1. Chapter 2: Atoms and Elements

2. Modern atomic theory development
Nature of atom theorized as early as the 5th century B.C. (Leucippus and Democritus)
It wasn’t until the 18th century that laws were discovered leading to the acceptance of modern atomic theory
Law of conservation of mass
Law of definite proportions
Law of mutiple proportions

3. Law of conservation of mass
Antoine Lavosier – in 1789, he performed various chemical reactions to formulate his law
Law - In a chemical reaction, matter is neither created or destroyed
Atom bonds are rearranged but the atoms themselves are not created or destroyed in the process
Mass you start with, you end up with

5. Law of definite proportions
Joseph Proust – in 1797, he found that elements composing a given compound always occur in a fixed proportion/ratio in all samples of that compound
Law – All samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements
18g of H2O breaks down into 16g of O2 and 2g of H2, O2:H2 ratio = 8:1
36g of H2O breaks down into 32g of O2 and 4g of H2, O2:H2 ratio = 8:1
Ratios are the same regardless of sample size for a pure compound

6. Practice
When CO2 decomposes into oxygen and carbon, it gives a gram ratio of 2.67:1 O2:C.
If a sample of CO2 decomposes to produce 5.84 grams of carbon, how much O2 is produced? How much was the original amount of CO2 in grams?
2. When a 32.4g sample of CO2 decomposes, how many grams of carbon are produced, and how many grams of O2 are produced?

7. Law of multiple proportions
John Dalton – in 1804 published the law of multiple proportions
Law – When two elements (call them A and B) form two different compounds, the masses of element B that combine with 1g of element A can be expressed as a ratio of small whole numbers.
1.33 grams of oxygen with 1 g of carbon = CO
2.67 grams of oxygen with 1 g of carbon = CO2
2.67/1.33 = 2
If you wanted to make CO3 it would take 3.99 grams of oxygen with 1g carbon
3.99/1.33 = 3

8. Atomic theory
John Dalton – In 1808, he connected the previous laws to describe the atom
Each element is composed of tiny indestructible atoms
All atoms of a given element have the same mass and properties that distinguish them from the atoms other elements
Atoms combine in simple whole number ratios
Atoms of one element cannot change into atoms of another element. A chemical reaction can only change how the atoms are bound to each other

9. Discovery of the electron
J. J. Thompson – In the late 1800’s performed experiments with a cathode ray tube
After applying a high voltage between two electrodes, a beam of particles were observed traveling from the negatively charged electrode to the positively charged electrode
These particles were negatively charged

10. Cathode ray continued Opposite charges attract, like charges repel
When electrical and magnetic fields are applied the negative particles were deflected
This allowed for calculation of the charge-to-mass ratio of the cathode ray particles
Electron discovered!!

11. Millikan’s oil drop experiment
Robert Millikan – In 1909, he set up oil drop experiment to determine charge of an electron
Fine oil droplets were sprayed into chamber using an atomizer
Oil droplets then allowed to be pass through a small hole under the influence of gravity into a lower region of the chamber
the droplets obtained electrons, produced from bombarding the air in the chamber with ionization radiation
Using an electric field, he was able to slow or reverse the negatively charged droplets that fell
Able to determine the charge of the electron
Using the charge-to-mass ratio from Thompson, Millikan determined the mass of the electron

12. Oil drop setup

13. Structure of an atom
J. J. Thompson – created the plum-pudding model to explain the atom
Negatively charged electrons were held within a positively charged sphere

14. Rutherford’s gold foil experiment
Ernest Rutherford – In 1909 he attempted to confirm the plum pudding model with a gold foil experiment
He fired large alpha particles at an ultrathin gold foil, expecting the particles to mostly pass through with only some being deflected slightly
Instead, some particles were deflected at large angles, even bouncing back directly at the source

16. Nuclear theory
The results from the experiment lead Rutherford to propose his nuclear theory
1. Most of the atoms mass is contained in a small positively charged core called the nucleus
2. Most of the atom is emptyspace where tiny negatively charged electrons are dispersed
3. The negatively charged electrons balance out the positively charged protons to give an overall neutral atom

17. Completion of nuclear theory
Hydrogen has 1 proton and helium has 2 protons, but helium is 4 times heavier??
Rutherford and James Chadwick explained the mass discrepancy by the presence of neutral particles called neutrons

18. Subatomic particles:the proton
Charge = ×10−19  C
Called “+1” for simplicity
Located in nucleus of atom
Mass = x 10–24 g
Approx. 1 atomic mass unit (amu)
1 amu = 1.66 x 10–24 g

19. The neutron No charge (0 C) Located in nucleus Mass = 1.675 x 10–24 g
Approx. = 1 amu

20. The electron Charge = –1.602 x 10–19 C Called “–1”
Located outside nucleus in an e- “cloud”
Mass = x 10-28g
Approx. = 0 amu

21. Elements!
Every element made of atoms, the number of protons determines the type of element
Number of protons = atomic number of element
Chemical symbol – abbreviation for each element
Usually abbreviation lines up with name, with latin based names being the exception
C = carbon, He = Helium
Pb = lead (latin:plumbum), Au = gold (latin:aurum)
For a neutral atom the protons and electrons are equal
Hydrogen has 1 proton and 1 electron when neutral

22. -- Also number of protons
Atomic number
-- Also number of protons
-- also number of electrons in neutral atom
Atomic symbol
-- 1 or 2 letter symbol
Average atomic weight or
Approximate Mass number
-- Also number of protons + number of neutrons

24. Can also represent by omitting atomic number (Z) since it is implied by the element
Which atom does the picture below represent?

25. Isotopes Atoms of the same element that have different atomic masses
Same number of protons
Different # of neutrons
Percent Abundance - amount of each isotope that occurs naturally

26. Average atomic mass
Weighted average of the mass of naturally occurring isotopes of a particular element reported in atomic mass units
Zi = mass number (mass in amu) of smallest isotope
% abd = percent abundance
[Zi(% abd) + Zi+1(% abd) + Zi+2(% abd) + ……]=Zamuaverage
Example-- [20(0.9092) + 21(0.0026) +22(0.0882)] = amu ~ amu

27. Problems
Lithium has two natural isotopes, 6Li and 7Li, which have percent abundances of 7.5% and 92.5% respectively. What is the average atomic mass of Lithium?

28. Using the percent abundances below, calculate the average atomic mass for Carbon from its three isotopes
C-12 = %
C-13 = 1.110%
C-14 = %
C-12
C-13
C-14

29. 3) Antimony (Sb) has two stable isotopes, 121Sb and 123Sb with masses of amu and amu, respectively. Calculate the percent abundances of these two isotopes

31. Ions
If a neutral atom loses electrons it will gain a positive charge corresponding to the number of electrons lost (anion)
If a neutral atom gains electrons it will gain a negative charge corresponding to the number of electrons lost (cation)
Li  Li+ + 1e-
Ca  Ca2+ + 2e-
F + 1e-  F-
O + 2e-  O2-
If The charge also can be referred to as the oxidation state of the atom
Ca2+ is in the +2 oxidation state

33. Dmitri Mendeleev (1834-1907) Russian chemist
Arranged elements in horizontal rows in order of increasing atomic weight
Started new rows in order to make columns of chemicals with similar characteristics
Left spaces open for elements yet to be discovered

34. Increasing atomic weight when going left to right

35. Similar properties within each column

36. Classification of the Periodic Table
Classification by Physical Properties

37. Metals Shiny Conduct electricity Ductile Can be drawn through wires
Malleable (Shapeable)
High M.P. & B.P
room temp
Except Hg

38. Non-Metals Don’t tend to conduct well Not usually ductile
Tend to be brittle
Low M.P. & B.P.
Many are gases at r.t.

39. Metalloids
Have chemical characteristics in between those of metals and non-metals
Includes elements: B (Boron), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te), Astatine (At)

40. Classification by Chemical Properties

41. Alkali Metals Group 1 (1A) Li, Na, K etc. Soft, shiny metals
Conduct heat and electricity
React violently with H2O
Form H2(g) and (basic) solutions

43. Akali(ne) Earth Metals
Group 2 (2A)
Be, Mg, Ca etc.
Not as reactive as Alkali Metals, but still quite reactive
Tend to make basic solutions when placed in water

44. Transition Metals Groups 3B-8B Tend to have high densities and B.P.
All are metals
Often used for electrical conduction
Often have vivid colors when in solution
Used for pigments

45. Colors of Transition Metal Compounds
Nickel
Cobalt
Copper
Zinc
Iron

46. Lanthanides Elements 57-71 Lanthanum (La) to Lutetium (Lu)
Commonly used in lasers
Can deflect UV and infrared rays

47. Actinides/Actinoids Elements 89-103 Actinium (Ac) to Lawrencium (Lr)
Only Actinium, Thorium (Th), and Uranium (U) occur naturally
Others created by neutron bombardment
Radioactive

48. Groups 3A – 6A No common name
Boundary between metals and non-metals occurs here
Contains the metalloids
Contain elements abundant in earth’s crust, atmosphere, and living things
C, N, O, Si

49. Halogens Group 7A Very reactive with many compounds
F, Cl, Br, I, At
Very reactive with many compounds
Like to form diatomic molecules
F2, Cl2, Br2,

50. Noble Gases Group 8A He, Ne, Ar, Kr, Xe, Rn Very unreactive
Don’t like to bond to other molecules
Generally not abundant

51. Problems Identify an element from each of the groups
1. Noble Gas (example-- He aka Helium)
2. Lathanide
3. Alkali metals
4. Halogens
5. Actinides
6. Alkaline metals
7. Element that is abundant in living organisms
Period and group

52. Chapter 2…. done.