Chapter Two - Lesson Three.pptx

Lesson Three: Subatomic Particles: Protons, Neutrons, and Electrons in Atoms (2.6.) Finding Patterns: The Periodic Law and the Periodic Table (2.7.) Chapter 2 Atoms and Elements  Subatomi...

Lesson Three: Subatomic Particles: Protons, Neutrons, and Electrons in Atoms (2.6.) Finding Patterns: The Periodic Law and the Periodic Table (2.7.) Chapter 2 Atoms and Elements  Subatomic Particles (1 of 2) Discovery of Electron and Thomson’s Plum Pudding Model J.J. Thomson discovered the electron (1897) and proposed the Plum Pudding Model. Robert Millikan’s oil-drop experiment (1909) measured the electron’s charge. Rutherford’s Experiment and Nuclear Model Rutherford’s gold foil experiment (1909) led to the discovery of the nucleus. Proposed the Nuclear Model (or Planetary Model) with a dense, positively charged nucleus and surrounding electrons. Discovery of Proton and the Role of Goldstein Eugen Goldstein (1886) discovered canal rays, indicating positive charges. Rutherford (1917) identified the proton as a positively charged particle in the nucleus. Mass Controversy and Discovery of Neutron Atomic mass discrepancy suggested another particle. James Chadwick (1932) discovered the neutron, explaining the missing mass.  Subatomic Particles (2 of 2) All atoms are composed of the same subatomic particles: Protons, Neutrons, Electrons Protons and neutrons have nearly identical masses. The charge of the proton and the electron are equal in magnitude but opposite in sign. The neutron has no charge. Table 2.1 Subatomic Particles Particle Mass (kg) Mass (amu) Actual Charge (C) Relative Charge Proton 1.00727 +1 Neutron 1.00866 - 0 Electron 0.00055 Atomic Number (Z) The most important number to the identity of an atom is the number of protons in its nucleus. The number of protons defines the element because every element has a different number of protons. The number of protons in an atom’s nucleus is its atomic number and is given the symbol Z. Mass Number (A) Mass number is the sum of the number of neutrons and protons in an atom. It is represented by the symbol A. A = number of protons (p) + number of neutrons (n) A=p+n Example: Chlorine has an atomic number of 17. One of its atoms has a mass number of 35. How many neutrons does this chlorine atom contain?  Try This Carbon has an atomic number of 6. One of its atoms has a mass number of 14. How many neutrons does this carbon atom contain? A. 6 B. 8 C. 14 D. 20  Try the Following Determine the number of protons, electrons, and neutrons in the following atoms. Isotopes (1 of 2) All atoms of a given element have the same number of protons; however, they do not necessarily have the same number of neutrons. Atoms with the same number of protons but a different number of neutrons are called isotopes. For example, Neon has 3 isotopes: Ne-20, Ne-21, & Ne-22. – All the three equal number of protons and have the same atomic number (p = Z = 10) – They may contain 10, 11, or 12 neutrons. 20 10 Ne 10 21 Ne Ne 22 10 Isotopes (2 of 2) The relative amount of each different isotope in a naturally occurring sample of a given element is roughly constant. These percentages are called the natural abundance of the isotopes. Advances in mass spectrometry have allowed accurate measurements that reveal small but significant variations in the natural abundance of isotopes for many elements.  Conceptual Connection 2.6 Argon has an atomic number of 18 and has two common isotopes. One of these isotopes has a mass number of 40. How many neutrons does this isotope contain? A. 18 B. 20 C. 22 D. 58  Conceptual Connection 2.7 Carbon has two naturally occurring isotopes: C-12 (natural abundance is 98.93%) and C-13 (natural abundance is 1.07%). If circles represent protons and squares represent neutrons, which image best represents the C-13 isotope? Ions: Losing and Gaining Electrons The number of electrons in a neutral atom is equal to the number of protons in its nucleus (designated by its atomic number Z). During chemical change, however, atoms can lose or gain electrons and become charged particles called ions. – Positively charged ions, such as are called cations. – Negatively charged ions, such as are called anions.  Conceptual Connection 2.8 How many electrons are present in the O-2 anion? A. 6 B. 8 C. 10 D. 18 2.7. Finding Patterns: The Periodic Law and the Periodic Table In 1869, Mendeleev noticed that certain groups of elements had similar properties - when elements are listed in order of increasing mass, these similar properties recurred in a periodic pattern. – To be periodic means to exhibit a repeating pattern. Mendeleev summarized these observations in the periodic law: When the elements are arranged in order of increasing mass, certain sets of properties recur periodically. Mendeleev’s Periodic Table (1 of 2) Mendeleev organized the known elements in a table. He arranged the rows so that elements with similar properties fall in the same vertical columns. Mendeleev’s Periodic Table (2 of 2) Mendeleev’s table contained some gaps, which allowed him to predict the existence (and even the properties) of yet undiscovered elements. – Mendeleev predicted the existence of an element he called eka-silicon. – In 1886, eka-silicon was discovered by German chemist Clemens Winkler (1838–1904), who named it germanium. Modern Periodic Table (1 of 2) Mendeleev’s classification was based on atomic weight, leading to inaccuracies with isotopes and element placement. In the modern table, elements are listed in order of increasing atomic number rather than increasing relative mass. The modern periodic table also contains more elements than Mendeleev’s original table because more have been discovered since his time. Modern Periodic Table (2 of 2)  Modern Periodic Table (3 of 2) The periodic table can also be divided into – main-group elements, whose properties tend to be largely predictable based on their position in the periodic table. – transition elements or transition metals, whose properties tend to be less predictable based simply on their position in the periodic table. Modern Periodic Table (2 of 2) The periodic table has 18 vertical columns called groups and 7 horizontal rows called periods. Main-group elements are in columns 1A–8A (groups 1, 2, 13–18). Subgroup elements or transition metals are in columns 3–12 (B groups).  Classification of Elements (1 of 2) Elements are broadly classified as metals, nonmetals, and metalloids. Metals: Lower-left side and middle of the periodic table. Good conductors of heat and electricity, malleable, ductile, often shiny, tend to lose electrons (e.g., Chromium, copper, strontium, lead). Nonmetals: Upper-right side of the periodic table. Varied states at room temperature (solids, liquids, gases), poor conductors of heat and electricity, not ductile or malleable, tend to gain electrons (e.g., Oxygen, carbon, sulfur, bromine, iodine). Metalloids (Semimetals): Along the zigzag line dividing metals and nonmetals. Exhibit mixed properties of metals and nonmetals, often semiconductors with temperature-dependent conductivity.  Classification of Elements (1 of 2) Elements in the periodic table are also classified according to their groups as alkali metals, alkaline earth metals, boron families, inert metals, halogens, and noble gases. Alkali Metals (Group 1A): Highly reactive metals (e.g., Lithium, sodium, potassium, rubidium, cesium). Alkaline Earth Metals (Group 2A): Reactive, but less so than alkali metals (e.g., Magnesium, Calcium, Strontium, Barium). Halogens (Group 7A): Very reactive nonmetals (e.g., Fluorine, Chlorine, Bromine, Iodine). Noble Gases (Group 8A): Mostly unreactive, chemically stable (e.g., Helium, Neon, Argon, Krypton, Xenon).  Explain the following 1.Why are Group 7A elements called halogens? 2.Why are Group 6A elements called chalcogens? 3.Why are Group 8A elements called noble gases? 4.Why are Group 2A elements called alkaline earth metals? Ions and the Periodic Table (1 of 2) A main-group metal tends to lose electrons, forming a cation with the same number of electrons as the nearest noble gas. – Alkali metals (group 1A) lose one electron and form 1+ ions. – Alkaline earth metals (group 2A) lose two electrons and form 2+ ions. A main-group nonmetal tends to gain electrons, forming an anion with the same number of electrons as the nearest noble gas. – Halogens (group 7A) tend to gain one electron and form 1- ions – Chalcogens (group 6A) tend to gain two electrons and form 2- ions Ions and the Periodic Table (2 of 2)