The Periodic Table

1. The Periodic Table • Antoine LaVoisier- the “father of chemistry” • By late 1700’s had compiled list of 23 elements • Chemical-based industry in 1800s greatly expanded the use and discovery of elements • In 1864, John Newlands: “Law of Octaves” • Chemical and physical properties repeat every 8 elements • Dmitri Mendeleev – Russian scientist • In 1872, presented the first periodic table of elements • The “father of the periodic table”

2. Mendeleev’s Periodic Table

3. Periodic Table Moseley: Noticed that the periodic table worked better if ordered by Atomic Number (Protons) rather than Atomic Mass Periodic Law: A predictable repetition of chemical and physical properties of the elements is found when arranged by atomic number.

5. Periodic Table Blocks

6. Periodic Brainteasers/Puns • What you do in a play? • What you do to a wrinkled shirt? • Policeman. • Superman’s weakness. • Your brother or mine • Name of a goofy convict • Not an exciting person • What a doctor does to his patients (3x).

7. Trends in Alkali Metals Francium is one of the rarest elements on Earth. It is the radioactive byproduct of the decay of Uranium and Actinium. It only has a half-life of 22 minutes, so it is very hard to analyze. How can we predict its physical properties?

8. Boiling Point Atomic number of Alkali Metal

9. Atomic Radius=1/2 distance between 2 nuclei

10. Atomic Radius Trends

11. Periodic Table Trends Element’s Atomic Radius determined by • Attraction of electrons by protons. • Repulsion between electrons. • Number of Principal Energy Shells. Net Effective Nuclear Charge: • Is the netpositivecharge experienced by electron in multi-electron atom. • Protons always added to nucleus, but electron positions can change as atomic number increases (energy levels). • Three factors control the net charge: size of energy levels, nuclear chargeand the shielding effect. • Thesedetermine trends in atomic properties

12. Shielding vs. Nuclear ChargeEffects on Periodic Table Trends • Shielding effect - inner electrons partially block outer electrons from the pull of the positively charged nucleus • The more principal energy levels, the more layers of inner electrons available to shield the valence electrons • Mainly affects properties of atoms going down a Group • s orbital e- can modestly shield p orbital e- on same level • Nuclear charge - attraction for all electrons by the positively charged nucleus • The higher the atomic number, the more protons in the nucleus, and the stronger the pull of the nucleus (greater nuclear charge) • Affects properties of atoms going left to right across a Period since those electrons are all on the same level and have a similar shielding effect.

13. Atomic Radius Trends • Across Periods: radius decreases as the increasing positive nuclear charge overwhelms the repulsive force of additional electrons. • Down Groups: radius increases as size of shells and shielding increase

15. Ions • An ionis an atom or bonded group of atoms with a positive or negative charge. • Atoms become ions when they either lose or gain electrons. Cats are Positive!

16. Ions • The size changes due to either greater repulsion of more electrons (anions), or increased protron attraction of fewer remaining electrons (cations).

17. Types of Ions and Effect on Radius • Cations = Positive Ions (atoms that have lost electron). Radius decreases. Ex: Ca+2 • Anions = Negative Ions (atoms that have gained electrons). Radius increases. Ex: O-2 Metals usually form cations. Non-metals usually form anions

18. Ionization energy is the energy required to remove an electron from an atom (form ion).

19. Ionization Energy: Increases across a Period • As you go across a period, one electron and one proton is added to the atom • But since the electron is added to the same energy level, the higher nuclear charge (from the additional protons in the nucleus) attracts the all of outer electrons more strongly. • Thus the energy required to remove an electron becomes larger across the period (left to right). Be B C N O F Li

20. Ionization Energy: Decreases down a group • As you move down a group, the size of the atom increases as it adds more energy levels. This increases the distance between the positive nucleus and electron. • This also causes a greater shielding effect, thus the electron is easierit is to remove. • The easier it is to remove an electron, the more reactive an element tends to be. Lithium Sodium Potassium Rubidium

22. Ionization Energy Trends

23. Electron Affinity X + e− → X− • Electron Affinity is effectively the opposite of ionization energy. • It is the energy released by an atom when it gains an electron (exothermic). • The halogens have the greatest electron affinity because gaining one electron gets them to noble gas configuration. • Electron Affinity generally increases left to right, but is erratic in behavior down groups..

24. Electron Affinities (kJ/mol) The first electron affinities of the group 17 elements

25. The electronegativityof an element indicates its relative ability to attract electrons in a chemical bond. Key to chemical bonding. • Fluorine is the most electronegative = 4

26. Overall Electronegativity Trend

27. NUMBER OF PROTONS INCREASES (Nuclear Charge Increases) • Radius DECREASES • Ionization energy INCREASES • Electron Affinity INCREASES Electronegativity INCREASES • LEVELS & SHIELDING INCREASE • Radius INCREASES • Ionization Energy DECREASES • Electronegativity DECREASES

28. The Octet Rule What makes atoms form certain ions? • Octet Rule: Atoms tend to gain, lose or share electrons in order to acquire a full set of eightvalence electrons. • Useful in determining what kind of ion an element is likely to form. They will add or lose electrons to get to s2p6. • Elements on the left side tend to lose electrons, elements on right tend to gain electrons to reach the noble gas configuration. • Na: 1s2 2s2 2p6 3s1 Obtains noble gas configuration (and a full valence octet) by losing an electron (Na+).

29. The Octet Rule • What are likely oxidation states (charges) of ions to form from the following? (HINT: add/lose fewest VALENCE electrons to get them to noble gas configuration) • Calcium: 1s2 2s2 2p6 3s2 3p64s2, thus Ca2+ ion • Potassium: • Chlorine: • Oxygen: • Phosphorus: • Zinc: • Iron: 1s2 2s2 2p6 3s2 3p6 4s2 3d6

31. Spiral Periodic Table