they have identical numbers of protons and electrons.

  8. 8. Atomic Mass • It is useful to compare the relative masses of atoms to a standard reference isotope. • Carbon-12 is the standard reference isotope. • Carbon-12 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. Some Elements and Their Isotopes

  9. Atomic Mass (cont.) • The atomic massof an element is a weighted average mass of the atoms in a naturally occurring sample of the element. • A weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature. • To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products. • For example, carbon has two stable isotopes: • C-12, which has a natural abundance of 98.89% • C-13, which has a natural abundance of 1.11%. Weighted Average Mass of a Chlorine Atom

  10. END OF SECTION 3

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