1. Chapter 11 & 12
The Periodic Table
&
Periodic Law

2. Meyer & Mendeleev
In 1869, published almost identical versions with the elements in order of increasing atomic mass and in columns with similar properties.

3. Father of the Periodic Table
Mendeleev
Father of the Periodic Table
Mendeleev is given more credit than Meyer BECAUSE:
He published his table first
He better demonstrated his table
Suggested some of the previously measured masses were incorrect
Left blanks for not yet discovered elements

4. Mosley
In 1913, using X-rays, he discovered a unique number of protons in the nuclei of atoms for each element.
Today the elements are arranged in order of increasing atomic number
There is a periodic repetition of chemical and physical properties of the elements when they are arranged in order of increasing atomic number
Periodic Law

5. Arrangement of the Periodic Table
Groups/Families
18 vertical columns (↑↓)
Two Labeling Systems
Number-and-letter system
1A through 8A columns (representative elements)
1B through 8B short columns (transition elements)
2. Number system, Group 1 to Group18
Periods
7 horizontal rows (↔)

6. Non-Metals
Create a Legend
Metalloid
Non-Metals
Metals
Metals

7. Metals Shiny Solid at room temperature
Good conductors of heat and electricity
Malleable
Ductile Group 1
Alkali Metals
Group 2
Alkaline Earth Metals
Groups 3-12
Transition Metals
Lanthanide & Actinide Groups
Inner Transition Metals or Rare Earth Metals

8. Group # Group Name # Valence e-
Valence Electrons-electrons in an atom’s outermost orbitals
Group # Group Name # Valence e-
1(1A) Alkali Metals
2(2A) Alkaline Earth Metals
Transition Metals varies
13(3A) Boron Group 3
14(4A) Carbon Group 4
15(5A) Nitrogen Group
16(6A) Oxygen Group 6
17(7A) Halogens
18(8A) Noble Gases

9. Nonmetals & Metalloids
Dull
Generally gases or brittle solids at room temperature
Poor conductors of heat and electricity
Metalloids
Elements with physical and chemical properties of both metals and nonmetals
Rest on the “stair-step” B Si As Te At Ge Sb Po
←Metals

10. ALL ELEMENTS IN GROUP 1 (ALKALI METALS) HAVE ONE VALENCE ELECTRON
Element Placement
Why are elements put into groups/families together?
Because they have similar chemical properties
Why do elements have similar chemical properties?
Because they have the same number of valence electrons
Group 1 – Alkali Metals
Period 2 Lithium 1s22s1 [He]2s1
Period 3 Sodium 1s22s22p63s1 [Ne]3s1
Period 4 Potassium 1s22s22p63s23p64s1 [Ar]4s1
ALL ELEMENTS IN GROUP 1 (ALKALI METALS) HAVE ONE VALENCE ELECTRON

11. Periodic Trends

12. Atomic Radius
Half the distance between two nuclei of identical atoms that are chemically bonded together
Down the group
atomic radius increases,
because…
Across the period
atomic radius decreases,
becauses….

13. H He B C N N F Ne Na Mg Mg Al Si Cl K Ca V Cr Ge Cs W
Atomic Radius Decreases
Atomic Radius Increases H He B C N N F Ne Na Mg Mg Al Si Cl K Ca V Cr Ge Cs W

14. Ionization Energy
The amount of energy required to remove an electron from the atom
(how tightly an atom holds on to its electrons)
A general term for the energy required to remove an electron from an orbital in an atom. Think of it also as the energy required to make a cation.
Down a group
ionization energy decreases, because…
Across a period
ionization energy increases, because…

15. Practice Ionization Energy
Which has the greater ionization energy?
Ne or Ar N or O Sc or Ti
Which has the smaller ionization energy?
Al, Si, P K, Rb, Sr Be, Mg, Ca

16. Be N O Ne Mg Al Si P Ar K Ca Sc Ti Rb Sr Ionization Energy Increases
Ionization Energy Decreases Be N O Ne Mg Al Si P Ar K Ca Sc Ti Rb Sr

17. Electron Affinity
The electron affinity of an element is the energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion.
The attraction to additional electrons
A fluorine atom in the gas phase, for example, gives off energy when it gains an electron to form a fluoride ion.
Down the group
electron affinity decreases, because
Across the period
electron affinity increases, because

18. Ionic Radius Octet Rule Ion
Atoms tend to gain, lose, or share electrons in order to achieve a full outer energy level (typically 8 are needed)
Groups 1 A – 3 A, loses Valence e-( 1 to 3 e-).
Group 4 A, share Valence e-.
Groups 5 A – 7 A, gain electrons ( 3 to 1 e-).
Noble Gases, Outer most shell are full. These elements don’t gain nor lose e-, Non-reactive.
Ion
An atom that has an overall charge due to the gaining or losing of electrons
Cation, positive charges (Groups 1A – 4A)
Anion, negative charges (Groups 5A – 7A)

19. Ionic Radius Comparisons
Metals have LOW ionization energy and electron affinity
They lose electrons to form positively charged ions
Positive charged ions are smaller than the original atom
Nonmetals have HIGH ionization energy and electron affinity
They gain electrons to form negatively charged ions
Negatively charged ions are larger than the original atom
Atom size decreases
Atom size increases

21. Electronegativity
The attraction an atom has for electrons in a covalent bond.
The ability of an atom to attract electrons in a chemical bond.
Down the group
Electronegativity values decrease, because
Across the period
Electronegativity values increase, because
*Noble gases are the exception to this rule.

22. You must subtract the values of electronegativity to determine it the bond is covalent, polar covalent or ionic
Electronegativity
Electronegativity of 0.0 to 0.3 is a non-polar covalent bond.
Covalent shares the electrons equally
Electronegativity of 0.4 to 1.9 is a polar covalent bond.
Polar is slightly negative on one side
Electronegativity of 2.0 to 4.0 is an ionic bond.
Ionic has electrons captured by one atom.