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Atomic History and Structure

Atomic History and Structure. Chapter 4. Early Theories of The Atom.

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Atomic History and Structure

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  1. Atomic History and Structure Chapter 4

  2. Early Theories of The Atom DemocritusDemocritus (b. c. 460 BC; d. c. 370 BC) postulated the existence of invisible atoms, characterized only by quantitative properties: size, shape, and motion. Imagine these atoms as indivisible spheres, the smallest pieces of an element that still behave like the entire chunk of matter. Thomson Rutherford Dalton “The PLUM PUDDING MODEL”

  3. Rutherford’s Experiment • Visit this website and click on the tutorial for “Section 3.2 Rutherford Experiment” Play and watch through all 5 parts • Visit this website and see what Rutherford’s Experiment looked like.

  4. What is an ATOM? • Atom- the smallest particle of an element that retains the properties of that element • Subatomic particles • Protons – Positive charge, found in the nucleus and have a mass of 1 amu. ( Identify) • Neutrons- No charge, found in the nucleus, and have a mass of 1 amu ( Isotopes) • Electrons- Negative charge, found in the energy levels outside of the nucleus, have relatively no mass ( Ions)

  5. By the Numbers… • Atomic number – • This determines the elements position in the periodic table ( Identifies the atom) • In atoms not ions • Atomic # = # protons= # of electron • Mass number- • Mass number = # protons + # neutrons • Why are electrons not included in the mass of an atom? • You will not look to the periodic table to determine the mass number, the number on the periodic table is an average.

  6. Isotopes- • An Isotope is the same element (at.#) with a different number of neutrons. • When naming an isotope you write name of the of the element, dash, then write the mass number • Example: Carbon-14 Symbol of element Atomic number from periodic table

  7. More about the mass • Knowing the mass number you can determine the number of neutrons in a specific atom • # of neutrons = mass number - # of protons • atomic mass unit (amu)- the mass of the carbon atom.

  8. Isotopic Notation • Isotope notation • Example #1 • Carbon-14 C • Carbon-12 C • Determine the number of protons, electrons and neutrons in an isotope. • Examples H H H • What number is different? Check your work!

  9. Calculating Atomic Mass • Atomic mass- weighted average mass of the isotopes of that element. • This is the decimal number on the periodic table. • To determine the atomic mass you must know what percent of each isotope of the element is found in nature and then it can be calculated. • Example: For Chlorine • 25% is chlorine-37 • 75% is chlorine – 35 • What is the average atomic mass of chlorine?

  10. To Calculate: • 25% is chlorine-37 • 75% is chlorine-35 • Take mass 37 x .25 = ans A • Take mass 35 x .75 = ans B • Add ansA + ansB = Average Atomic mass for Chlorine

  11. Ions-a charged particle • An atom that has either gained or lost an electron.Electrons are lost and gained to make ions • When they are gained (-Neg) ions (anion) • When they are lost (+ Pos) ions (cations)

  12. Ion Examples • Example: • What is the charge of an ion that has 11 protons and 10 electrons. Write the isotope notation for this atom. • Tell the number of P, E and N in the following ions. p=7 n=8 e=10 p=12 n=12 e=10

  13. History of Periodic Table • Dmitri Mendeleev – listed the elements in several vertical columns in order of increasing atomic MASS. He left blanks in the table for elements that were not discovered yet. • Henry Mosley – Arranged elements on the periodic table in order of increasing atomic NUMBER. http://www.chemheritage.org/EducationalServices/chemach/ppt/lm04.html

  14. Arrangement of Periodic Table • Periodic law – the properties of elements are periodic functions of their atomic number • Periods– horizontal rows; 7; correspond to energy levels • Groups/Families – vertical columns; “group A” Roman numerals correspond to the number of valence electrons • Group I – Alkali metals, Group II – Alkaline earth metals, Group VII – Halogens, Group VIII – Noble gases

  15. Metals, Nonmetals, and Metalloids • Metals – everything to the left of the stairstep; including aluminum; does not include hydrogen • Properties: Have luster (shiny), good conductors of heat and electricity, malleable ( able to be pounded into sheets), ductile (able to be pulled into a wire), tend to lose electrons in chemical reactions, most are solids • Transition metals – middle block over to stairstep • Inner transition metals – bottom 2 rows; sometimes called “lanthanide series” and “actinide series”

  16. Metals, Nonmetals, and Metalloids • Nonmetals – everything to the right of the stairstep; includes hydrogen • Properties – Dull, poor conductors, brittle, tend to gain or share electrons in chemical reactions, most are gases • Metalloids– either side of the stairstep; does not include aluminum

  17. Periodic Table colored to show metals, nonmetals, and metalloids http://hyperphysics.phy-astr.gsu.edu/hbase/pertab/pertab.html

  18. Answers to practice problems • Example • Isotopes of Hydrogen The mass number differs due to the difference in neutrons! Back to isotopic notation

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