1. Periodic Table
Chapter 5

2. Organizing the elements
Section 1

3. Let’s Review! Chemical properties
Any property that can only be tested by changing the chemical make-up of the substance.
Physical properties
Any property that can be tested without changing the chemical make-up of the substance
Atomic mass
Mass of protons and neutrons
Atomic number
Unique to each element, same as number of protons

4. Dmitri Mendeleev 1870: 63 elements known to man
He organized them in order of their atomic mass, and saw a pattern from their properties.
Was working on this while Thomson and Rutherford were still “exploring” the atom

5. Dmitri Mendeleev
Arranged his table with repeating properties in columns, starting a new row each time the chemical properties repeated
Left blank spaces in his table, concluding that these spaces were elements that hadn’t been discovered yet.
Based on the patterns and the other elements around the blank space, he predicted the properties of those elements

6. What he called ekasilicon – it was discovered a few years later
An Example
What he called ekasilicon – it was discovered a few years later
Prediction
Atomic Mass
72 amu
Density
5.5 g/mL
Appearance
dark gray metal
Melting Point
high melting point
Germanium
72.6 amu
5.3 g/mL
gray metal
937o C

7. Mendeleev’s Table

8. Henry Moseley
Mendeleev’s table worked because as protons increase, atomic mass should increase, but if there are fewer neutrons it could decrease
Errors arose because he was arranging the table with the wrong number
1910: Discovered atomic number and rearranged the periodic table using this number, it fell into perfect order

9. Periodic Law
Periodic Law: physical and chemical properties of the elements are periodic functions of their atomic numbers
In other words, when the elements are arranged by their atomic numbers, you should see chemical and physical properties repeating themselves

10. Rows Left to right – called periods
Elements in the same periods show patterns left to right - conductivity/reactivity change, elements become less metallic
The period # indicates the number of energy levels each atom in the row has

11. Columns Top to bottom – called groups
Elements in a group have similar chemical properties and show trends top to bottom
The elements in the same group (column) have the same number of valence electrons

12. Exploring the periodic table
Section 2

13. The periodic table is organized by atomic number
Remember
The periodic table is organized by atomic number
For a neutral atom, the number of protons equals the number of electrons

14. Valence Electrons
The trends found in the periodic table are a result of electron arrangement, specifically, the number of valence electrons
Valence Electron: electrons in the outermost energy level

15. Valence Electrons
The group number of an element will tell you the number of valence electrons it has
Group 1: 1 valence electron
Group 2: 2 valence e- ’s
Skip 3-12
Group 13: 3 valence e- ’s
Groups 14-18: 4, 5, 6, 7, and 8 valence e- ’s respectively.

17. Ion
When a neutral atom gains/loses electron(s) through bonding, the atom is no longer neutral and has an excess charge
It becomes an ion
Ion: a charged atom

18. Ion
All atoms want 8 valence electrons in the outer shell – this makes the shell full and the atom stable
Elements close to having 8 tend to be the most reactive.
Elements already “full” are considered inert, they don’t react because they don’t need to
The number of electrons an atom can gain or lose is equal to the number of valence electrons it has

19. Let’s Practice Group 16 Give Up? or Gain? Group 13 Give Up? or Gain?

20. Ion Protons = positive charge Electrons = negative charge
p+ # CANNOT change, but e- # can
So…
If an atom GAINS electrons, is it more positive, or negative?
If they LOST electrons?

21. ION Cation: atoms that LOSE electrons, becoming more positive
Anion: atoms that GAIN electrons, becoming more negative

22. ion How do we know if an atom is an ion?
Cations have a +, and anions have a – superscript
If an atom has gained 3 electrons
It has 3 MORE negative particles than positive particles, it is more negative = Al3-
If an atom has lost 3 electrons
It has 3 LESS negative particles than positive particles, it is more positive = Al3+

23. The Periodic Table
Section 3

24. The Periodic Table
Divided into three major categories based on general properties
Metals, Nonmetals, Metalloids (semiconductors)

25. Metals Like to give up valence electrons
Physical Properties: high luster (shiny), conductive (heat and electricity), malleable (bendable), ductile (stretchable), high density, high melting point
Chemical Properties: Most will react with oxygen

26. Nonmetals Like to gain electrons
Physical Properties: dull, don’t conduct, brittle, low density, low melting points
Can be solid, liquid or gas at room temperature depending on the element.

27. Metalloids (Semiconductors)
Share properties of both metals and nonmetals
Can be shiny or dull, conduct ok, ductile and malleable or brittle
These elements have become really important because of the computer revolution
Computer chips are made out of semiconductors (normally Si)

28. Families The periodic table can be further broken down into families
Families of elements have similar properties because they have the same number of valence electrons

29. families Metals: Groups 1-12
Alkali Metals, Alkaline-Earth Metals, Transition Metals
Groups contain both nonmetals/metalloids
Nonmetals: Oxygen, Nitrogen, Carbon, Sulfur, Phosphorus, and Selenium
Metalloids: Boron, Silicon, Germanium, Arsenic Antimony, and Terellium
The group is named by the first element in the column
Nonmetals: Groups 17-18
Halogens, Noble Gases

30. Hydrogen
Hydrogen is in group 1 but is not an alkali metal, because it is only 1 proton and 1 electron (no neutrons)
Its properties are closer to a nonmetals than to a metal
it is a colorless, odorless, explosive gas with oxygen

31. Group 1: Alkali Metals (excluding H), 1 valence e-
Very reactive, especially with water
Soft, shiny white metals (can be cut with a knife!)

32. Group 2: Alkaline-Earth Metals
2 valence e-
Not as reactive as alkali, but still very reactive.
Magnesium is used in flash bulbs

33. Groups 3-12: Transition Metals
1 or 2 valence e-
Most are silver and not that reactive so they have more everyday uses.

34. Groups 3-12: Transition Metals
Two bottom rows, or innertransition metals
Lanthanide Series: also called rare- earth metals
Actinide Series: very radioactive and not easily found in nature

35. Groups 13-16 Boron Group: Group 13, 3 valence e-
Aluminum is most common and abundant element on the planet.
Carbon Group: Group 14, 4 valence e-
Pure carbon can be diamonds, soot, or graphite, silicon and germanium are used for computer chips

36. Groups 13-16 Nitrogen Group: Group 15, 5 valence e-
Nitrogen makes up 78% of the air, Phosphorus is in soaps, and Arsenic is a well known poison
Oxygen Group: Group 16, 6 valence e- ’s
Oxygen makes up 21% of the air and is necessary for things to burn

37. Group 17: Halogens 7 valence e-
All nonmetals (can be solid, liquid or gas)
Extremely reactive with alkali metals
“Chlorine” added to pools as a disinfectant is a compound containing Chlorine, by itself chlorine is a green gas

38. Group 18: Noble Gases 8 valence e- ’s (except Helium)
Full outer shell of electrons
All are gases and extremely non- reactive (inert) and found in the atmosphere
“Neon” lights contain a variety of Noble Gases