1. Inside the Atom

2. Atoms
The word atom comes from the Greek atomos, meaning “indivisible.”
You cannot divide a carbon atom into smaller pieces and still have carbon.
Atoms compose all ordinary matter—if you want to understand matter, you must begin by understanding atoms.

3. Experiencing Atoms
Atoms are incredibly small, yet they compose everything.
Atoms are the pieces of elements.
Properties of the atoms determine the properties of the elements.
Tro's "Introductory Chemistry", Chapter 4

4. Experiencing Atoms There are about 91 elements found in nature.
Over 20 have been made in laboratories.
Each has its own, unique kind of atom.
They have different structures.
Therefore they have different properties.
Tro's "Introductory Chemistry", Chapter 4

5. Imaging and Moving Individual Atoms
In spite of their small size, atoms are the key to connecting the macroscopic and microscopic worlds.
An atom is the smallest identifiable unit of an element.
There are about
91 different naturally occurring elements, and
over 20 synthetic elements (elements not found in nature).

6. Early Ideas about the Building Blocks of Matter
Leucippus (fifth century b.c.) and his student Democritus (460– 370 b.c.) were first to propose that matter was composed of small, indestructible particles.
Democritus wrote, “Nothing exists except atoms and empty space; everything else is opinion.”
They proposed that many different kinds of atoms existed, each different in shape and size, and that they moved randomly through empty space.

7. Early Building Blocks of Matter Ideas
Plato and Aristotle did not embrace the atomic ideas of Leucippus and Democritus.
They held that
matter had no smallest parts.
different substances were composed of various proportions of fire, air, earth, and water.

8. Early Building Blocks of Matter Ideas
Later scientific approach became the established way to learn about the physical world.
An English chemist, John Dalton (1766–1844) offered convincing evidence that supported the early atomic ideas of Leucippus and Democritus.

10. Modern Atomic Theory and the Laws That Led to It
The theory that all matter is composed of atoms grew out of observations and laws.
The three most important laws that led to the development and acceptance of the atomic theory are as follows:
The law of conservation of mass
The law of definite proportions
The law of multiple proportions

11. The Law of Conservation of Mass
Antoine Lavoisier formulated the law of conservation of mass, which states the following:
In a chemical reaction, matter is neither created nor destroyed.
Hence, when a chemical reaction occurs, the total mass of the substances involved in the reaction does not change.
This law is consistent with the idea that matter is composed of small, indestructible particles.

12. The Law of Definite Proportions
In 1797, a French chemist, Joseph Proust made observations on the composition of compounds.
He summarized his observations in the law of definite proportions:
All samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements.
The law of definite proportions is sometimes called the law of constant composition.
For example, the decomposition of 18.0 g of water results in 16.0 g of oxygen and 2.0 g of hydrogen, or an oxygen-to-hydrogen mass ratio of:

13. The Law of Multiple Proportions
In 1804, John Dalton published his law of multiple proportions.
When two elements (call them A and B) form two different compounds, the masses of element B that combine with 1 g of element A can be expressed as a ratio of small whole numbers.
An atom of A combines with either one, two, three, or more atoms of B (AB1, AB2, AB3, etc.).
Carbon monoxide and carbon dioxide are two compounds composed of the same two elements: carbon and oxygen.
The mass ratio of oxygen to carbon in carbon dioxide is 2.67:1; therefore, 2.67 g of oxygen reacts with 1 g of carbon.
In carbon monoxide, however, the mass ratio of oxygen to carbon is 1.33:1, or 1.33 g of oxygen to every 1 g of carbon.

14. The ratio of these two masses is itself a small whole number.

15. John Dalton and the Atomic Theory
Dalton’s atomic theory explained the laws as follows:
1. Each element is composed of tiny, indestructible particles called atoms.
2. All atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements.
3. Atoms combine in simple, whole-number ratios to form compounds.
4. Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms only change the way that they are bound together with other atoms.

17. Parts of the Atom

18. The Parts of an Atom Atoms themselves come with several smaller parts
Atoms consist of subatomic particles
Names of the subatomic particles
Electron
Proton
Neutron

19. Understanding the Geography of the Atom
Atoms are mostly empty space
Virtually % of atoms mass is found in the nucleus
What is the nucleus?
Center of the atom
Very small in volume compared to atoms
Contains protons and neutrons

20. Particles that live in the nucleus
Protons: positively charged particles
Neutrons: particles that have the same mass as protons yet are electrically neutral
Protons and Neutrons EACH weigh 1.67 x kg

21. The Electron Discovered by J.J. Thompson Negatively charged particle
Found orbiting around the nucleus on specific energy levels

22. The Modern Atom
We know atoms are composed of three main pieces—protons, neutrons, and electrons.
The nucleus contains protons and neutrons.
The nucleus is only about cm in diameter.
The electrons move outside the nucleus with an average distance of about 10-8 cm.
Therefore, the radius of the atom is about 105 times larger than the radius of the nucleus.
Tro's "Introductory Chemistry", Chapter 4

23. The Nature of Electrical Charge
Electrical charge is a fundamental property of protons and electrons.
Positively and negatively charged objects attract each other.
Like charged objects repel each other.
+ to +, or  to .
When a proton and electron are paired, the result is a neutral charge.
Because they have equal amounts of charge.
Tro's "Introductory Chemistry", Chapter 4

24. Subatomic Particles
The charge of the proton and the electron are equal in magnitude but opposite in sign. The neutron has no charge.

25. A model for the Atom

27. Significance of the Nucleus
Each element on the periodic table has a different number of protons
Each nucleus has a different mass due to differing numbers of protons and neutrons

28. What’s in the nucleus? Strong Nuclear Force: Atomic Number (Z)
# of protons in the nucleus of an atom
Atomic Mass Number (A)
# of protons and neutrons COMBINED in the nucleus of an atom
Strong binding force within the nucleus caused by the addition of neutrons to positive protons, which allows the nucleus be so tightly packed.
Strong Nuclear Force:

29. Reading an element on the periodic table
Element name
Atomic Number
Element Symbol
Average Atomic
Mass

30. How elements are written
Element name – mass
This tells you the name of the element on the periodic table
This tells you the atomic mass of the element
Example
Bromine-80
Bromine has an atomic mass of 80
Bromine has an atomic # of 35
Bromine therefore has 35 protons, 35 electrons, and 45 neutrons

31. How elements are written: Practice
Give the # of protons, neutrons and electrons for:
Carbon-12 has
6 protons, 6 electrons, 6 neutrons
Sodium-23 has
11 protons, 11 electrons, 12 neutrons
Lithium-7 has
3 protons, 3 electrons, 4 neutrons
Boron-11 has
5 protons, 5 electrons, 6 neutrons

32. Writing The Symbol Way Mass number atomic number Elementcharge Example
3919 K
Potassium has 19+’s, 19-’s, 20 neutrons and a mass of 39
2713 Al
Aluminum has 13+’s, 13-’s, 14 neutrons and a mass of 27
23892 U+1
Uranium has 92+’s, 91-’s, 146 neutrons and a mass of 238 Ag
Silver has 47+’s, 47-’s, 60 neutrons and a mass of 107
3216 S-2
Sulfur has 16+’s, 18-’s, 16 neutrons and a mass of 32

33. Atomic Mass: The Average Mass of an Element’s Atoms
Atomic mass is sometimes called atomic weight or standard atomic weight.
The atomic mass of each element is directly beneath the element’s symbol in the periodic table.
It represents the average mass of the isotopes that compose that element, weighted according to the natural abundance of each isotope.

34. Elements of different mass numbers
Isotope: elements with the same number of protons but differing numbers of neutrons
Example 1:
Example 2:

35. Atomic Mass
Naturally occurring chlorine consists of 75.77% chlorine-35 atoms (mass amu) and 24.23% chlorine-37 atoms (mass amu). We can calculate its atomic mass:
Solution:
Convert the percent abundance to decimal form and multiply it with its isotopic mass:
Cl-37 = (36.97 amu) = amu
Cl-35 = (34.97 amu) = amu
Atomic Mass Cl = = amu

36. Atomic Mass
In general, we calculate the atomic mass with the equation:

38. Lab: Chi-Square Analysis of M&M’s
Introduction: how many of you like M&M’s?
How many of you are curious about the colors of M&M’s in the bag?

39. What if I told you
Each bag of M&M’s is designed to have a preset amount of M’s based on statistics.
Acquire worksheet for a group of 4 students
Prepare a null hypothesis: your hypothesis should state whether or not you feel your bag of M&M’s will be similar to what is listed in the table above. We are using plain M&M’s.

40. The Experiment
Each team will get a separate bag of M&M’s from Professor Berger
Open the bag and spread/sort the M&M’s into piles separating the different colors
Count the number of each color present and record on your data table

41. Statistically Speaking
Lets look at a spreadsheet perhaps

42. The rest of the chart
Difference: get the net positive answer for your difference
Difference squared: square the difference
Take difference squared divided by expected
For total column add the 6 column’s difference squared divided by expected to get the total

43. Degrees of Freedom Deals with the # of categories – 1 you have
How many degrees of freedom do we have?
Compare your chi-square (total box) and see which probability it falls into. Do you accept or reject the null hypothesis as a result?

44. Accept or Reject
If you said your bag would follow the manufacturer’s data:
Accept null hypothesis if chi-squre is <9.24
Reject null hypothesis is chi-square is >9.24
If you said your bag would NOT follow the manufacturer’s data:
Accept null hypothesis if chi-squre is >9.24
Reject null hypothesis is chi-square is <9.24
Answer the questions below on your worksheet

45. What about all the bags?
Professor Berger’s been working hard too you know ;)
Copy the table posted on the screen
Answer the questions below on your page, turn in paper when you’re done. Make sure all group members name’s are on the first page.
Congrats you have finished lab #1