1. Atomic Structure
HW: read CH3 (and 11.4c)

2. Early models of the atom
History of the atom (5 min)
Atoms are defined as the smallest particle of an element that retains the chemical identity of that element
Democritus, Dalton, Faraday, Franklin, Thompson, Rutherford, Moseley, Bohr all had significant contributions
(*you don’t need to write down slides with *)

3. *Early models of the atom
Democritus (450 BC) first individual to say that matter was made of “stuff” called atoms.
Now atoms are defined as the smallest particle of an element that retains the chemical identity of that element

4. *Early models of the atom
Dalton (1803) came up with 4 postulates that explained the properties of matter that is called the atomic theory of matter. Dalton thought atoms were hard and round.
Each element is composed of extremely small particles called atoms
All atoms of a given element are identical, but they differ from those of another element
Atoms are neither created nor destroyed in any chemical reaction
A given compound always has the same relative numbers and kinds of atoms

5. *Early models of the atom
Franklin (1750ish) concluded an object could have one of two kinds of electric charge (positive/negative) and the charges repel each other
Faraday (1850ish) thought the atom structure was related to electricity (electrical charges)

6. *Early models of the atom
Thomson (1910ish) –worked on cathode ray tubes and he concluded that these tubes contain negative particles that come from the cathode end (electrons).
Decided the atom had all its particles evenly distributed throughout it. Known as the “plum pudding” model of an atom.
3 min

7. *Early models of the atom
Rutherford (1900ish) Atoms had neutral charge and electrons (negative charge) were in atoms so he set out to prove there were positively charged particles.
He shot a beam of alpha particles (positive charge) at a thin piece of gold. Most of the particles pass straight through the foil but there were a few that scattered in every direction.
These particles bounced off a small positively charged core (nucleus) at the center of the atom.
Recreation 4 min

8. *Early models of the atom
Bohr – planetary model where electrons orbit around the nucleus
Electron location is in energy levels
1st energy level can hold 2 electrons, 2nd can hold 8 electrons, 3rd can hold 18 electrons…

9. Atomic number
*Moseley: Explained Dalton’s postulate of “All atoms of a given element are identical, but they differ from those of another element” when he discovered that each element had a unique number of protons. The number of protons became the atomic number and the way to differentiate and categorize elements
Example: Oxygen’s atomic # is 8, so it has 8 protons
There are also 8 electrons (neutral atom)

10. What we know(?) about the atom
Atoms are composed of protons (positive), neutrons (neutral), and electrons (negative). In a neutral atom, protons and electrons are equal.
The protons and neutrons are found in the nucleus. Thus, the nucleus has a positive charge.
The mass of atomic particles is so small that instead of measuring in grams, the mass is measured in atomic mass units (amu).
Just how small is the atom(5 min)

11. The periodic table of elements
118 known elements 
Arranged by atomic number
Also give the atomic mass (average mass)
F = chemical symbol (Fluorine)
9 = atomic number (number protons)
= atomic mass (number protons and neutrons)
3 min song

14. protons Protons = p+ Lives in nucleus Positive charge
Mass of 1 AMU (same as neutrons)
# protons is unique to each element

15. neutrons Neutrons = n0 Lives in nucleus No charge (neutral)
Mass of 1 AMU (same as proton)
Helps stabilize nucleus

16. electrons Electron = e- Lives in electron shell or cloud
Negative charge
Much smaller than p+ or n0 (1/1836 AMU)
Can turn atom into a charged ion

17. Determining numbers Number of protons = atomic number
Number of electrons = number of protons (atomic number)
Number of neutrons = atomic mass (rounded to nearest whole number) minus atomic number

18. YOU TRY! Determine the number of protons, neutrons, and electrons in
A) 88 Ra 226
B) 12 Mg 24
C) 79 Au 197
D) 63 Eu 152

19. isotopes
Atoms of the same element but different numbers of neutrons are called isotopes.
Have the same atomic number but different atomic mass.
The nature of each atom primarily depends on the number of electrons so atoms with different neutrons are barely indistinguishable from each other. The only difference is that the one with more neutrons is heavier.
Example: 
75.5% abundance % abundance

20. YOU TRY! Determine the number of protons, neutrons, and electrons in
A) 6 C 12
B) 6 C 13
C) 51 Sb 121
D) 51 Sb 123

21. Isotopes and atomic mass
Naturally occurring chlorine that is put in pools is 75.53% 35Cl and 24.47% 37Cl. Calculate the average atomic mass (note: this is why mass in the periodic table is not whole number)
1) Convert % to a decimal
2) multiply mass by decimal
3) Add all products
4) Check to see if it is reasonable!
Copper used in electric wires comes in two types: 63Cu and 65Cu. 63Cu has an abundance of 69.09%. 65Cu has an abundance of 30.91%. Calculate the average atomic mass.

22. Development of the periodic table
Dimitri Mendeleev- constructed first periodic table by arranging elements according to similarities in properties
This table was able to predict the chemical properties of unknown elements and the existence of unknown elements
Henry Moseley- arranged elements by order of atomic number, currently arranged this way
History of table 5 min

23. Modern periodic table Arrangement
A. Horizontal rows are called periods; vertical columns are called groups
B. Each group is identified by a numeral and the letter A or B
Periodic Law: when elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties.
A. The properties of an element are related to the element’s electron configuration.

25. Characteristics of the Periodic Table
All elements can be classified as
A. metals
B. nonmetals
C. metalloids/semimetals
Metals (left side of the table)
A. Characteristics: high luster (shine) and high electrical conductivity
Nonmetals (right side of table)
A. Characteristics: non-lusterous and poor electrical conductors
3 min song

26. Metals
Conductors
Lose electrons
Malleable and ductile

27. Nonmetals
Brittle
Gain electrons
Covalent bonds

28. Semi-metals or Metalloids

29. *group characteristics
Group 1: Alkali Metals
A. low density and melting point, soft (can be cut with knife)
B. very reactive with water
C. uses: form hydroxides (NaOH) that assist paper making, petroleum refining
Group 2A: Alkaline Earth Metals
A. when reacted with water, produce alkaline solutions
B. extracted from mineral ore (“earth” metals)
C. uses: Sr (red color in fireworks), magnesium used in jet engines, calcium essential to body

30. Alkali Metals

31. Alkaline Earth Metals

32. *group characteristics
Group 3A: Aluminum Group
A. includes both metals and nonmetals
B. uses: aluminum, structural materials, gallium in thermometers
Group 4A: Carbon Group
A: uses: carbon (graphite and diamond, living things are based on carbon), silicon second most abundant element, tin has metal uses
Group 5A: Nitrogen Group
A. uses: N and P are essential to living organisms (DNA, bones, teeth), less common but used in metals- arsenic, antimony, and bismuth

33. *group characteristics
Group 6A: Oxygen/Chalcogen Group
A. uses: oxygen is most abundant element (photosynthesis, medicine, manufacture of steel), S is essential to living organisms and to make rubber, selenium is a semiconductor, tellurium is rare
Group 7A: Halogens and Hydrogen
A. fairly abundant, exists as salts with group 1A elements
B. free halogens are very reactive, they usually exist in molecules of 2 atoms (H2, Br2, F2)
C. uses: fluorine (healthy teeth), chlorine (blood and other fluids), iodine (thyroid gland, early antiseptic)

34. Halogens

35. *group characteristics
Group 8A: Noble Gases
A. tend not to combine with other elements (inert)
B. uses: weather balloons, neon lights
Groups 1B-8B: Transition Metals
A. typical metals, metallic luster, conduct electricity, most are colored
B. uses: light bulb filament (W), wire (mostly Cu), hemoglobin (Fe)

36. Noble Gases

37. Transition metals

38. Inner Transition Metals

39. ions
(this is where Bohr model was successful)
If atom gains or loses electrons (usually through bonding) then the atom is no longer neutral.
Lose electron = positive charge = cation
Gain electron = negative charge = anion (ant)
Look at the GROUP number. The electrons available for movement is the group number. All elements follow the octet rule: they gain or lose electrons to attain an electron configuration of the nearest noble gas.

40. ions Example: Magnesium (atomic #12) loses 2 electrons
Magnesium now  Mg 2+
Example: Fluorine (atomic #9) gains 1 electron
F 1-
NOTE: Fl is an atom; Fl1- is an ion

41. White boards On your whiteboard Label groups # 1-8 at top
Label the ionic charge for each group
Once done, be ready for specific element practice

42. YOU TRY!
A) What is the ionic charge of K?
B) What is the ionic charge of O?
C) What is the ionic charge of C?

43. YOU TRY!
An ion has a net charge of -1. It has 18 electrons and 20 neutrons.
A) How many protons are present?
B) What is its atomic number?
C) What is its symbol?
D) What is its mass?
E) What is the charge on the nucleus?

44. YOU TRY!
An ion has a net charge of +2. It has 10 electrons and 13 neutrons.
A) How many protons are present?
B) What is its atomic number?
C) What is its symbol?
D) What is its mass?
E) What is the charge on the nucleus?

45. YOU TRY! How many protons, neutrons, electrons are in the following:
3170 Ga3+
55133 Cs1+
1632 S2-
53127 I1-

46. Ionic compounds
An ionic compound is formed between a metal and nonmetal.
The overall charge must be neutral
Example: If Li and N combine, what formula is for the compound to be neutral?
1) Li charge = 1+
N charge = 3-
2) Need 3 Li molecules to become neutral; number is subscript
3) Formula = Li3N

47. Ionic compounds Give the formula for the ion pairs: A) K and I
B) Mg and N

48. Ionic compounds
If given a name, it is the same concept, just one extra thinking step.
Write the formula for:
A) Cesium nitride
B) Barium chloride
C) Boron iodide

49. *Periodic trends ATOMIC RADIUS Atoms get larger going down a group
Atoms get smaller left to right across a period
more protons as you move to the right; the more protons give more atomic pull. The strong attractive force shrinks orbitals so the atom is smaller

50. *Periodic trends IONIC SIZE
Recall: Ions are elements that gain or lose electrons 
More electrons the size becomes larger
more repulsion's, spread out electrons, increase size
Fewer electrons the size becomes smaller
reduces repulsion's, electrons pulled closer to nucleus

51. *Periodic trends IONIZATION ENERGY
Energy needed to remove one of its electrons
Atoms with high I.E. hold onto their electrons tightly
Atoms with low I.E. easily lose electrons
I.E decreases as you move down a group
I.E. increases as you move from left to right 
think of octet rule for optimum number of electrons

52. *Periodic trends ELECTRONEGATIVITY
Ability to attract electrons, related to Ionization Energy
Fluorine most electronegative (wants electron the most)
Left side of table least electronegative (not want electrons)