1. The Picture Worth a Thousand Words
The Periodic Table
The Picture Worth a Thousand Words

2. The Periodic Table
Function: organizational tool which classifies the over 100 elements in a systematic way; allows for easy comparison of the elements and provides substantial information by illustrating patterns Imagine reading the content of the periodic table in dictionary format

3. Organization of the Periodic Table
The periodic table visually classifies elements by multiple traits
number of protons (and electrons)
groups and periods
metals, nonmetals, and metalloids
main-group elements (representative elements), transition metals, inner transition metals
s, p, d, f blocks
valence electrons and electron configurations
periodic trends

4. Key to the Periodic Table
Every element is represented by a box containing the atomic number, the element symbol, the element name, and the atomic mass of the element

5. Overview of Periodic Table
Elements are arranged by increasing atomic number in rows known as periods and columns known as groups (or families)

6. Metals, metalloids, and nonmetals

7. Metals, Nonmetals, and Metalloids
Classification as metals, nonmetals or metalloids is based on specific properties
Most important to this classification is ability to conduct electricity or heat
Metals are good conductors
Nonmetals are poor conductors
Metalloids are intermediate in nature

8. Metals, nonmetals, and metalloids
Property
Metals
Metalloids
Nonmetals 1
good conductors of electric current and heat
ability to conduct electric current varies with temperature
poor conductors of electric current and heat 2
usually solid at room temperature (except Hg)
often gases at room temperature; when solids typically brittle 3
most are malleable
have properties of metals and nonmetals; when
have low boiling points 4
most metals are ductile
conductors they do not conduct heat or 5
reactivity varies; some highly reactive, others do not react easily
electricity as well as the
metals
reactivity varies; some highly reactive, some are extremely unreactive 6
usually have luster (shiny)
when solids typically dull-looking

9. Properties of Metals Good conductors of electricity and heat
burner on stove is made of metal
handles on cookware are not metal
Usually solids at room temperature (often hard)
exception is Hg (mercury) – liquid
Most are malleable – can be formed into sheets/foils
Most are ductile – can be drawn into wires
Reactivity varies (most react with acids)
some are highly reactive
some do not react easily
Usually shiny (luster)
Au (gold)
Ag (silver)

10. Properties of Nonmetals
Poor conductors of electricity and heat Often gases at room temperature Have low boiling points reason they are gases at room temperature Reactivity varies some are highly reactive – halogens some are highly unreactive – noble gases If solids they are typically dull-looking

11. Properties of Metalloids
Metalloids have properties of metals and nonmetals; their properties tend to fall between metals and nonmetals
do not conduct heat or electricity as well as the metals
properties vary widely
Ability to conduct electricity varies with temperature
semiconductors (used in electronics)
used in computer chips and solar cells

12. Metals Nonmetals Metalloids
Shiny dull solids varied
Solid liquid/gas solid Be Ca Ba C P S Si Ge Bi Os

13. Main Group Elements and Others
main group (representative elements): the s and p blocks
transition metals: the d block
inner transition metals: the f block

14. Main Group Elements most clearly illustrate periodic patterns
most abundant naturally occurring elements
most important elements for living things, including most common components of organisms: O, C, N, H, P, Ca

15. Transition Metals
Metals (with the properties of metals) Can have multiple oxidation states d orbitals allow them to behave differently, and to be less predictable Form complexes with water

16. Inner Transition Metals
Display very similar reactivity to each other Actinides are all radioactive unstable

17. s, p, d, and f Blocks

18. Special Groups
Group = column Group members typically have similar properties Four groups have special names: Alkali metals highly reactive Alkaline earth metals reactive Halogens highly reactive Noble gases highly unreactive

19. Special Groups A A L L K K N O I I B N M L E E E T G A R H S S T A H L

20. Versions of the Periodic Table
Updated with new elements
Updated with more accurate atomic masses
Presented with labels and color-coding
Different arrangements for different purposes

21. Versions of the Periodic Table
Different versions emphasize different features
They are all correct, and none is absolute

22. Origins of the Periodic Table
Mendeleev: created one of the earliest versions of the periodic table
Arranged elements by increasing mass
(right or wrong?)
Grouped elements with similar properties

23. Mendeleev’s 1872 Periodic Table
Predicted undiscovered elements and their properties

24. Mendeleev’s Contribution
Mendeleev developed Periodic Law (Modern) Periodic Law: the elements, when listed in order of their atomic numbers, fall into recurring groups, so that elements with similar properties occur at regular intervals Periodic law is observed by all members of a column having the same ending to their electron configuration

25. The Big Picture Periodic Table shows patterns
Multiple types of patterns can be found
Periodic repetition of chemical and physical properties of the elements as you move across a period (row)
Or down a group (column)
This trait of periodic repetition of the properties of elements when they are arranged by increasing atomic number is known as periodic law.

26. Electron Configurations
Pattern across a period and pattern down a group

27. Periodic Trends
Certain properties show a systematic variation across a period and down a group:
metallic character
valence electrons
atomic radii (atomic size)
ionic radii
electronegativity
ionization energy

28. Metallic Character

29. Valence Electrons
Alkali metals
Halogens

30. Valence Electrons
Equal number of valence electrons for all members of a group
For main-group elements (IA-VIIIA) the valence electrons are equal to the group number
Exception: He
(He only has 2 e-)

31. Atomic Radius
As protons are added across a period, the nucleus has a larger positive charge Na +11 Al +13 S +16 Ar +18 As electrons are added across a period, they continue to fill the same energy level, n The result? Larger effective nuclear charge on the electrons from left to right across a period

32. Atomic Radius
atomic radius increases
atomic radius increases

33. electronic structure of ion**
Ionic Radii
Na+
Mg2+
Al3+
P3-
S2-
Cl-
no of protons 11 12 13 15 16 17
electronic structure of ion**
2,8
2,8,8
ionic radius (nm)
0.102
0.072
0.054
0.212
0.184
0.181
**How many electrons does each have?
Na+ has 10, Mg2+ has 10, Al3+ has ?
Cl- has 18, S2- has 18, P3- has ?

34. Ionic Radius
The positive ions (Cations) Each ion has exactly the same electronic structure, and is equal to the structure of Ne, with n = 2. However, the number of protons for each ion is increasing. That causes a larger effective nuclear charge that pulls the electrons towards the nucleus and causes the ionic radii to decrease.

35. Ionic Radius
The negative ions (Anions) Each ion has the same electronic structure again, but this time equal to the structure of Ar with n = 3. Compared to the cations, there is an extra layer of electrons, so the ionic radii increase.

36. Ionic Radii
Cations become smaller than their atoms Anions become larger than their atoms

37. Electronegativity
Electronegativity: a measure of the ability of an atom in a compound to attract electrons More electronegative elements are more able to pull electrons from another atom towards their own nucleus

38. Electronegativity
Fluorine is the MOST electronegative element

39. Electronegativity
Think: Are Noble Gases electronegative? What is their electron configuration? How does this impact tendency to gain or lose electrons? Noble gases DO NOT follow the trend for electronegativity

40. Ionization Energy
Ionization energy: the energy required to remove an electron from an atom in the gaseous state Remember, atomic radius decreases across a period due to greater effective nuclear charge Electrons are closer to the nucleus and harder to remove
From top to bottom within a group, the valence electrons are in a higher energy, farther shell
This results in electron shielding, making it easier to remove electrons

41. Ionization Energy

42. First Ionization Energies

43. Ionization Energy
Atoms can lose more than one electron Always requires more energy to remove a second electron (increased effective nuclear charge) Peaks change element WHY?

44. Subsequent Ionization Energies
Each electron removed changes the peak

45. Electron Affinity
The energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion.
F(g) + e F-(g) Ho = kJ/mol
The negative sign of Ho indicates that energy is given off. This process is exothermic.

46. Electron Affinity
Electron affinity is more or less the opposite of ionization energy Halogens: greatest electron affinity Noble gases: lowest electron affinity (requires E)

47. Learning Check How do electron affinity and ionization energy compare?
F electron affinity
F ionization energy

48. Learning Check How do electron affinity and ionization energy compare?
F + e F electron affinity
F F+ + e ionization energy
Electron affinity is the energy required to gain an electron. Ionization energy is the energy required to lose an electron.