1/8/09 Warm Up:

1. 1/8/09 Warm Up: • The observed regularities in the properties of the elements are periodic functions of their • Atomic numbers • Mass numbers • Oxidation states • Nonvalence electrons

2. Pop Quiz • Elements in Mendeleev’s periodic table were arranged according to their • Atomic number • Atomic mass • Relative activity • Relative size

3. 2. Most of the groups of the periodic table of elements contain • Nonmetals only • Metals only • Nonmetals and metals • Metals and metalloids

4. 3. An element is a gas at room temperature. It could be • A metal or a metalloid • A metal or a nonmetal • A metalloid or a nonmetal • A nonmetal only

5. 4. Atoms of metals tend to • Lose electrons and form negative ions • Lose electrons and form positive ions • Gain electrons and form negative ions • Gain electrons and form positive ions

6. 5. Which property is generally characteristic of metallic elements? • Low electrical conductivity • High heat conductivity • Existence as brittle solids • Existence as molecular solids

7. 6. When a potassium atom reacts with a bromine atom, the bromine atom will • Lose only 1 electron • Lose 2 electrons • Gain only 1 electron • Gain 2 electrons

8. 7. At room temperature, potassium is classified as • A metallic solid • A molecular solid • A network solid • an ionic solid

9. 8. Which list contains 2 metalloids • Ga, Ge, Sn • Si, P, S • C, Si, Ge • B, C, N

10. 9. Which element in period 4 is classified as an active nonmetal? • Ga • Ge • Br • Kr

11. 10. Which element in period 4 is classified as an active metal? • K • V • Br • Ge

12. Allotropes Buckyball • Some nonmetals can exist in 2 or more forms in the same phase EXAMPLES • Carbon • Graphite • Diamond • Buckminsterfullerene • Oxygen • O2 • O3 (Ozone)

13. Periodic Table Trends

14. Trends in Atomic Size • First problem: Where do you start measuring from? • The electron cloud doesn’t have a definite edge. • They get around this by measuring more than 1 atom at a time.

15. Atomic Size } • Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule. Radius

16. ALL Periodic Table Trends • Influenced by three factors: 1. Energy Level • Higher energy levels are further away from the nucleus. 2. Charge on nucleus (# protons) • More charge pulls electrons in closer. (+ and – attract each other) 3. Shielding effect

17. What do they influence? • Energy levels and Shielding have an effect on the GROUP • Nuclear charge has an effect on a PERIOD

18. #1. Atomic Size - Group trends H • As we increase the atomic number (or go down a group). . • each atom has another energy level, • so the atoms get bigger. Li Na K Rb

19. #1. Atomic Size - Period Trends • Going from left to right across a period, the size getssmaller. • Electrons are in the same energy level. • But, there is more nuclear charge. • Outermost electrons are pulled closer. Na Mg Al Si P S Cl Ar

20. Rb K Period 2 Na Li Atomic Radius (pm) Kr Ar Ne H Atomic Number 3 10

21. Ions • Some compounds are composed of particles called “ions” • An ion is an atom (or group of atoms) that has a positive or negative charge • Atoms are neutral because the number of protons equals electrons • Positive and negative ions are formed when electrons are transferred (lost or gained) between atoms

22. Ions • Metals tend to LOSE electrons, from their outer energy level • Sodium loses one: there are now more protons (11) than electrons (10), and thus a positively charged particle is formed = “cation” • The charge is written as a number followed by a plus sign: Na1+ • Now named a “sodium ion”

23. Ions • Nonmetals tend to GAIN one or more electrons • Chlorine will gain one electron • Protons (17) no longer equals the electrons (18), so a charge of -1 • Cl1- is re-named a “chloride ion” • Negative ions are called “anions”

24. #2. Trends in Ionization Energy • Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom). • Removing one electron makes a 1+ ion. • The energy required to remove only the first electron is called the first ionization energy.

25. Ionization Energy • The second ionization energy is the energy required to remove the second electron. • Always greater than first IE. • The third IE is the energy required to remove a third electron. • Greater than 1st or 2nd IE.

26. Shielding • The electron on the outermost energy level has to look through all the other energy levels to see the nucleus. • Second electron has same shielding, if it is in the same period

27. Ionization Energy - Group trends • As you go down a group, the first IE decreases because... • The electron is further away from the attraction of the nucleus, and • There is more shielding.

28. Ionization Energy - Period trends • All the atoms in the same period have the same energy level. • Same shielding. • But, increasing nuclear charge • So IE generally increases from left to right.

29. First Ionization energy Atomic number

30. Driving Forces • Full Energy Levels require lots of energy to remove their electrons. • Noble Gases have full orbitals. • Atoms behave in ways to try and achieve a noble gas configuration.

31. Trends in Ionic Size: Cations • Cations form by losing electrons. • Cations are smaller than the atom they came from – not only do they lose electrons, they lose an entire energy level. • Metals form cations. • Cations of representative elements have the noble gas configuration before them.

32. Ionic size: Anions • Anions form by gaining electrons. • Anions are bigger than the atom they came from – have the same energy level, but a greater area the nuclear charge needs to cover • Nonmetals form anions. • Anions of representative elements have the noble gas configuration after them.

33. Ion Group trends • Each step down a group is adding an energy level • Ions therefore get bigger as you go down, because of the additional energy level. Li1+ Na1+ K1+ Rb1+ Cs1+

34. Ion Period Trends • Across the period from left to right, the nuclear charge increases - so they get smaller. • Notice the energy level changes between anions and cations. N3- O2- F1- B3+ Li1+ Be2+ C4+

35. #3. Trends in Electronegativity • Electronegativity is the tendency for an atom to attract electrons to itself when it is chemically combined with another element. • They share the electron, but how equally do they share it? • An element with a big electronegativity means it pulls the electron towards itself strongly!

36. Electronegativity Group Trend • The further down a group, the farther the electron is away from the nucleus, plus the more electrons an atom has. • Thus, more willing to share. • Low electronegativity.

37. Electronegativity Period Trend • Metals are at the left of the table. • They let their electrons go easily • Thus, low electronegativity • At the right end are the nonmetals. • They want more electrons. • Try to take them away from others • High electronegativity.

38. The arrows indicate the trend: Ionization energy and Electronegativity INCREASE in these directions

39. Atomic size and Ionic size increase in these directions: