1. SOL Review 2 Atomic Structure And The Periodic Table

2. 1. John Dalton Dalton’s Atomic Theory 1.Atoms are tiny, indivisible particles 2.Atoms of the same element are identical 3.Atoms of different elements can combine in simple whole number ratios to form compounds 4.Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of a different element

3. 2. Joseph Thomson (Plum Pudding) JJ Thompson discovered the electron by passing an electric current through a cathode ray tube, and deflecting it with a magnet. With the discovery of the first subatomic particle, the atom was no longer “indivisible”. Scientists thought that the positive and negative particles inside an atom were evenly distributed, like in a “plum pudding”

4. 3. Ernest Rutherford – The Nuclear Atom Rutherford used the gold foil experiment to conclude that atoms are made up mostly of empty space, with a small dense nucleus that has a positive charge The electrons in Rutherford’s model orbit the nucleus in a “planetary” fashion

5. 4. Bohr’s Model Niels Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus. Each orbit corresponds to a specific energy level,

6. According to Bohr’s model, an electron will absorb a discrete quantum of energy to jump from one orbital to another orbital of higher energy When the electron falls back down to its ground state it gives off this energy, which we see as a distinct band of color that corresponds to the amount (quantum) of energy released

7. 5. Schrodinger Model Erwin Schrodinger came up with a mathematical equation that described the behavior of the electrons in a hydrogen atom. The quantum mechanical model of the atom comes from the solutions to this equation. This model does not give an exact path for electrons like Bohr’s model. Instead, the electrons occupy a cloud-like region around the nucleus.

8. Atomic Orbitals An orbital is a region of space in which there is a high probability of finding an electron These regions are assigned quantum numbers to identify their shape and energy level

9. Quantum Numbers 1st Quantum Number is the principal energy level 2 nd Quantum Number is the sublevel type, each with a different shape Sublevels: s, p, d, f 1 st level (n=1)s 2 nd level (n=2)s p 3 rd level (n=3)s p d 4 th level (n=4)s p d f

10. Aufbau Principle Lower energy levels fill first Hund’s Rule Every orbital in a subshell is filled with one electron before any one orbital is doubly occupied Pauli Exclusion Principle Each atomic orbital holds a maximum of two electrons and each electron must have opposite spin direction.

11. Order in which electron orbitals are filled

13. The Periodic Table Columns are groups ; rows are periods Metals o left side o high luster o good conductors o Most are ductile/malleable Nonmetals o right side o Non-lusterous o poor conductors o Most are gases at room temp. Metalloids o along the staircase; o properties of both

16. The period number of an element signifies the highest energy level an electron in that element occupies. The (representative) group number of an element signifies the number of valence electrons in the highest energy level.

17. Periodic Properties An element’s position is related to its atomic # and its electron configuration

18. Periodic Trends 1. Atomic Radius Atomic radius (size) decreases across a row/period Atomic radius increases down a column/group 2. Ionization Energy 1 st ionization energy – energy needed to remove the first electron from an atom Increases across row/period Decreases down a column/group 3.Electronegativity Tendency for atoms to attract electrons when combined w/ another element Follows the same trend as ionization energy: increases across a row, decreases down a group

19. Periodic Trends

20. Valence Electrons Valence electrons are the electrons in the highest occupied energy level of an atom. These are the electrons that determine the reactivity of an element. The number of valence electrons in the atom of an element is equal to its group number. Examples: oxygen (group 6A = 6 valence e - ) calcium (group 2A = 2 valence e - )

21. Electron Dot Diagram

22. Octet Rule OCTET RULE: Atoms of all elements tend to achieve the electron configuration of Noble gases by either losing or gaining electrons to have a complete octet of electrons in outer level. o Metals achieve octet by forming cations: losing their valence electrons to leaving a compete octet on a previous energy level. Metals become positively charged ions. o Nonmetals achieve an octet by forming anions: gaining electrons to complete an octet on the higher/outer energy level. Nonmetals become negatively charged ions.