The Periodic Table

1. The Periodic Table Unit 4

2. I. History A. DmitirMendeleev • Russian chemist, 19th century •Arranged elements by their properties •Arranged by increasing atomic mass •Groups: vertical groups-elements have similar properties •Periods : horizontal rows •Periodic Law: Properties of the element are a periodic function of their atomic mass O Now arranged by atomic number O Iodine and tellurium were out of order

4. B. Henry Mosely • British physicist (1887-1915 years of accomplishment) • Developed the modern periodic table • Used x-rays to determine atomic number

5. II. Elements • Arranged based properties • 109elements-mostly naturally occurring • Any element greater than 83 is radioactive

6. III. Metals, Nonmetals, Metalloids A. Metals (H is NOT a metal) • Make up 2/3rds of the periodic table • Shiny • Solids (not Hg) • Malleable • Ductile • Good conductor of heat and electricity • Mobile electrons • Tend to lose electrons to become ions

8. B. Nonmetals • Not shiny • Gas, liquid, and solids • Not malleable or ductile • Brittle • Poor conductors of heat and electricity • Tend to gain electrons to become negative ions

10. C. Metalloids or semi-metals • In between in properties On stairs B, Si, Ge, As, Sb, and, Te

11. IV. Groups or Families A. Alkali Metals (group 1) • Very active metals, reactivity increases as you go down a group • React violently with water • Always found in compound

12. B. Alkaline Earth Metals (group 2) • Active but not as much as group 1 • Reactivity of metals increases as you go down a group

13. C. Transition Metals (group 3-12) •Can lose 1-3 electrons to become ions •Multiple oxidation states •High melting points O Hard solid at Standard Temperature and Pressure (STP: 0 C and 1 atm) •Mercury (Hg) is the exception-liquid •Form colored ions in solution • Reactivity of metals increases as you go down a group

14. D. Halogens (group 17) •“salt-formers” •Tend to bond with group 1 and 2 •Very active non-metals •Only group containing all three states of matter at room temperature

15. E. Noble Gases (group 18) •Non-reactive •Inert gases •8 valence electrons •Octet rule: all elements “want” 8 valance electrons •He exception- only 2 valence electrons

16. Q1. Explain the placement of an unknown element in the periodic table Q2. Compare and contrast metals, metalloids, nonmetals. Q3. Why are noble gases non-reactive

17. V. Trends A. Electron Configuration • Although arranged by atomic number, there are significant trendsfor electron configuration • Groups: same number of valance electrons (valence electrons determine how an element will react with other elements/ compounds) O Draw Li, Na, K •Periods: same number of principle energylevels O Draw Na, Mg, Al

18. •Lewis Dot Diagrams • Used to determine the type(s) of covalent bonds that an element may make in certain situations • Used to predict the type of ion that an atom might make when it forms an ion. • Each dot diagram consists of an elemental symbol, which represents the kernel of the atom, and a group of 1-8 dots which shows the configuration of the valence shell electrons (outer-most electron shell of the atom).

19. Order for placing dots, two dots can start on any side, continue either clockwise or counter clockwise, fill one dot at a time • Remember that each side can only hold up to two dots • The number of valance electrons can be determined using the group number

20. Q4. Explain how the number of energy levels containing electrons can be used to determine the Period the element would be found on the periodic table Q5. Draw a Lewis electron-dot structure for Na, Be, Al, Ne Q6. Distinguish between valence and non-valence electrons, given an electron configuration the following electron configurations. 2-1, 2-8-7, 2-8-18-8

21. B. Atomic radii (Size)-Table S • Measure of the size of the atom •Atomic radii measured as half the distance between 2 nuclei • r = ½ d •Groups: increase the number of principle energy levels as you go down • Radii increase •Period: same number of principle energy levels as you go across • Number of protons increases • More pull for electrons • Radii decrease

23. C. Ions •Octet Rule: all atoms in nature want to “look like” noble gases (8 valance electrons) •Metals • Few valence electrons • Loseelectrons become + ions •Nonmetals • Close to 8 • Want to gain electrons become – ions

24. D. Shielding •The electron shielding effect is the effect where core electrons block valence electrons from the nuclear charge of the nucleus. •If you increase the number of principle energy levels, shielding increases •Positive and negative charges attract each other so the more effective charge the electrons gets, the more attraction there is between the nucleus and the outer electrons. So as the effective nuclear charge increases, the atom and it's radii becomes smaller •As the shielding becomes stronger, the nuclear charge decreases and the size of the atom increases-More shielding, bigger atom

25. Q7. Explain the trends of periods, in terms of nuclear charge and electron shielding seen on the periodic table

26. E. Ionization Energy •Energy to remove the most loosely bound electron from a neutral gaseous atom •Trend in periods • From left to right, there is an increase in the number of protons which results in the nuclear charge increasing, the electrons are more strongly attracted and more energy is needed to remove them from the atom

27. •Trends in Group • Ionization energy decreases because valance electrons in each successive element are at a higher energy level and farther from the nucleus

29. F. Electronegativity • Electronegativity value of an atom is a measure of its attraction for electrons when bonded to another atom •Table S •Periods: from left to right shows an increase in electronegativity • Group: the highest electronegativity value is found at the top. Attraction for bonded electrons is less towards the bottom of the group