1. Test 4: Chapter 4 – Atomic Structure
Chemistry 2009

2. The Atom Video
Watch the video and answer questions that follow along on a worksheet.
Workbook p.19

3. Discovery Lab – Observing Electrical Charge
Read the discovery lab on p.87.
Copy and complete procedures 2-3. Record your observations.
Copy and answer the analysis questions.
Turn in your lab report to the bin.

4. Section 1: Early Theories of Matter
Democritus ( BC) – Greek Philosopher
John Dalton ( ) – English School Teacher
J. J. Thomson ( ) – English Physicist
Robert Millikan ( ) – American Physicist
Ernest Rutherford ( ) – English Physicist
Henry Moseley ( ) – English Scientist
James Chadwick ( ) – English Physicist

5. Democritus
1st known person to propose the idea that matter was not infinitely divisible
Ahead of his time, he was criticized by Aristotle
Was a philosopher, not a scientist
Main Ideas of Theory:
Matter is composed of empty space through which atoms move.
Different kinds of atoms have different sizes and shapes.
The differing properties of matter are due to size and, shape, and movement of atoms.
Apparent changes in matter result from the changes in groupings of atoms and not from changes in the atoms themselves.

6. John Dalton Revised Democritus’ ideas based on his own research.
Main Ideas:
All matter is composed of extremely small particles called atoms.
All atoms of a given element are identical, having the same size, mass, and chemical properties,
Atoms of a specific element are different from atoms of any other element.
Atoms cannot be created or destroyed. Atoms arte indivisible.
Different atoms combine in simple, whole number ratios to form compounds.
In a chemical reaction, atoms are separated, combined, or rearranged.

7. Democritus vs. Dalton Similarities: Differences:
Matter is composed of atoms.
Democritus states that matter is composed of empty space through which atoms move, Dalton makes no such claims.
Atoms are indestructible and indivisible.
Changes in matter are due to changes in groupings of atoms.

8. Were they right? Yes and No!
Wrong -
Atoms are divisible. They can be broken down into smaller particles – protons, neutrons, electrons, quarks, etc.)
Atoms of the same element can have slightly different masses. Isotopes of an element have differing numbers of neutrons, which give them different masses.
Right –
Matter is composed of small particles called atoms.
Atoms of a specific element are different from atoms of any other element.
Atoms combine in simple, whole number ratios to form compounds.

9. Our definition of an atom:
The smallest particle of an element that still retains all of the properties of the element.

10. Section 2:
Subatomic Particles and the Nuclear Atom

11. Discovering the Electron
Cathode ray – a ray of radiation that originates from the cathode and travels to the anode of a cathode ray tube.
Discovered by English Physicist Sir William Crookes
Led to the invention of television
By the end of the 1800’s scientists were convinced:
Cathode rays were actually a stream of charged particle.
The particles carried a negative charge.
Late 1800’s J.J. Thomson identified the particles as the first subatomic particle, the electron.
1909 – Robert Millikan determine the charge as -1 and the mass to be 9.1 x grams or 1/1840 the mass of a hydrogen atom.

12. Electron
A negatively charged, fast-moving particle with an extremely small mass that is found in all forms of matter and moves through the empty space surrounding an atom’s nucleus.

13. Plum-Pudding Atomic Model
JJ Thomson
Negatively charged electrons are distributed throughout a uniform positive charge.
Picture chocolate chip cookie dough (chocolate chips = electrons; dough = positive charge)

14. The Nuclear Model
1911 – Ernest Rutherford – Gold Foil Experiment (p.95 fig 4.11)
A narrow beam of alpha particles was aimed at a thin sheet of gold foil. A zinc-sulfide coated screen surrounding the gold foil produced a flash of light whenever it was struck by an alpha particle.
Results showed most alpha particles passing through the gold foil with little or no deflection, but some were deflected at small and large angles.
Rutherford then concluded that an atom existed of mostly empty space through which the electrons moved and a tiny, dense region centrally located within the atom that contained all of the atom’s positive charge and virtually all of its mass, which he called the nucleus.

15. 1920 – Rutherford refined the concept of the nucleus and concluded that it contained positively charged particles called protons. The charge on the proton (1+) was equal to but opposite that of an electron.
1932 – James Chadwick showed that the nucleus contained another subatomic particle that was neutral. A neutron has a mass nearly equal to that of a proton, but carries no electrical charge.

16. Current Atomic Model Electron Cloud Model
Atoms consist of a cloud of fast moving, negatively charged electrons surrounding a tiny, extremely dense nucleus containing positively charged protons and neutral neutrons.

17. Properties of Subatomic Particles
Know this table, it can also be found on p.97 in your text.
Particle
Symbol
Location
Relative Electrical Charge
Relative Mass
Actual Mass (g)
Electron e-
Space surrounding the nucleus 1-
1/1840
9.11 X 10-28
Proton p+
Nucleus 1+ 1
1.673 x 10-24
Neutron n0
nucleus
No Charge
1.675 x 10-24

18. Section 3:
How Atoms Differ

19. Atomic Number
1910’s – Henry Moseley discovered that atoms of each element contain a unique positive charge in their nuclei. (The number of protons in an atom identifies it as an atoms of a particular element.)
Atomic Number – the number of protons in an atom of a given element.
Example: Hydrogen
Atomic Number = 1 SO, #p+=1
Remember, all atoms are neutral, so… the number of protons must be equal to the number of electrons.
Atomic Number = #p+ = #e-

20. Isotopes
All atoms of a particular element must have the same number of protons and electrons, but the numbers of neutrons on the nuclei may differ.
Isotopes – atoms with the same number of protons, but different numbers of neutrons.
Most elements are found as a mixture of isotopes, but the relative abundance of each isotope is usually constant.
Isotopes differ in mass, but generally have the same chemical behaviors.

21. Mass Numbers
Mass Number – the sum of the protons and neutrons in the nucleus
To identify isotopes the mass number follows the element’s name.
Example- Potassium-39 and Potassium-40
#n0 = Mass Number - Atomic Number

22. Mass of Individual Atoms
Because of the extremely small masses of the subatomic particles, chemists have developed a method of measuring the mass of an atoms relative to the mass of a specifically chosen standard – the carbon-12 atom.
1 carbon-12 atom = 12 atomic mass units
1 atomic mass unit (amu) is defined as 1/12th the mass of a carbon-12 atom.
Atomic Mass – is the weighted average mass of the isotopes of that element.

23. Class Assignment: Copy and complete practice problems 11-13 on p.99.
Copy and complete practice problem 14 on p.101.
Copy and complete practice problems on p.104

24. Lab Activities
MiniLab p.102 – Modeling Isotopes: Copy and complete lab. Answer analysis questions.
ChemLab p.109 – Very Small Particles: Copy and complete the lab. Answer Analyze and Conclude questions. Answer Real-World Chemistry questions.
Turn in both labs when complete.

25. Homework:
Workbook p.20-23