1. The Periodic Table
Chapter 5

2. Development of the Periodic Table
Section 1
Development of the Periodic Table

3. Objectives
Discuss how Dobereiner and Newlands helped to develop the periodic table
Evaluate the contributions of Mendeleev’s periodic table
State the periodic law

4. Why?
We have 115 currently known/synthesized elements, each with unique properties
Think about having a key ring with over 100 keys on it? Organization is needed!

5. Forerunners of the Periodic Table
Early 1800s: J. W. Dobereniner began organizing elements into sets of three
Sodium, lithium, potassium; chlorine, bromine, iodine

6. Forerunners of the Periodic Table
1865: J.A.R. Newlands noticed a repeating pattern present in every eighth element when arranged by atomic mass
Called the pattern the law of octaves; was scorned by his colleagues for the connection to music

7. Forerunners of the Periodic Table
1869: Dmitri Mendeleev and Lothar Meyer both publish a new classification system
Mendeleev’s system is better explained and more publicized, thus he tends to get the credit

8. Forerunners of the Periodic Table
Mendeleev followed Newlands’ lead and arranged elements by mass
Saw a periodic repetition of patterns
Thus, he called his table the periodic table of elements

9. Forerunners of the Periodic Table
Look at page 161
Compare Mendeleev’s periodic table to that of the table we have today. What are some similarities? What are some differences?

10. Forerunners of the Periodic Table
So why does Mendeleev get all the credit and not Newlands?
Mendeleev broke the pattern to put elements of like properties in the same column
Suggested that some masses were incorrect

11. Forerunners of the Periodic Table
Mendeleev also did something remarkable for the time: he predicted the properties of unknown elements with decent accuracy
Confirmed when germanium was discovered and fit the properties of Mendeleev’s ekasilicon

12. The Periodic Law Mendeleev still had some elements out of order
Did not know about atomic number
A student of Rutherford’s, H.G.J. Moseley realized the significance of the proton

13. The Periodic Law How was Mendeleev still correct?
Look at your periodic table. Compare mass trends to atomic number trends. What do you see?
Mendeleev accidentally found the pattern, despite not looking at the correct property

14. The Periodic Law The periodic law is the basis for the periodic table
“When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern.”
This would be a good thing to know!

15. Objectives in Review
How did Dobereiner and Newlands help develop the periodic table?
How did Mendeleev contribute to the modern periodic table?
State the periodic law

16. Reading The Periodic Table
Section 2
Reading The Periodic Table

17. Objectives
Explain why elements in the same group on the periodic table have similar properties
Distinguish between metals, nonmetals, and metalloids on the periodic table
Identify the location of the four blocks on the periodic table

18. Organizing the Squares
Our current periodic table has reached 118 squares, each representing a unique element
Aligned in vertical columns, called families or groups
Aligned in horizontal rows, called periods

19. Organizing the Squares
Notice that the number of elements in each period increases
How many elements in Period 1? 2? 4? 6?
Note that in Period 6, the rare earth metals are placed below the periodic table to conserve space

20. Labeling and Naming Groups
Three schemes to label families
We will use the International Union of Pure and Applied Chemistry (IUPAC) standard, numbering periods 1-18.

21. Labeling and Naming Groups
Some families also have names
Group 1 = Alkali
Group 2 = Alkaline Earth
Group 17 = Halogens
Group 18 = Noble Gases
More on that in the next chapter!
Yes Kyler, that is also where you get to see an explosion or two

22. Labeling and Naming Groups
Hydrogen is an oddball
Scientists argue about where it actually belongs; some put it in 17
Other scientists put helium in 2

23. Metals, Nonmetals, and Metalloids
Most of our elements are metals
Shiny
Good conductors of heat and electricity
Mostly solid at room temperature (Hg being an exception)
Malleable and ductile

24. Metals, Nonmetals, and Metalloids
The nonmetals are on the far right
Dull
Poor conductors of heat or electricity
Brittle
Can be gases, liquids, or solids at room temperature
Wide variety of physical properties

25. Metals, Nonmetals, and Metalloids
Metals on the left, nonmetals on the right (except H)
Metalloids occupy the dividing line
Also called semimetals
Located on the dense black line dividing the metals and nonmetals

26. Metals, Nonmetals, and Metalloids
Metalloids have properties of both metals and nonmetals
Typically brittle and dull, but can conduct electricity
Used as semiconductors

27. Electron Configurations
Remember that electrons in the highest energy level are the atom’s outermost electrons
Called valence electrons
Elements in a family have similar properties because they have similar configurations

28. Electron Configurations
For example, elements in Family 1 have 1 valence electron
This helps us shorten up our electron configurations!
Let’s take a look at how to abbreviate Family 1

29. Electron Configurations
H = 1s1
Li = [He]2s1
Na = [Ne]3s1
K = [Ar]4s1
Rb = [Kr]5s1
Cs = [Xe]6s1

30. The s-, p-, d-, and f-block Elements
Look at the shape of the periodic table
It can be divided up into the s-block on the left, the d-block on the middle, the p-block on the right, and the f block below
Look at the widths of each

31. Objectives in Review
Explain why elements in the same group on the periodic table have similar properties
Distinguish between metals, nonmetals, and metalloids on the periodic table
Identify the location of the four blocks on the periodic table

32. Section 3
Periodic Trends

33. Objectives Define the term periodic trend
Discuss four important periodic trends
Distinguish between ionization energy, electron affinity, and electronegativity

34. Periodic Trends
Many properties of the elements change in a predictable way as you move throughout the periodic table
Called a periodic trend

35. Atomic Radius
It may be helpful to you to draw mini periodic tables and draw arrows showing trends
The atomic radius is the distance from the center of the nucleus to the outermost electron of an atom
Not the most accurate measurement in the world…

36. Atomic Radius Two distinct trends
Atoms get larger moving from top to bottom
Atoms get smaller moving from left to right
See page 175

37. Atomic Radius As you go down, the more orbitals there are
So why does an atom get smaller as we go left to right…?
Remember that protons and electrons attract each other
Increased protons and electrons = more attractions

38. Ionic Size Similar trend as before…only now we have ions
If there are more protons than electrons, the size is greatly reduced
If there are more electrons than protons, the size increases
Why?
Look at page 176

39. Ionization Energy The energy needed to remove one electron
Measured in the number of joules needed to remove one electron
Atoms with high ionization energy hold onto electrons, while atoms with low ionization energies lose them easily

40. Ionization Energy
Typically represented as a group of atoms, known as the mole
1 mol = × 1023 atoms
Ionization energy is typically represented in kJ/mol

41. Ionization Energy Exact opposite trend for atomic radius
The closer to the nucleus the electrons are, the harder it is to pull them away

42. Successive Ionization Energies
What about to remove more than one electron?
Successive ionization energy is the energy required to remove electrons beyond the first
Take a look at the chart on page 180. What do you notice about the changes in energy?

43. Successive Ionization Energies
An atom holds strongly to the electrons in its noble gas inner core
Family 1 tends to form 1+ ions, Family 2 tends to form 2+ ions, and so forth
Elements on the right hand side do not want to give up electrons

44. Electron Affinity
The energy change that occurs when an atom gains an extra electron
Again, typically represented by kJ/mol
Suppose we have a mole of neon atom. It takes 29 kJ/mol of energy to add one electron to each atom

45. Electron Affinity Let’s look at another Fluorine takes -328 kJ/mol
Wait, negative?
Atoms with a negative electron affinity release energy to gain an electron

46. Electron Affinity Look at page 182
Not as defined of a pattern, but typically nonmetals have a higher electron affinity than metals do

47. Electron Affinity
Using ionization energies and electron affinities, we can formulate important principles about atoms
The octet rule states that atoms tend to gain, lose, or share electrons in order to gain a full set of valence electrons

48. Electron Affinity
For most energy levels, there are 8 valence electrons—2 in the s and 6 in the p orbital
H and He only have a maximum of 2
Elements on the right tend to become negative, elements on the left tend to become positive

49. Electronegativity
Electronegativity is the ability to attract electrons in a chemical bond
No units; it is relative to other atoms
Increases from left to right and from bottom to top
How does this relate to atomic radii?
What about the noble gases?

50. Objectives in Review Define the term periodic trend
What are the four important periodic trends?
Distinguish between ionization energy, electron affinity, and electronegativity