1. Ch. 3: Atoms
3.1 Foundations

2. History Democritus named the most basic particle
atom- means “indivisible
Aristotle
didn’t believe in atoms
thought matter was continuous

3. History by 1700s, all chemists agreed:
on the existence of atoms
that atoms combined to make compounds
Still did not agree on whether elements combined in the same ratio when making a compound

4. Law of Conservation of Mass
mass is neither created or destroyed during regular chemical or physical changes

5. Law of Definite Proportions
any amount of a compound contains the same element in the same proportions by mass
No matter where the copper carbonate is used, it still has the same composition

6. Law of Multiple Proportions
applies when 2 or more elements combine to make more than one type of compound
the mass ratios of the second element simplify to small whole numbers

7. Law of Multiple Proportions

8. Dalton’s Atomic Theory
All mass is made of atoms
Atoms of same element have the same size, mass, and properties
Atoms can’t be subdivided, created or destroyed
Atoms of different element combine in whole number ratios to make compounds
In chemical reactions, atoms are combined, separated, and rearranged.

9. Modern Atomic Theory Some parts of Dalton’s theory were wrong:
atoms are divisible into smaller particles (subatomic particles)
atoms of the same element can have different masses (isotopes)
Most important parts of atomic theory:
all matter is made of atoms
atoms of different elements have different properties

10. Ch. 3: Atoms
3.2 Structure of Atom

11. Structure of Atom Nucleus: Electron Cloud:
contains protons and neutrons
takes up very little space
Electron Cloud:
contains electrons
takes up most of space

12. Subatomic Particles includes all particles inside atom
proton
electron
neutron
charge on protons and electrons are equal but opposite
to make an atom neutral, need equal numbers of protons and electrons

13. Subatomic Particles
number of protons identifies the atom as a certain element
protons and neutrons are about same size
electrons are much smaller
nuclear force- when particles in the nucleus get very close, they have a strong attraction
proton + proton
proton + neutron
neutron + neutron

14. Atomic Radius size of atom
measured from center of nucleus to outside of electron cloud
expressed in picometers (1012 pm = 1 m)
usually pm

15. Example
An atom has a radius of 140 pm. How large is that in meters?

16. 3.2b: Important Discoveries
Ch. 3 Atoms
3.2b: Important Discoveries

17. Discovery of Electron
resulted from scientists passing electric current through gases to test conductivity
used cathode-ray tubes
noticed that when current was passed through a glow (or “ray”) was produced

18. Discovery of Electron Noted Qualities of Ray Produced:
existed- there was a shadow on the glass when an object was placed inside
had mass- the paddle wheel placed inside, moved from one end to the other so something must have been “pushing” it

19. Discovery of Electron Noted Qualities of Ray Produced:
negatively charged- the rays behaved the same way around a magnetic field as a conducting wire
negatively charged- were repelled by a negatively charged object

20. Discovery of Electron
all of these led scientists to believe there were negatively charged particles inside the cathode ray
Milliken found the mass of the electron

21. Discovery of Electron
J.J. Thomson (English 1897) did more experiments to actually make the discovery
he found ratio of charge of this particle to this mass of the particle
since the ratio stayed constant for any metal that contained it, it must be the same in all of the metals

22. Are electrons the only particles?
since atoms are neutral, something must balance the negative charge
since an atom’s mass is so much larger than the mass of its electrons, there must be other matter inside an atom

23. Discovery of Nucleus
Rutherford discovered the nucleus by shooting alpha particles (have positive charge) at a very thin piece of gold foil
he predicted that the particles would go right through the foil at some small angle

24. Discovery of Electron

25. Discovery of Nucleus
some particles (1/8000) bounced back from the foil
this meant there must be a “powerful force” in the foil to hit particle back
Predicted Results Actual Results

26. Discovery of Nucleus Characteristics of “Powerful Force”:
dense- since it was strong enough to deflect particle
small- only 1/8000 hit the force dead on and bounced back
positively charged- since there was a repulsion between force and alpha particles

27. Find the element Sodium on your periodic table
Find the element Sodium on your periodic table. What do you know about atoms of sodium from the information on the table?

28. Ch. 3 Atoms
3.3 Counting Atoms

29. Atomic Number (Z) number of protons
All atoms of the same element have the same atomic number
located above the symbol in the periodic table
order of the elements in the periodic table

30. Isotopes atoms of the same element with different numbers of neutrons
most elements exist as a mixture of isotopes
What do the Carbon isotopes below have in common? What is different about them?

31. Mass Number sum of particles in nucleus A = #p + #n
Hydrogen isotopes have special names:
protium
deuterium
tritium
What do the
prefixes
in their names
come from?

32. Designating Isotopes Hyphen notation: Nuclear Symbol notation:
Name - mass number
ex. Carbon – 13
Nuclear Symbol notation:

33. Examples 7 protons, 8 neutrons Nitrogen-15 17 electrons, 19 neutrons
Chlorine- 36

34. Examples
Z=5, 6 neutrons
Boron- 11
A=75, 42 neutrons
Arsenic- 75

35. Examples Carbon- 13 Xenon-131 Sodium-24 Oxygen- 15 #p #n #e Z A

36. Examples Carbon- 13 Xenon-131 Sodium-24 Oxygen- 15 #p #n #e Z A

37. A neutral atom contains 34 electrons and has an A of 59
A neutral atom contains 34 electrons and has an A of 59. Write the nuclear symbol notation and hyphen notation for this isotope.

38. Ch. 3 Atoms
3.3 Counting Atoms

39. Relative Atomic Mass
since masses of atoms are so small, it is more convenient to use relative atomic masses instead of real masses
to set up a scale, we have to pick one atom to be the standard
since 1961, the carbon-12 nuclide is the standard and is assigned a mass of exactly 12 amu

40. Relative Atomic Mass
atomic mass unit (amu)- one is exactly 1/12th of the mass of a carbon-12 atom
mass of proton= amu
mass of neutron= amu
mass of electron= amu

41. Relative Atomic Mass
the mass number (A) and the relative atomic mass are very close but not the same because
relative atomic mass includes electrons
the proton and neutron masses aren’t exactly 1 amu

42. Average Atomic Mass
weighted relative atomic masses of the isotopes of each element
each isotope has a known natural occurrence (percentage of that elements’ atoms)

43. Calculating Average Atomic Mass
Naturally occurring copper consists of:
69.71% copper-63 ( amu)
30.83% copper-65 ( amu)
( x )+( x )
=63.55 amu

44. Calculating Average Atomic Mass
An element has three main isotopes with the following percent occurances:
#1: amu, 90.51%
#2: amu, 0.27%
#3: amu, 9.22%
Find the average atomic mass and determine the element.

45. Calculating Average Atomic Mass

46. The mole
a unit for measuring a very large amount- like number of atoms or molecules in a sample
like one dozen (1 dozen = 12 things)
except bigger: mole = 6.022x1023 things
Why 6.022x1023 ?
6.022x1023 is the number of atoms in exactly 12 g of carbon-12

48. The mole
6.022x1023 is called Avogadro’s Number in honor of all of his contributions to chemistry
can be used as a conversion factor between a number of things and mole

49. Molar Mass the mass of one mole of pure substance in grams per mole
numerically equal to average atomic mass
under the symbol on the periodic table
can be used as a conversion factor between moles and grams

50. Conversion Factors # Atoms Grams Moles Use Molar Mass:
grams per mole
Use Avog.’s Number:
atoms per mole

51. Gram  Moles use molar mass Ex. 32.3 g Na = ? mol Na
Ex mol Fe = ? g Fe

52. What is the mass in grams of 3.5 mol of Cu?
Grams of Fe in 2.25 mol of Fe?
Grams of K in mol of K?
How many moles are in 11.9g Al?
Number of moles in 5g of Ca?
Number of moles in 3.60 x g Au?

53. # Atoms  Moles use Avogadro’s Number Ex: 1.40 mol Na = ? Na atoms
Ex: 3.4x1023 atoms Fe = ? mol Fe

54. # of moles of Ag in 3.01 x 1023 atoms
Moles of Pb in 1.50 x atoms Pb?
Moles of Sn in 2500 atoms of Sn?

55. 3.01x1023 Ag atoms ( 1 mol Ag) =
6.022x1023
0.500 mol Ag
1.50x1012 atoms Pb ( 1 mol Pb) =
2.49x10-12 mol Pb
2500 Sn atoms ( 1 mol Sn) = 4.2x10-21

56. Grams  # Atoms Ex: 0.0326 g N = ? atoms of N
use both: Avogadro’s # and molar mass
Ex: g N = ? atoms of N
Ex: 2.01x1041 atoms of H = ? g H

57. 1. Mass of 1.20x108 atoms of Cu?
2. Mass of 7.7x1015 atoms of Ni?
3. # of S atoms in 4g of S?