1. Atoms and Elements
Chapter 2

2. Atomic theory
John Dalton, 1808: matter is made of tiny indestructible particles called atoms
What evidence persuaded John Dalton that matter was made of atoms?

3. Evidence for Atoms Boyle’s Law (1660s)
Gases can be compressed (PV = constant)
Suggests that a gas is made of particles with space between
Conservation of mass (Lavoisier, 1789)
In a chemical reaction, matter is neither created nor destroyed

4. Evidence for Atoms Law of Definite Proportions (Proust, 1797)
All samples of a compound, regardless of source or how prepared, have the same proportions of their constituent elements (constant composition)
Law of Multiple Proportions (Dalton, 1803)
When two elements (A and B) form two different compounds, the masses of B that combine with 1 g of A can be expressed as a ratio of small whole numbers.

5. Law of Multiple Proportions

6. Evidence for Atoms Combining Volumes of Gases
When two gases combine to form a new compound, the volumes that combine will be in a ratio of small whole numbers.

7. A New System of Chemical Philosophy: John Dalton, 1808
Each element is composed of tiny indestructible particles called atoms.
All atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements.
Atoms combine in simple, whole-number ratios to form compounds.
Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms change the way that they are bound together with other atoms to form a new substance.

8. Cathode ray tube

9. Discovery of the electron
J.J. Thompson (1897)
Cathode rays are deflected by an electric or magnetic field
Cathode ray = beam of negatively charged particles (electrons) coming out of cathode metal atoms
Determined electron mass-to-charge ratio = – x 10–9 g/C

10. Millikan oil-drop experiment

11. Charge of the electron
Robert Millikan (1909)
Drop charges are integer multiples of x 10–19 Coulomb

12. Radioactivity
Some elements spontaneously emit radiation called radioactivity
Three types of radioactivity:
Alpha (a) = heavy, +2 charge
Beta (b) = light, –1 charge
Gamma (g= high-energy photon

14. The gold foil experiment
Ernest Rutherford (1910)
Bombarded gold foil with alpha particles
Most went straight through
Some slightly deflected
A few strongly deflected
“about as credible as if you had fired a 15-inch shell at a piece of tissue paper, and it came back and hit you”

15. Alpha particle scattering explained

16. Rutherford’s nuclear model
Mass & positive charge concentrated in nucleus
Most  particles miss the nucleus & are not deflected: most of the atom is empty space
Some  particles come near the nucleus & are deflected: nucleus is positively charged and very dense
Tiny, lightweight electrons circle the nucleus, like planets around the sun

17. Composition of the atom
Protons
+1 charge
x 10–24 g = amu
Neutrons
0 charge
x 10–24 g = amu
Electrons
–1 charge
x 10–24 g = amu

18. Atomic number and mass number
Atomic number (Z) = protons
Mass number (A) = protons + neutrons
A 35Cl atom has 17 protons and 18 neutrons
Also shown as chlorine-35

19. Ion charge
When an atom loses or gains an electron, it becomes an ion
Ion charge is shown in the upper right corner of the atomic symbol
If no charge is shown, the charge is zero
mass number, A
protons + neutrons 23 11
Na1+
ion charge
protons – electrons
atomic number, Z
protons

20. Give the structure (p+, n0, e–) of these atoms
40K1+
235U
34S2–

21. Give the symbol of each atom or ion
35 p+, 44 n0, 36 e–
79Br1–
47 p+, 62 n0, 46 e–
109Ag1+
26 p+, 28 n0, 26 e–
54Fe

22. Isotopes
Atoms with the same number of protons but different numbers of neutrons are isotopes of the same element
6Li and 7Li are isotopes of lithium
Both are the element lithium
6Li has 3 protons, 3 neutrons
7Li has 3 protons, 4 neutrons

23. A Mass Spectrometer

24. Mass Spectrum of Neon
20Ne 90.48%
21Ne %
22Ne %

25. Average atomic mass
Element’s atomic mass = weighted average of masses of all naturally-occurring isotopes of that element

26. Average Atomic Mass Calculate the average atomic mass of Ne:
ISOTOPE ISOTOPIC MASS (amu) ABUNDANCE
20Ne amu %
21Ne amu %
22Ne amu %

27. Average Atomic Mass Calculate the average atomic mass of Ne:
ISOTOPE ISOTOPIC MASS (amu) ABUNDANCE
20Ne amu %
21Ne amu %
22Ne amu %
(0.9048)( amu) = amu
(0.0027)( amu) = amu
(0.0926)( amu) = amu
20.18 amu

28. How to calculate the average atomic mass
Magnesium occurs as three isotopes: 24Mg, mass amu (78.99%) 25Mg, mass amu (10.00%) 26Mg, mass amu (11.01%)
The average atomic mass of Mg is ( x ) + ( x ) + ( x ) = amu
Be careful about sigfigs!

29. Atomic mass
Two natural isotopes of antimony exist % exists as 121Sb (mass amu), and the rest is 123Sb (mass amu). What is the atomic mass of antimony?
The abundances must total 100%, so abundance of 123Sb = 100 – 57.3 = 42.7%
(0.573 x ) + (0.427 x ) = = amu

30. Percent abundance
Two natural isotopes of copper exist, 63Cu ( amu) and 65Cu ( amu). What is the abundance of each isotope?
The abundances must total 100%
If x = 63Cu abundance, then 65Cu = 1–x
Average mass (from periodic table) = amu

31. Percent abundance 63Cu = 69.15% and 65Cu = 30.85%
x = 63Cu ( amu) & 1–x = 65Cu ( amu)
Average mass = amu
63Cu = 69.15% and 65Cu = 30.85%

32. A unit for counting atoms
We have units for describing mass, volume, temperature, and so forth, but
we need a unit for counting the number of items
For large items like eggs or donuts, we can use dozen: one dozen = 12 items
For tiny items like atoms or molecules, we need the mole!

33. What is a mole?
One mole is defined as the number of atoms in exactly 12 grams of 12C
One mole contains x 1023 particles
This number is called Avogadro’s number (NA)
There is Avogadro’s number of particles in a mole of any substance

34. Calculations with Avogadro
How many atoms of gold are present in mol Au?
How many moles of Pb atoms are 8.27 x 1022 atoms of Pb?

35. Mass on the periodic table
One atom of Cl
weighs
35.45 amu OR
One mole of Cl atoms weighs grams
The molar mass of Cl is g/mol

36. Calculations with molar mass
What is the mass of 1.38 mol Al?
How many moles are in 35 g of Zn?

37. Chemistry toolbox You now have 2 mole tools in your toolbox:
Avogadro’s number
6.02 x 1023 particles = 1mol
Relationship between particles (atoms/molecules) & moles
Molar mass
The atomic mass on the periodic table, in grams = 1 mol
Relationship between grams & moles

38. Avogadro and molar mass
What is the mass of 2.35 x 1024 atoms of Cu?

39. Avogadro and molar mass
How many He atoms are present in a 22.6 g sample of He gas?

40. Mole concept Examples 2-7A, 2-7B, 2-8A, 2-8B
How many atoms of gold are present in 5.07 x 10–3 mol Au?
How many 206Pb atoms are present in 8.27 x 10–3 mol of Pb? (Pb contains 24.1% 206Pb )
What is the mass of 2.35 x 1024 atoms of Cu?
How many He atoms are present in a 22.6 g sample of He gas?

41. Mole concept Examples 2-9A and 2-9 B
How many Pb atoms are in a piece of lead with volume cm3? Density of Pb = g/cm3
0.100 mg of Re contains 2.02 x 1017 atoms of 187Re. What is the percent abundance of 187Re in this sample?

42. Exercise 30
Show that mass is conserved within experimental error in this experiment:
10.00 g calcium carbonate was dissolved in mL hydrochloric acid (d = g/mL). The products were g solution and 2.22 L carbon dioxide gas (d = g/L)

43. Exercise 31
In one experiment, 2.18 g sodium were reacted with g chlorine. All the sodium was used up, and 5.54 g sodium chloride were produced. In a second experiment, 2.10 g chlorine were reacted with g sodium. All the chlorine was consumed, and 3.46 g sodium chloride were produced. Are these results consistent with the law of constant composition?

44. Exercise 38 In an experiment like the oil drop experiment,
drop 1 carried a charge of 6.41 x 10–19 C
drop 2 had 1/2 the charge of drop 1
drop 3 had twice the charge of drop 1
drop 4 had a charge of 1.44 x 10–18 C
drop 5 had 1/3 the charge of drop 4
Do these data support the value of x 10–19 C for the electron charge? Could Millikan have inferred the electron charge from these data?

45. Exercise 42 For the ion 228Ra2+, with a mass of 228.030 u, determine
The numbers of protons, neutrons, and electrons
The ratio of the mass of this atom to that of an atom of 16O ( u)

46. Exercise 49 There are three naturally occurring isotopes of magnesium:
Calculate the weighted average atomic mass of magnesium.

47. Exercise 54
Boron has a weighted average atomic mass of u. What are the percent natural abundances of its two isotopes, 10B ( u) and 11B ( u)?

48. Exercise 58 Determine The number of Ar atoms in 5.25 mg argon
The molar mass of an element if the mass of 2.80 x 1022 atoms of that element is 4.24 g
The mass of a sample of aluminum that contains the same number of atoms as g zinc

49. Exercise 60
What is the total number of atoms present in a 75.0 cm3 sample of plumber’s solder, an alloy containing 67% Pb and 33% Sn by mass and having a density of 9.4 g/cm3?

50. Early Observations of Mass 1770s
Lavoisier - Conservation of Matter
Careful measurements showed that mass is neither created nor destroyed in a chemical reaction.
The total mass before and after reaction is the same

51. Conservation of mass Examples 2-1A and 2-1B
A g sample of magnesium reacts with g nitrogen gas. The sole product is magnesium nitride. After the reaction the mass of unreacted nitrogen is g. What mass of magnesium nitride was produced?
A 7.12g sample of magnesium is heated with 1.80 g of bromine. All the bromine is consumed and 2.07 g of magnesium bromide is the only product. What mass of magnesium is left over?

52. Early Observations of Matter 1799
Proust - Constant Composition
All samples of a given compound contain exactly the same proportion of elements by mass.
This is the Law of Definite Proportions

53. Constant composition Examples 2-2A and 2-2B
This information was provided in the example:
A g sample of magnesium combines with oxygen to yield g magnesium oxide
What is the mass of magnesium in g of magnesium oxide?
You wish to make exactly 2.00 g of magnesium oxide. What masses of magnesium and oxygen must you combine?

54. Early Observations of Matter 1803
John Dalton - Law of Multiple Proportions
When the same elements form different compounds, the combining ratios in the compounds are always in a whole number relationship.
Compound II has exactly twice as much O per C as compound I, by mass
mass of Oxygen
reacting w/ 1 g of Carbon
cmpd I g O/g C
cmpd II g O/g C

55. Definite and Multiple Proportions
3:2 3:2:3

56. The Electron: Part I J.J. Thompson - 1897
Cathode rays can be deflected by electric and magnetic fields
The charge-to-mass ratio of an electron can be determined: 1.76 x 108 C/g

57. Cathode Rays Deflected by
Thomson (1897)
Cathode Rays Deflected by
Electric Field

58. Cathode Rays Deflected by
Thomson (1897)
Cathode Rays Deflected by
Magnetic Field

59. Balancing Deflections:
Thomson (1897)
Balancing Deflections:
Measures the e/m Ratio

60. Millikan’s Oil-drop Experiment
Determines the charge of one electron

61. Radioactivity
Some elements spontaneously emit radiation called radioactivity
Three types of radioactivity:
Alpha (a) = He2+
Beta (b) = e1–
Gamma (g) = no mass or charge
Marie Curie
1876–1934

62. Alpha Particle Scatter Explained
Rutherford’s Nuclear Model

63. Atomic composition Examples 2-3A and 2-3B
Write an appropriate symbol for the species with 47 protons, 61 neutrons, and 47 electrons
Determine the numbers of protons, neutrons, and electrons in an ion of sulfur-35 that carries the charge 2–

64. Atomic Masses
Calculate the atomic mass of copper, Cu, from the following data:
ISOTOPE ISOTOPIC MASS (amu) ABUNDANCE
63Cu amu %
65Cu amu %

65. Atomic Masses
Calculate the atomic mass of copper, Cu, from the following data:
ISOTOPE ISOTOPIC MASS (amu) ABUNDANCE
63Cu amu %
65Cu amu %
63Cu: x =
65Cu: x =
63.5
the average atomic mass of Cu is 63.5 amu

66. Periodic Table
Dmitri Mendeleev developed the modern periodic table in 1869,based on his belief that element properties are periodic functions of their atomic masses
We now know that element properties are periodic functions of their ATOMIC NUMBERS