1. Chapter Three: Periodic Table

2. Understanding the Periodic Table
Atomic Mass
Selected Oxidation Numbers
Name
of Element
Atomic Number
Electron Configuration

3. Understanding the Periodic Table
For Example:
Atomic Mass
Selected Oxidation Numbers
12.011 C -4 +2 +4 6
Atomic Number
2-4
Electron Configuration
SUMMARY: All the elements or atoms have the same pattern.

4. Parts of the Periodic Table: Periods
Periods are horizontal rows on the periodic table.
There are seven periods on the periodic table.
Period one has one principle energy level and as the period number increases so does the number of principle energy levels.

5. Parts of the Periodic Table: Groups
Groups are vertical columns on the periodic table.They are numbered from 1-18.
Elements in the same group have similar properties because they have the same amount of valence electrons.
As you go down a group radius increase, electronegativity decreases and ionization energy generally decreases (because valence electrons are further away from positive protons and more electrons are repelling the valence electrons). [Table S]
Group 1
Group 2
Group 17 Li Na K Rb Cs Fr Be Mg Ca Sr Ba Ra F Cl Br I At
1 valence electron
2 valence
electrons
7 valence
electrons

6. Periodic Law
The Periodic Law states that the properties of the elements are a period function of their atomic number. This means that, when the elements are arranged by atomic number, those with similar properties will be at regular intervals; these elements will be in the same group.

7. Metals
All elements to the left of the staircase (zigzag line) are metals. They have low ionization energy (little energy is needed to remove an electron) and low electronegativity (little attraction for electrons). [Shown in Table S]
Other properties of metals:
Tend to lose electrons to form positive ions
Have metallic luster (shine)
Malleable (can be made into sheets)
Ductile (can be made into wires)
Good conductor of heat and electricity
At room temperature most metals are solids except Hg which is a liquid.

8. Nonmetals
All elements to the right of the staircase (zigzag line) except for Group 18. They have high ionization energy (a lot of energy is needed to remove an electron) and high electronegativity (big attraction for electrons). [Shown in Table S]
Other Properties of Nonmetals
Tend to gain electrons when reacting with metals, forming positive ions
They share electrons when combined with nonmetals
Lack metallic luster and are brittle in the solid phase
Poor conductors of heat and electricity.

9. Metalloids
All elements touching the staircase (zigzag line) on two sides with the exception of Al and Po. These elements have properties of both metals and nonmetals.

10. Noble Gases
All elements in Group 18. They have 8 valence electrons (with He as an exception, which has 2) and because of this, are stable. Noble gases can bond with only themselves (He-He) and only Kr, Xe, and Rn are able to bond with F and O. The forces of attraction which bonds these elements are called dispersion forces and are very weak.
As you go down the group dispersion forces increases; therefore, the boiling point increases.

11. Special Types of Elements
Alkali Metals - located in Group 1 and are the most reactive metals.
Alkaline Earth Metals - located in Group 2 and are the second most reactive metals.
Transition Metals - located from groups 3-11 and have a color in solution.
Halogens - located in Group 17 and are the most reactive non-metals.

12. Periodic Table: Most Active Metals and Nonmetals
The most active metal is located in the bottom left hand corner. It is Fr because it is easier to lose one valence electron than two, if it was in Group 2, or three, if it was in Group 3.
The most active nonmetal is located in the top right hand corner. It is F because it is easier to gain one valence electron than it is to gain more than one.

13. Ions of Metals and Nonmetals
When a metal (Na) loses and electron loses an electron, a nonmetal (Cl) gains an electrons.
The Na loses an electron and becomes an Na+ and then causes the Na to have a complete outer shell.
The Cl gains an electron and becomes Cl- causing it to have a full outer shell.
An ion and an atom only differ in the number of electrons.

14. Allotropes
Allotropes are forms of the same element that have molecular formulas.
For example O2 and O3 or carbon which differ in crystalline structure (arrangement of atoms) such as a diamond, graphite, or coal.
Allotropes have different properties.

15. Physical Properties of Elements
A physical property is a characteristic of a substance that you can notice without changing the substance into anything else.
Examples:
Color
Odor
Solubility
Density
Boiling Point
Melting Point
Hardness
Conductivity
Phase

16. Chemical Properties Of Elements
Chemical properties of an element explain how an element reacts in a chemical reaction.
Examples:
Burning
Mixing with acid or water
Distillation

17. Lewis Dot Structure
Use to show how many valence electrons an element has.
Group 1 2 13 14 15 16 17 18
Element X
X means elements
Means 1 valence electron

18. Sample Questions
The chemical properties of the elements are periodic functions of their atomic (1) masses (2) weights (3) numbers (4) radii
Boron and arsenic are similar in that they both (1) have the same ionization energy (2) have the same covalent radius (3) are in the same family of elements (4) are metalloids
(3) - You learned that properties of the elements are periodic functions of the periodic table.
(4) - Boron and arsenic are touching the zigzag line which means they are metalloids.

19. Sample Questions
3. Which of the following substances is the best conductor of electricity? (1) NaCl(s) (2) Cu(s) (3) H2O(l) (4) Br2(l)
4. Which statement best describes Group 2 elements? (1) they have one valence electron and they form ions with a 1+ charge (2) they have one valence electron and they form ions with a 1- charge (3) they have two valence electrons and they form ions with a 2+ charge (4) they have two valence electrons and they form ions with a 2- charge
3) (2) Metals are on the left of the zigzag line and are good conductors of electricity and only Cu is to the left.
4) (3) Group 2 has 2 electrons in the last shell, or 2 valence electrons. Also, look at the oxidation number. It is 2+.

20. Sample Questions
5. Which electron configuration represent the first two elements in Group 17 of the Periodic Table? (1) 2-1 and 2-2 (2) 2-2 and 2-3 (3) 2-7 and (4) 2-8 and 2-8-7
6. Which element is a liquid at room temperature? (1) K (2) I2 (3)Hg (4) Mg
5) (3) Look in the Periodic Table and look at the electron configuration of the first two elements in Group 17.
6) (3) It is an exception to the rule that metals are solids at room temperature.

21. Ionization Energy
It is the amount of energy needed to remove an electron. The smaller the amount of ionization energy, the easier it is to lose an electron.
Nonmetals have high ionization energy because they are trying to gain electrons to become more like the noble gases.
Metals have low ionization energy because they are trying to lose electrons to become more like the noble gases.

22. Ionization Down A Group
Look at Group 1. The elements Li, Na, K, and Rb are listed in order of atomic number in table S. Look at how the ionization energy decreases as you go down the group.
Ionization decreases as you go down a group because electrons are negative and protons are positive and the further away the electrons are from the nucleus, the less attraction it has to the protons so, therefore, it takes less energy to remove an electron.

23. Ionization Across A Period
Look at Period 2. The elements Li, Be, B, C, N, O, F, and Ne are listed in order of atomic number in Table S. Look at how the ionization energy increases as you go across the period.
Ionization increases because there are more protons attracting the electrons, making it harder to remove an electron.

24. Electronegativity
Electronegativity is the attraction (or “desire”) for electrons. The larger the electronegativity, the more the atom attracts electrons.
Electronegativity is a scale from 0-4, 0 being the least desire for electrons and 4 being the most.
Back to the most attractive elements, F has the highest electronegativity because it is the closet nonmetal to being a noble gas by gaining one electron. Fr has the lowest electronegativity because it is the closet metal to being a noble gas if it loses an electron.

25. Electronegativity Down A Group
Look at Group 1. The elements Li, Na, K, and Rb are listed in order of atomic number in Table S. Look at how the electronegativity decreases as you go down the group.
Electronegativity decreases because the elements try to lose electrons to become more like the noble gases, or more stable. This is because noble gases have a full (8 electrons) outer shell.

26. Electronegativity Across A Period
Look at Period 2. The elements Li, Be, B, C, N, O, F, and Ne are listed in order of their atomic number in Table S. Look at how the electronegativity increases, with the exception of Ne which has no desire for electrons.
This is because the elements want to be more like a Group 18 element.

27. Atomic Radius
An atomic radius is the distance from the nucleus to the outermost electron.
The atomic radius for all elements on the periodic table can be seen in Table S.

28. Atomic Radius Down A Group
Look at the elements in Group 1. The elements Li, Na, K, and Rb are listed in order by atomic number in Table S. Look at how the atomic radius increases.
The atomic radius increases because there are more levels of electrons, making the atom bigger.

29. Atomic Radius Across A Period
Look at the elements Li, Be, B, C, N, O, F, and Ne in Period 2. Look at how the atomic radius generally decreases.
This is because there are more protons which makes there be more attraction, pulling the electrons closer to the nucleus.

30. Ionic Radius
For a metal, it loses electrons and becomes smaller. They become positive and are called cations. For example, Li loses an electron and becomes Li+. The ionic radius becomes smaller because it has less electrons.
For a nonmetal, it gains electrons and become bigger. They become negative and are called anions. For example, F gains an electron and become F-. The ionic radius becomes larger because it has more electrons.

31. REMEMBER
Atomic mass, Atomic number, Oxidation states, and Electron Configurations are all listed in the Reference Table On The Periodic Table.
Electronegativity, Atomic Radius, Ionization Energy, and Ionic Radius are all listed on Reference Table S.

32. Sample Questions
A diatomic element with a high first ionization energy would most likely be a (1) nonmetal with high electronegativity (2) nonmetal with low electronegativity (3) metal with high electronegativity (4) metal with low electronegativity
1) (1) - Compare the ioniztion energy in Group 1 and Group 17. Nonmetals have high ionization energy and electronegativity. Diatomic elemetns are in Group and they are to the right of the zigzag line.

33. Sample Questions
2. An atom of which of the following elements has the smallest atomic radius? (1) Li (2) Be (3) C (4) F
3. Which atom has a radius larger than the radius of its ion? (1) Cl (2) Ca (3) S (4) Se
2) (4) Look at Table S.
3) (2) Ca is a metal and a metal loses an electron an becomes an ion. Also it is the one that is different, all the other choices are nonmetals.