1. The Periodic Table
Topic 5
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2. Searching For an Organizing Principle
Chlorine, bromine, and iodine have very similar chemical properties.
Chlorine, bromine, and iodine have very similar chemical properties. The numbers shown are the average atomic masses for these elements.

3. Mendeleev’s Periodic Table
6.1
Mendeleev’s Periodic Table
Mendeleev’s Periodic Table (1869)
How did Mendeleev organize his periodic table?
The Abbreviated History of the Periodic Table for Regents Chemistry

4. Click on pix for history

5. A. Dmitri Mendeleev (1869, Russian)
I. HISTORY
A. Dmitri Mendeleev (1869, Russian)
Organized elements by
increasing ATOMIC MASS.
Elements with similar chemical properties were grouped together.
There were some discrepancies.

6. B. Henry Moseley ORGANIZED ELEMENTS BY INCREASING ATOMIC NUMBER.
Resolved discrepancies in Mendeleev’s arrangement.

7. 6.1
The Periodic Law
In the modern periodic table, elements are arranged in order of increasing atomic number.
In the modern periodic table, the elements are arranged in order of increasing atomic number. Interpreting Diagrams How many elements are there in the second period?

8. When elements are arranged in order of INCREASING ATOMIC #, elements with similar chemical properties appear at regular intervals.

9. 6.1
The Periodic Law
The periodic law: When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties.
The properties of the elements within a period change as you move across a period from left to right.
The pattern of properties within a period repeats as you move from one period to the next.

10. Periodic Table of the Elements Intro
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11. A. Arrangement of Table II. ORGANIZATION OF THE ELEMENTS
1. Horizontal rows
Called PERIODS
All elements in the same period have the same number of ENERGY LEVELS in their atomic structure
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12. 2. Vertical Columns Called GROUPS OR FAMILIES
All elements in the same group have the
same number of VALENCE ELECTRONS, therefore lose or gain the SAME number of electrons, form similar CHEMICAL FORMULAS and have similar CHEMICAL PROPERTIES
ex. XCl Group 2:
Be +2 Cl -1 = BeCl 2
Mg +2 Cl = MgCl2
Remember: When writing formulas, use the criss-cross rule to cancel out oxidation states

13. III. Comparing Metals, Nonmetals & Metalloids
Elements on the Periodic Table are divided into three subgroups called METALS, NONMETALS and METALLOIDS (semimetals).
Increase nonmetallic properties
Increase metallic properties
Decrease metallic properties

14. METALS: located on the LEFT SIDE of the periodic table (except H); MORE THAN 2/3 of all elements
1. Chemical properties
tend to LOSE ELECTRONS EASILY
have LOW IONIZATION ENERGY (energy needed to remove electrons)
Metallic character INCREASES as ionization energy decreases.
have LOW ELECTRON AFFINITY (attraction for electrons)
form POSITIVE IONS when combining with other atoms
FRANCIUM most reactive metal: See Table J

15. 2. Metals Physical Properties
good conductors of heat and electricity
LUSTROUS - reflect light, shine when they are polished
MALLEABLE - can be rolled or hammered into sheets
DUCTILE - can be drawn into wires
are SOLIDS at room temperature except for MERCURY (liquid)

16. B. NONMETALS Chemical properties
located on the right side of the periodic table
Chemical properties
tend to GAIN electrons to form NEGATIVE IONS
have high electron affinities (electronegativity)
produce COVALENT bonds by SHARING electrons with other nonmetals
FLUORINE most reactive nonmetal: see Table J

17. 2. Nonmetals Physical Properties
exist as gases, molecular solids, or network solids at room temperature except BROMINE (liquid)
BRITTLE - (shatters when struck)
DULL - does not reflect light even when polished
POOR CONDUCTORS of heat and electricity
Allotropes: Different SHAPE & PROPERTIES forms from the same element.
CARBON: coal; diamond, graphite
OXYGEN: O2; O3 (OZONE)

18. C. METALLOIDS
Found lying on the jagged line between metals and nonmetals flatly touching the line (except Al and Po).
B,Si,Ge,As, Sb, & Te
Exhibit properties of both metals and nonmetals
Behave as nonmetals but their conductivity is like metals
SEMICONDUCTORS – Si and Ge

19. IV. Periodic Trends A. Periodic Law
When elements are arranged in order of increasing atomic #, elements with similar properties appear at regular intervals.

20. 1) Ionization Energy
Energy needed to remove the most loosely bound electron from a neutral gaseous atom
X + energy X+ + e-

21. Trends in Ionization Energy
6.3
Trends in Ionization Energy
First ionization energy tends to increase from left to right across a period and decrease from top to bottom within a group. Predicting Which element would have the larger first ionization energy—an alkali metal in period 2 or an alkali metal in period 4?

22. Trends in Ionization Energy
IE increases as you move across a period
Why?
The nuclear charge (atomic #) is increasing therefore greater attraction of the nucleus for electrons hence harder to remove an electron

23. Trends in Ionization Energy
IE decreases as you move down a group
Why?
Atom size increases making the outermost electron farther away from the nucleus therefore making it easier to remove
Shielding increases

24. Ionization Energy cont.
Why opposite of atomic radius?
In small atoms, e- are close to the nucleus where the attraction is stronger
Why small jumps within each group?
Stable e- configurations don’t want to
lose e-

25. 3. Ionization Energy – Table S
First Ionization Energy
Increases UP and to the RIGHT
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26. Ionization Energy cont.
Successive Ionization Energies
Large jump in I.E. occurs when a CORE e- is removed.
Mg 1st I.E kJ
2nd I.E. 1,445 kJ
Core e- 3rd I.E. 7,730 kJ

27. 2. Atomic Radius – Table S ½ the distance between nuclei
© 1998 LOGAL
½ the distance between nuclei
Increases DOWN
Decreases Across
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28. Atomic Radius cont. Why is it smaller to the right?
Increased nuclear charge(# p+) without additional shielding pulls e- in tighter
Why is it larger going down?
Higher energy levels have larger orbitals
Shielding - core e- block the attraction between the nucleus and the valence e-

29. 3. Electronegativity – Table S
Electronegativity- ability for an atom to attract electrons
Based on a scale of 4, Fluorine having the greatest electron affinity
A. Metals
lose e-
Form Cations (+)
get smaller
© 2002 Prentice-Hall, Inc.
B. Nonmetals
gain e-
Form Anions (–)
Get larger
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30. Trends in Electronegativity
6.3
Trends in Electronegativity
Representative Elements in Groups 1A through 7A

31. 4. Melting/Boiling Point – Table S
Highest in the middle of a period.

32. Periodic Trends Summary (use reference Table S for data comparison)
Across a period
Up a group
Ionization energy
increases
Electronegativity
Atomic radii
decreases
Metallic properties
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33. IV. Classification Alkali Metals Alkaline Earth Metals
Transition Metals
Halogens
Noble Gases
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34. Group 1: Alkali Metals extremely reactive (not found free in nature)
form stable ionic compounds
react with water to form a base
react with air to form oxides
react with acids to form salts
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35. Group 2: Alkaline Earth Metals
reactive (not found free in nature) - form stable ionic compounds
react with water to form a base
react with air to form oxides
react with acids to form salts
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36. Groups 3-11: Transition Metals
multiple positive oxidation states
Lose electrons from two outermost energy levels
Ions form colored solutions

37. Group 15 – unique features
Members range from typical nonmetals (nitrogen and phosphorus) through metalloids (arsenic and antimony) to metals (bismuth)
Nitrogen
Forms stable diatomic molecules with a triple bond
Component of protein
Forms some unstable compounds that are used as explosives
Phosphorus
Component of nucleic acids (DNA, RNA)
More reactive than nitrogen at room temperature

38. Group 16 – unique Features
Members range from typical nonmetals (oxygen and sulfur) through metalloids (selenium and tellurium) to metals (polonium)
Solids except oxygen
Oxygen can exist as O2 and O3 (it is an allotrope)
Polonium is radioactive

39. Group 17: Halogens very reactive nonmetals - high electronegativity
not found free in nature
form diatomic molecules when free
react with metals to form salts (halides)
Found in all three phases (s, l, g) due to differences in Van der Waals forces (these are weak)

40. Group 18: Noble Gases Have complete outer shells
Almost inert (not reactive); stable
Krypton, xenon, and radon form compounds with oxygen and fluorine
Referred to as monatomic gases

41. TODAY……..
1. 1. Assemble into your groups according to Take your notes and 5 questions with you. Using your notes, discuss & complete the Teachback WS questions pertaining to your group.
Use the post-it notes for any questions you may have THAT MAY NEED ANSWERING

42. Once finished, begin RB questions 46-90.
3. BEGIN your teachback, this NOT a silent lesson – use your low voices to teach the other members. Ask your questions to ascertain clarity. At the end of this session, everyone in the group should have the Teachback WS completed.
Use the post-it notes for any questions you may have THAT NEED ANSWERING..
Once finished, begin RB questions
Cooperation and diligence is necessary…..
I will be watching and listening.

43. TEACHBACK PROJECT REVIEW
Alkali Metals
Alkaline Earth Metals
Transition Metals
Halogens
Noble Gases
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44. GROUP 1 1. The name of this group is ALKALI METALS
2. Does this group contain metals or nonmetals? METALS
3. Alkali metals (lose or gain)l LOSE electrons becoming (positive or negative) ions? POSITIVE
4. What happens to the reactivity of the elements in this group as the atomic number increases. (increases or decreases) INCREASES
4. Are they (more or less) MORE reactive than all of the elements in Group 2 and why? THEY HAVE LOW IONIZATION ENERGIES

45. Group 1 continued:
5. Can these compounds be found in nature in the elemental or combined state? COMBINED STATE IN THE FORM OF A SALT
6. What type of compounds do they normally form (ionic or covalent)? IONIC (M + NM)
7. If element Y represents an alkali metal, what is it’s general formula for the reaction with a:
Chloride: YCl Oxide: Y2O
8. What is the most reactive metal in this group? FRANCIUM OR CESIUM

46. GROUP 2 1. The name of this group is ALKALINE EARTH METALS.
2. Does this group contain metals or nonmetals? METALS
3. They (lose or gain) LOSE electrons & form (positive or negative) ions POSITIVE ?
4. Describe the reactivity of the elements in this group as the atomic number increases (increases or decreases) INCREASES
5. Are they (more or less) LESS reactive than all of the elements in Group 1 and why? _________________________________________

47. Group 2 continued:
5. Can these compounds be found in nature in the elemental or combined state? COMBINED STATE IN THE FORM OF A SALT
6. What type of compounds do they normally form (ionic or covalent)? IONIC (M + NM)
7. If element Z represents an alkaline earth metal, what is it’s general formula for the reaction with a:
Chloride ZCl Oxide: ZO

48. GROUP 15 1. What is this group referred to as? NITROGEN GROUP
2. Name the diatomic element in this group? NITROGEN
3. Classify each element in this group as a metal, non-metal or semi-metal (metalloid).
metals: BISMUTH
nonmetals: NITROGEN, PHOSPHORUS
semi-metal: ARSENIC. ANTIMONY
4. What happens to the reactivity of a non-metal as the atomic number increases (increases or decrease) DECREASES .
5. Which is the most reactive non-metal in this group? PHOSPHORUS
6. Is nitrogen a (diatomic or monatomic) DIATOMIC molecule & what type of bond is found in nitrogen? TRIPLE COVALENT

49. GROUP 16
1. Classify each element in this group as a metal, non-metal or semi-metal.
metals: POLONIUM nonmetals: OXYGEN, SULFUR, SELENIUM semi-metal: TELLERIUM
2. Name the diatomic element in this group. OXYGEN
3. Define an allotrope? DIFFERENT FORMS OF AN ELEMENT IN THE SAME PHASE WITH HAVING DIFF CHEM & PHYSICAL PROPERTIES
4. Which element(s) in this group is an allotrope? PHOSPHORUS, SULFUR, OXYGEN (O2, O3)
5. What type of element is Polonium? RADIOACTIVE METAL

50. TRANSITION ELEMENTS (groups 3B-12)
Which element is a liquid at room temperature? MERCURY (Hg)
2. What are the four main characteristic chemical properties of transition elements?
MULTIPLE POSITIVE OXIDATION STATES
IONS FORM COLORED SOLUTIONS
LOSE ELECTRONS FROM TWO OUTERMOST ENERGY LEVELS
UNFILLED D ORBITALS

51. GROUP 17 1. The name of this group is HALOGENS
2. Classify the elements in this group: NONMETALS
3. Halogens (lose or gain) GAIN electrons becoming (positive or negative) ions? NEGATIVE
4. Why is astatine not included much in these discussions? NOT ENOUGH AVAILABLE TO STUDY
5. What would the general formula of a Group 17 element (represented by X) combined with magnesium of
group 2? MgX2

52. Group 17 continued:
6. What is the most reactive element in this group? FLUORINE
7. Can these compounds be found in nature in the elemental or combined state? COMBINED STATE AS SALTS
8. What type of salts do these elements form? HALIDES
9. For each state of matter, list the element(s) in this group.
solid: IODINE liquid: BROMINE gas: CHLORINE, FLUORINE
10. What type of forces of attractions account for the different states of matter that exist and the high MP’s and BP’s as you go down the group? VAN DER WAALS FORCES
(weak forces that get stronger as you go down the group )

53. GROUP 18 The name of this group is called NOBLE GASES.
What type of molecules do these gases form? (monatomic or diatomic) MONATOMIC
Describe the electron arrangement in the outermost energy level of all these elements. STABLE OCTECT – INERT GAS STRUCTURE
Which element has only two electrons? HELIUM
Describe the reactivity of the elements in this group. THEY ARE UNREACTIVE (Kr and Xe can be forced to react with F in lab settings)