1. The Periodic Table http://www.periodicvideos.com/videos/091.htm
Chapter 6
The Periodic Table

2. Section 1: History of the Periodic Table
Many scientist noticed patterns amongst elements when organized by their properties: A) Johann Dobereiner (early 1800’s)– recognized “triads” – groups of three elements with similar properties. Ex/ Ca, Sr, Ba; S, Se, Te; and Cl, Br, I B) John Newlands (1865) recognized rows of elements with similar properties repeated every 8th element; pattern called Law of Octaves (oct = 8)

3. Dmitri Mendeleev (1869)
“father” of the periodic table because his organization of elements (by atomic mass) at the time left gaps in the table for undiscovered elements at the time.
Ex/ ekaaluminum (eka = “one beyond”) temporary name for an element that Mendeleev predicted would be below aluminum (now known as Ga)

4. D) Henry Moseley (1914) His work led to organization of the periodic table by atomic number

5. The Periodic Law
States that an element’s physical and chemical properties are a function of their atomic number. Therefore, when elements are placed in order of atomic number, the elements with similar properties occur at regular intervals (groups!) **Note: Elements in a group are similar because of their similar valence electrons. Valence electrons found in the highest occupied energy level (outermost shell) of an atom; these electrons determine chemical reactivity.

6. Copy each question in notes
1)How many valence electrons are in the following elements: a) calcium b) sulfur c) bromine Hint: write the electron configuration. 1) a) Calcium: 1s22s22p63s23p64s2 Answer: 2 valence electron 1) b) sulfur: 1s22s22p63s23p4 Ans. 6 valence electrons 1) c) Ans. 7

7. 4) In which period and group is the element whose electron configuration is [Ne] 3s23p1 ?
Ans. Period 3 group 13
5) Given 1s22s22p63s23p64s23d10 4p5 Give the period and group numbers.
Ans. Period 4 and group 17
6) Write the outer electron configuration for a group 2 period 4 element.
Ans. 4s2

8. Class work: copy each question on back of Periodic Table & keep for discussion
1)How many valence electrons does potassium have?
2)How many valence electrons does chlorine have?
3)How many valence electrons does Kr have?
4) In which period and group is the element whose electron configuration is [Ar]4s2 ?
5) Given 1s22s22p63s23p64s23d9 . Give the period and group numbers.
6) Write the outer electron configuration for a group 1 period 7 element.

9. The Modern Periodic Table
Periodic Table: arrangement of elements in order of atomic number so that elements with similar properties fall in the same column, or group.
Using your text as a reference list notes in your notebook about the following groups and areas of the periodic table
Noble Gases, transition metals, inner-transition metals, representative elements, alkali metals, alkaline earth metals, halogens, metals, nonmetals, metalloids,
s-block, p-block, d-block, f-block, lanthanides, actinides,

10. Define each term in your notes & note location on Periodic Table
Periodic Table Noble Gases, transition metals, inner-transition metals, representative elements, alkali metals, alkaline earth metals, halogens, metals, nonmetals, metalloids, s-block, p-block, d-block, f-block, Lanthanides & actinides

11. Periods
What determines the length of each period on the periodic table? Answer: Determined by the number of electrons that can occupy the sublevels being filled in that period. sublevels in order Period # # of elements in period of filling: 1 2 1s 2 8 2s 2p 3 8 3s 3p s 3d 4p s 4d 5p s 4f 5d 6p s 5f 6d 7p

12. The Modern Periodic Table
Periodic Table: arrangement of elements in order of atomic number so that elements with similar properties fall in the same column, or group.

13. The s-Block (group 1) Group 1 are called Alkali metals
Extremely reactive metals (hence stored in oil)
Outermost energy level contains 1 electron in an s orbital
Soft; cuts easily with a knife
Not found in nature in elemental state, only in chemical bonds with other elements.
React with water to form hydrogen gas & a metallic hydroxide (alkalis)

14. The s-Block (group 2) Group 2 are called the alkaline-earths
Contain a pair of electrons in the s orbital (2)
Harder, denser, & stonger than group 1
Less reactive than alkali metals
Too reactive to be found as free elements

15. d-Block Elements: Groups 3-12
d-block are also known as transition metals
d sublevel first appears when n = 3
10 electrons fit in d-orbitals
Good conductors of electricity, high luster
Reactivity varies; but some are so unreactive that they exist free in nature (ex/ Platinum & Gold)

16. p-Block Elements: Groups 13-18
Consists of all elements in groups 13-18, except Helium
Electrons add to a p sublevel only after the s sublevel in the energy level is filled.
p-block contains all 3 element types (metal, nonmetal, and all 6 metalloids – B, Si, Ge, As, Sb, & Te)

17. p-Block Halogens - Group 17 Most reactive nonmetals.
React vigorously with metals to form salts ex/ sodium chloride NaCl – table salt
7 valence electrons

18. p-Block Noble Gases: group 18 1894-1900 all noble gases identified
Inert (non-reactive)
All have an octet (exception is He with a duet – 2 electrons) of electrons in their valence level (octet means 8 valence electrons)

19. f-block (below the P.T.) Lanthanides – Begins after La (Lanthanum);
14 elements below the table (part of f-block)
Atomic numbers 58 Cerium to 71 Lutetium
Period 6 (no group numbers assigned – between groups 3 & 4)
Actinides – Begins after Ac (Actinium); below the table too (also part of f-block); all radioactive
Atomic numbers 90 Thorium to 103 Lawrencium
Period 7 (no group numbers assigned – between groups 3 & 4)

20. Main-group or Representative Elements
s and p block together elements represent the main-group
Predictive chemistry for these blocks
Valence electrons are easily determined by their group number
Group number valence electrons

21. Periodicity – repeating patterns
Periodicity with respect to atomic number can be observed in any group of elements.
For example:
Group Group 18
Atomic number: Difference: at. #: diff.

22. Section 3 Electron Configuration and Periodic Properties(Trends)
Atomic radius one-half the distance of adjacent nuclei of identical atoms that are bonded together
(a measure of size of atoms)
refer to fig pg. 175
Period Trend: Atomic radius decreases L to R*
**The trend to smaller atoms across a period is caused by the increasing positive charge on the nucleus.
Group Trend: Atomic radius increases down a group

23. Example questions:
1. Which of the following elements has the largest atomic radius: Li, O, C, F? Ans. Li 2. Which has the smallest atomic radius: Li, K, Cs? Ans. Li

24. Atoms can either lose or gain electrons to become ions.
ion – an atom (or group of bonded atoms) that have a positive or negative charge.
Metals lose electrons to form positively charged ions called cations.
Ex/ Sodium atom: 11 protons & 11 electrons (neutral) Na
1s22s22p63s1
Sodium ion: 11 protons & 10 electrons (positive charge) Na+ 1s22s22p6 (stable electron configuration; like Ne)

25. Nonmetals gain electrons to form negatively charged ions called anions.
Ex/ Chlorine atom: 17 protons & 17 electrons (neutral) Cl
1s22s22p63s23p5
Chloride ion (name changes for neg ions add –ide):
17 protons & 18 electrons (negative charge) Cl-
1s22s22p63s23p6 (stable configuration like Ar)

26. Ionization Energy (IE)
IE – energy required to remove an electron from a neutral atom (kJ/mol) (or First Ionization Energy, IE1 )
fig in textbook
Period Trend: IE’s of main-group elements increases L to R across a period
Group Trend: IE’s of main-group elements decrease down a group
2nd and 3rd IE’s increase because each successive electron removed from an ion feels increasing stronger effective nuclear charge (pos. nucleus attracting neg. cloud) see p. 177

27. In each of the following pairs of elements, choose the element that has the higher first ionization energy. a) Ca and Ba b) Ca and Br c) Ca and K d) Ca and Mg Ans. a) Ca b) Br c) Ca d) Mg

28. Trend: Ionic Radii – size of ions
*same as Atomic Radius trends
fig pg. 179
Period Trend decreasing L to R across a period
Group Trend increasing down a group

29. Electron Affinity
EA – The energy change that occurs when an electron is acquired by a neutral atom (kJ/mol)
Period Trend - EA increases L to R across a period.
Group Trend - EA decreases down a group
Element with highest electron affinity is Fluorine
( kJ/mol). The higher neg. value the higher the affinity for electrons.

30. Electronegativity
EN - Measure of the ability of an atom in a chemical compound to attract electrons in a chemical bond.
0 to 4.0 scale (no units) Fluorine is the most electronegative!
table 6.2 pg 181
Period Trend – EN inc. L to R across a period
Group Trend – EN dec. down a group or remain about the same.