1. Atomic Structure, Periodic Table, & Lewis Structures

2. Basic Definitions
Atom – smallest unit of an element that retains the properties of that element
Atoms - made up of several subatomic particles: protons, neutrons, and electrons

3. Protons, Neutrons, & Electrons
+1 charge
found in the nucleus of the atom
Neutrons
no charge
found in the nucleus of an atom
Electrons
-1 charge
found on the outside of the nucleus
Nucleus
made up of protons and neutrons
overall + charge

4. Atomic Structure

5. Atomic Numbers
Atomic number (Z) - number of protons in the nucleus of an atom of that element.
The number of protons determines the identity of an element
The number of protons for an element CANNOT be changed.

6. Atomic Numbers
Atoms have no overall electrical charge so the number of protons must equal the number of electrons.
So, the atomic number of an element also tells the number of electrons in a neutral atom of that element.
The number of electrons can be changed when determining the charge of an ion.

7. Masses
The mass of a neutron is almost the same as the mass of a proton.
The sum of the protons and neutrons in the nucleus is the mass number (A) of that particular atom.
Isotopes of an element have different mass numbers because they have different numbers of neutrons, but they all have the same atomic number (number of protons & electrons)

8. Isotopes
When writing isotopes, the atomic number (or number of protons) will appear at the bottom left of the formula
The mass number (number of protons plus neutrons will appear at the top left of the formula.
The element symbol will appear to the right of the numbers
The different number of neutrons has NO bearing on chemical reactivity

9. Writing the Names of Isotopes
When writing the name of an isotope, you will write the name of the element – the mass number
For example 126 C would be named:
Carbon - 12

10. Try the following Name Symbol # Protons # Neutrons # Electrons Mass #
Carbon – 11
197 Au 79 1 2 25 55
Oxygen - 15

11. Try the following Name Symbol # Protons # Neutrons # Electrons Mass #
Carbon – 11 11 C 6 5
Gold - 197
197 Au 79
118
Hydrogen – 3 3 H 1 2
Manganese - 55 55 Mn 25 30
Oxygen - 15 15 O 8 7

12. Try this one Name Symbol # Protons # Neutrons # Electrons Mass #
Iodine

13. Try this one Name Symbol # Protons # Neutrons # Electrons Mass #
Iodine
130
I -1 53 77 54

14. Atomic Mass
Atomic mass –the weighted average mass of all the naturally occurring isotopes of that element.
The number is usually located at the bottom of the periodic table and has decimal places

15. Calculating Atomic Mass

16. Calculating Atomic Mass
Copper exists as a mixture of two isotopes.
The lighter isotope (Cu-63), with 29 protons and 34 neutrons, makes up 69.17% of copper atoms.
The heavier isotope (Cu-65), with 29 protons and 36 neutrons, constitutes the remaining 30.83% of copper atoms.

17. Calculating Atomic Mass
First, calculate the contribution of each isotope to the average atomic mass, being sure to convert each percent to a fractional abundance.

18. Calculating Atomic Mass
The average atomic mass of the element is the sum of the mass contributions of each isotope.

19. Try this one…
Calculate the atomic mass of germanium.
72.59 amu

20. You can tell many things from an isotope formula
Hydrogen has three naturally occurring isotopes in nature: Hydrogen – 1, Hydrogen – 2, and Hydrogen – 3.
Which is the most abundant in nature?
Hydrogen – 1
Which is the heaviest?
Hydrogen - 3

21. Periodic Table
Periodic Table – arrangement of elements in order of increasing atomic number with elements having similar properties in vertical columns
Groups – vertical columns
Periods – horizontal rows

22. Group Names Group Name 1A Alkali Metals 2A Alkaline Earth Metals 6A
Chalcogens
Halogens
Noble Gases

23. Groups
The group tells you the number of valence electrons that the element has
Valence electrons - electrons in the outermost shell of the atom
All group 1A elements have 1 valence electron. Likewise, all group 8A elements have 8 valence electrons.

24. Characteristics
Elements in the same group exhibit similar chemical characteristics due to the fact that they all have the same number of valence electrons.
The most stable number of valence electrons is 8
This is called an octet

25. Charges
Every element wants 8 valence electrons to become stable. They will gain or lose valence electrons to form an octet
For example…Group 1A elements have 1 valence electron. They can either gain 7 electrons to have an octet or lose 1.
What is easier?
Lose 1
If an element loses 1 electron (1 negative charge) what charge will the resulting ion have? +1

26. Charges Let’s go to group 7A. This group has 7 valence electrons
It can either loose 7 or gain 1
What is the easiest?
Gain 1
What will be the resulting charge if the element gain 1 electron (1 negative charge)? -1

27. Physical States and Classes of the Elements
The majority of the elements are metals. They occupy the entire left side and center of the periodic table.
Nonmetals occupy the upper-right-hand corner.
Metalloids are located along the boundary between metals and nonmetals.

28. Metals
Metals are elements that have luster, conduct heat and electricity, and usually bend without breaking.
All metals except mercury are solids at room temperature; in fact, most have extremely high melting points.
The periodic table shows that most of the metals are not main group elements.

29. Transition Metals
The elements in Groups 3 through 12 of the periodic table are called the transition elements.
All transition elements are metals.
Many transition metals can have more than one charge

30. Non Metals
Most nonmetals don’t conduct electricity, are much poorer conductors of heat than metals, and are brittle when solid.
Many are gases at room temperature; those that are solids lack the luster of metals.
Their melting points tend to be lower than those of metals.

31. Metalloids
Metalloids have some chemical and physical properties of metals and other properties of nonmetals.
In the periodic table, the metalloids lie along the border between metals and nonmetals.

32. Lewis Structure
Lewis Structures – shows how the valence electrons are arranged among the atoms of a molecule
There are rules for Lewis Structures that are based on the formation of a stable compound
Atoms want to achieve a noble gas configuration

33. Octet & Duet Rules Octet Rule – atoms want to have 8 valence electrons
Duet Rule – H is the exception. It wants to be like He & is stable with only 2 valence electrons

34. Steps for drawing Lewis Structures
Sketch a simple structure with a central atom and all attached atoms
Add up all of the valence electrons for each individual atom
If you are drawing a Lewis structure for a negative ion add that many electrons to create the charge
If you are drawing a Lewis structure for a positive ion subtract that many electrons to create the charge

35. Steps for drawing Lewis Structures
Subtract 2 electrons for each bond drawn
Complete the octet on the central atom & subtract those electrons
Complete the octet on the surrounding atoms & subtract those electrons
Get your final number
If 0  you are done!
If +  add that many electrons to the central atom
If -  need to form multiple bonds to take away that many electrons

36. Examples
CCl4
Sketch a simple structure with a central atom and all attached atoms Cl │
Cl – C – Cl

37. Examples Add up all of the valence electrons for each individual atom
4 + 4(7) = 32
Subtract 2 electrons for each bond drawn
32-8 = 24
Complete the octet on the central atom & subtract those electrons
Done

38. Examples
Complete the octet on the surrounding atoms & subtract those electrons
24 – 24 = 0
Final number = 0…DONE!
Final structure is… __
│Cl │
__ │ __
│Cl – C – Cl │ │

39. Examples HF

40. Examples
NH3

41. Examples
NO+

42. Exceptions to the octet rule
Sometimes the central atom violates the octet rule and has more or less than 8 valence electrons
Keep using the same rules to draw Lewis Structures

43. Examples
SF4
ICl3
XeF4
ICl-

44. Resonance
When more than one Lewis Structure can be written for a particular molecule
Resonance structure – all possible Lewis structures that could be formed
The actual structure is the average of all of the structures
You MUST show all structures!

45. Examples
SO3
NO2-
NO3-