1. CH 2: Atoms, Molecules, and Ions

2. Chapter Outline History of chemistry (2.1 – 2.4)
Chemical laws – start with slide 8 for content, 3-7 FYI slides
Path to the atom
Modern atomic structure (2.5)
Molecules vs. Ions (2.6)
Naming molecular and ionic compounds (2.8)
Introduction to the periodic table (2.7)

3. History of Chemistry
Greek Philosophers: 5th Century BCE (BCE = before the common era - replaces BC)
The Greek philosophers were the first to reflect on the nature of matter. 
They proposed that all matter is made out of first 4 elements -- earth, air , water, fire.  Aristotle added a fifth element, plasma (also called ether). 

4. Greek Philosophers
Democritus had an alternate view of matter.  He proposed that matter was  made up of tiny particles called atoms. 
His "theory" was not well accepted at the time.

5. Alchemy Alchemy: ~600-1600's CE
(CE = common era, replaces AD)
Alchemy developed at about the same time in China, India, and Greece. It spread into Europe in the 8th century.

6. Alchemy Alchemists had two pursuits
Search for a means to convert “base” metals into gold
Search for the elixir of life
Substance that would lead to immortality

7. Alchemy Advances from Alchemy Many new substances where identified
Plaster of Paris, nitric acid….
New lab techniques and equipment developed
New medicines identified

8. Modern Chemistry, ~1600 on
First chemists/physicists to use scientific method
Boyle - elements
Lavoisier – law of conservation of matter
Proust - law of definite proportions
Dalton – law of multiple proportions, atomic theory
Avogadro - hypothesis
Thomson – charge to mass ratio for an electron
Millikan – charge on the electron
Bequerel and the Curies - radioactivity
Rutherford – nuclear atom

9. Modern Chemistry Robert Boyle: ~1660
Proposed a substance to be an element unless it can be broken down into simpler substances.
Proposed one of the gas laws – CH 5
Boyle – studied the relationship between pressure and volume of a gas
Boyle still believed in alchemy – that metals could be converted into gold

10. Lavoisier: ~1760 Law of Conservation of Matter
Matter is neither created nor destroyed in a chemical reaction.
Called the father of modern chemistry
Beheaded for his political believes

11. Proust: late 1700s Law of Definite Composition/Proportions
A given compound always contains the same proportion of elements by mass.

12. Law of Multiple Proportions
John Dalton: ~1800
Law of Multiple Proportions
When two elements form more than one compound, the ratios of the masses of the second element that combines with one gram of the first element can always be reduced to small whole numbers.

13. Law of Multiple Proportions Example
Consider two g samples 2 different compounds containing only C and H. Compound A: 27.2 g of C
Compound B: 42.9 g of C.
Show how this data illustrates the law of multiple proportions.

14. Dalton also proposed the first table of atomic masses
Most masses later need revision
Dalton is best known for proposing Atomic Theory

15. Dalton’s Atomic Theory
Elements are made up of tiny particles called atoms.
Atoms are indivisible and indestructible
Atoms of a given element are identical
Atoms of different elements differ in some fundamental way(s)

16. Dalton’s Atomic Theory
Compounds form when atoms of different elements combine with each other.
A given compound always has the same relative number and types of elements.

17. Dalton’s Atomic Theory
Chemical reactions occur when atoms change how they are bound to each other.
Individual atoms are not changed, just rearranged

18. Avogadro: 1811 Avogadro's Hypothesis
At the same temperature and pressure equal volumes of gases contain the same number of particles.
Based on Guy-Lussac’s data
See page 41/42

19. From Dalton to Atomic Structure
Dalton’s atomic theory lead to much research on the nature of the atom.
This research showed the atom to made up of smaller particles.

20. J.J. Thomson
Thomson measured the deflection of a cathode ray beam in electrical and magnetic fields of known strengths.
Applied electrical field
Cathode ray +
(+)
Metal electrode
(-)
Metal electrode -
Cathode ray tube experiment, pg 43

21. Thomson found the cathode rays were attracted by the positive charge and repelled by negative
These findings clearly indicated that the rays consisted of negatively charged particles.
Today we know these particles as electrons.

22. Thomson measured the deflection of the beam in a magnetic field and more! From his data he determined the charge:mass ratio for an electron
e = x 108 C/g m
e = charge on the electron in coulombs
M = mass of an electron in grams
Cathode ray tube experiment

23. Thomson also found that the cathode ray particles were identical regardless of source.
Concluded all elements contain these negative particles (electrons)

24. J.J. Thomson
Thomson:
identified cathode ray beams as a stream of negatively charged particles
calculated the charge to mass ratio for these negatively charged particles
proposed the existence of positively charged particles
To balance the negative charge of the electrons

25. Millikan ~1909
Millikan’s oil drop experiment allowed him to determine the charge on an electron
This charge can be plugged into Thomson’s formula and the mass of the electron calculated
Mass electron = x kg
Page 44

26. Radioactivity Becquerel, Marie and Pierre Curie: ~1896
Henri Becquerel - observed the natural emission of energy/rays by uranium.
Marie and Pierre Curie studied “Becquerel's rays”.
The Curies’ findings suggested that matter was composed of smaller particles than atoms.
The Curies coined the term radioactivity to describe the rays emitted.

27. Radioactivity Three types of radioactivity were identified:
gamma rays - very high energy light
beta particles - high energy electrons
alpha particles - He+2 particles
2 protons and 2 neutrons

28. Plum Pudding Model of the Atom
In the early 1900’s the accepted model of the atom was called the plum pudding model of the atom
Electrons (tiny and negatively charged) were pictured to be dispersed in a ‘cloud’ of positive charge.
Proposed by JJ Thomson and Lord Kelvin in 1904

29. Rutherford and the Nuclear Atom
In 1911 an experiment conducted in Ernest Rutherford’s lab showed the “plum pudding” model to be incorrect.
Experiment was conducted by Geiger and Marsden and the findings interpreted by Rutherford.
See page 45

30. Rutherford’s Atom First to propose a nuclear atom.
An atom has a dense positive center containing all of positive charge and most of the mass of the atom – the nucleus
Electrons are scattered in the empty space around the nucleus
Electrons occupy a volume that is huge as compared to the size of the nucleus.

31. A New Model of the Atom Expected based on Plum pudding model
Rutherford’s model
Based on ”his” results

32. Modern Atomic Structure
Rutherford continued to study the atom and the positive matter of the atom.
1919, + particle named the proton
~1932 James Chadwick proposed the existence of a third subatomic particle, the neutron.

33. Subatomic Particles Subatomic Particle Charge Mass, amu
Location in atom
Electron
(e-) -1
0 amu
Outside of nucleus
Proton (p) +1
~1 amu
Nucleus
Neutron (n)

34. Mass of Subatomic Particles
Protons and neutrons have ~ the same mass (in the range of kg).
Mass of each and of individual atoms is often expressed in amu rather than grams
Atomic mass unit (amu) – 1/12 the mass of a carbon-12 atom

35. Mass of Subatomic Particles
The mass of the electron (10-31 kg) is tiny as compared to that of the proton and neutron (10-27 kg) .
Therefore, the electron’s mass is considered to be ~0 amu when calculating the mass of an atom.

36. Subatomic Particles and the Elements
Each element has a unique number of protons.
Number of protons defines the element.
Atomic # = # protons 6 C

37. Subatomic Particles
Since atoms are neutral, for every proton there is a/n _________.
When atoms interact to form compounds, it is their ___________ that interact.

38. Terms
Mass number = sum of the # of protons and the # neutrons in the nucleus of an atom
FOR MOST ELEMENTS THE MASS NUMBER IF NOT ON THE PERIODIC TABLE.
You will be given enough information to determine mass number or number of neutrons.

39. Terms Isotopes = atoms of a given element that differ in mass number
Isotopes have the same number of _____________.
Isotopes differ in the number of _______.

40. B Isotopes Writing atomic symbols for isotopes pg 46 11 Mass #
Symbol for element 5
Atomic #

41. FAQ - Isotopes C Po When is mass number found on the periodic table?
What’s the atomic mass? Is it the same as the mass number? C
(209) Po

42. Molecules and Ions (2.6)
Atoms of different elements combine to form compounds
Atoms in compounds are held together by chemical bonds.
Bonds involve interactions of the bonding atoms’ ________

43. Bonding There are two types of bonds:
Covalent bonds – bonding atoms share electrons
Atoms are always nonmetal atoms
Covalently bonded atoms form molecules
Ways to represent molecules
Chemical formula; H2O
Structural formula O
H H

44. Bonding Ionic bonds – attractive force among oppositely charged ions
Bond formed between metal cations and nonmetal anions
No molecules involved

45. Ions - Terms Ion – charged atom or group of atoms
Formed when atoms gain or lose electrons
Cation – positively charged ion
Formed when an atom _______ electrons
Anion – negatively charged ion
Formed when an atom ______ electrons

46. Ions
Describing ion formation
Cation example:
Anion example:

47. Naming Binary Compounds
Binary compounds – compound composed of 2 elements
NaCl CO
CO2

48. Types of Binary Compounds
Type I binary ionic compounds
Metal forms only one ion
Type II binary ionic compounds
Metal forms more than one ion
Use roman numerals to indicate the charge on the ion
Type III binary covalent compounds
Compound between 2 nonmetals

49. Types I Binary Compounds
Compound between a metal and a nonmetal
Metal forms only one ion
Name the cation and then the anion.
Name of the cation is the name of the element
Name of the anion is the name of the nonmetal with the ending changed to “ide”

50. Monoatomic cations to know
Group #
Charge on ion
examples IA +1
Na1+ sodium (ion)
K potassium (ion)
IIA +2
Mg2+ magnesium (ion)
IIIA metals +3
Al3+ aluminum (ion)

51. Monoatomic anions to know
Group #
Charge on ion
examples VA -3
N3- nitride (ion)
P3- phosphide (ion)
VIA -2
O2- oxide (ion)
S sulfide
VIIA -1
F1- fluoride (ion)
Cl1- chloride (ion)
Br1- bromide (ion)
I iodide (ion)

52. Practice
Name  chemical formula
Chemical formula  name