Atoms and Periodic Table Review

1. Atoms and Periodic Table Review

2. Dalton’s Concept • John Dalton, an English schoolteacher proposed the following ideas about matter: 1. Matter is made up of atoms. *2. Atoms cannot be divided into smaller pieces. *3. All the atoms of an element are exactly alike. 4.Different elements are made of different kinds of atoms. • Dalton pictured an atom as a hard sphere that was the same throughout.

3. Thomson’s Atomic Model • Thomson pictured a sphere of positive charge. • The negatively charged electrons were spread evenly among the positive charge. • The atom is neutral. • Plum pudding model (chocolate chip cookie)

4. Rutherford’s Gold Foil Experiment

5. Rutherford’s Model • Most of atom is empty space • Electrons surrounding • Nucleus- very dense positively charged area.

6. The Proton • In 1920 scientists identified the positive charges in the nucleus as protons. • A proton is a positively charged particle present in the nucleus of all atoms.

7. Neutron • Electrons are very light weight • Mass of atom should be equal to the number of protons. • But mass is about 2X the number of protons • Hypothesized a new particle. neutron (NEW trahn), would have the same mass as a proton and be electrically neutral.

8. The Electron Cloud Model • The new model of the atom allows for the somewhat unpredictable wave nature of electrons by defining a region where electrons are most likely to be found. • Electrons travel in a region surrounding the nucleus, which is called the electron cloud.

9. Atoms by the Numbers • Atomic Number- the number of protons in an atom. (Z) • The mass number of an isotope is the number of neutrons plus protons. Also called the atomic mass. Designated A

10. Isotope IdentificationName of element followed by mass number identifies the isotope. • Hyphen notation • Carbon-14 or C-14 • Mass number • Atomic number • 14 6 C C

11. Development of the Periodic Table– Mendeleev’s Table of Elements • A Russian chemist, Dmitri Mendeleev (men duh LAY uhf), published his first version of the periodic table 1869. • When Mendeleev arranged the elements in order of increasing atomic mass, he saw a pattern. • He noticed repeating patterns of properties.

12. Moseley’s Contribution • Although Mendeleev’s table correctly organized most of the elements, a few elements (cobalt and nickel, tellurium and iodine) seemed out of place. • In the early twentieth century, the English physicist Henry Moseley realized that Mendeleev’s table could be improved by arranging the elements according to atomic number rather than atomic mass.

13. Today’s Periodic Table • In the modern periodic table, the elements still are organized by increasing atomic number. • The rows or periods are labeled 1-7. • A period is a row of elements in the periodic table whose properties change gradually and predictably.

14. Today’s Periodic Table • The periodic table has 18 columns of elements. • Each column contains a group, or family, of elements. • A group contains elements that have similar physical or chemical properties.

15. Element Key Color- book only not reference sheet Blue – Metal Green – Metalloid Yellow - Nonmetal

16. Common elements - Uncommon Names • Sodium Na • Potassium K • Iron Fe • Copper Cu • Silver Ag • Tin Sn • Antimony Sb • Tungsten W • Gold Au • Mercury Hg • Lead Pb

17. Review Electron Configuration • The properties of an element are determined by the arrangement of electrons in the outer energy level. • The number of electrons in the outer energy level of main group elements can be determined by the group number of the element. • Group 1 and 2: number of electrons = group number. • Group 13-18: number of electrons = group number - 10

18. Groups • Group 1 – alkali metals 1 outer electron • Group 2 – alkaline earth metals - 2 outer electrons • Group 3-12 – Transition Elements (Medals) • Group 13 – 3 outer electrons • Group 14 – 4 outer electrons • Group 15 – 5 outer electrons • Group 16 – 6 outer electrons • Group 17 – halogens - 7 outer electrons • Group 18 – noble gases – 8 outer electrons