1. WARM-UP
Studying atoms is difficult because they are too small to see or directly observe even with the best scientific tools. Write a similar example of something that can not be studied directly.

2. Chapter 4: Atomic Structure
Atom- The basic building block of matter
Learning Targets-
I can describe the structure of atoms.
I can describe how structure of an atom affects it’s properties.
I can create a timeline that shows the developments that lead to the current model of the atom.

4. Ancient Greek Models of atoms
BC Democratis- all matter consisted of extremely small particles that could not be divided he called them atoms from the Greek word Atomos which means “uncut” or “indivisiable” BC Aristotle- no limit to the number of times that matter could be divided Accepted until the 1800s

5. The existence of atoms wasn’t scientifically proven until the early 1800s.
John Dalton, , English chemist and teacher
Studied the behavior of gases in air
Concluded that gas contains individual particles
No matter how large or small the sample the ratio of the elements in compounds is always the same

6. Dalton’s atomic theory
All elements are composed of atoms
All atoms of the same element have the same mass and atoms of different elements have different masses
Compounds contain atoms of more than one element
In a particular compound, atoms of different elements always combine in the same way

7. Daltons atomic model Tiny solid spheres with different masses
Dalton’s theory explained data from many experiments and thus became widely accepted

8. Although atoms are incredibly small they can now be observed with a scanning tunneling microscope

9. Charged materials
Some materials when rubbed gain the ability to attract or repel other materials
Gain either a positive or negative charge
Object with like charges repel or push apart, objects with opposite charges attract or pull together
Charged particles can flow creating an electric current

10. Thomson’s model of an atom
J.J. Thomson, an English physicist
Wires connect the metal disks at opposite ends of a empty glass tube, one disk becomes negatively charged and one becomes positively charged
A glowing beam appears in the space between the plates
The beam is repelled by a negatively charged metal disk or attracted by a positively charged plates brought near it

11. Thomson’s model
Thomson hypothesized the ray was a stream of negatively charged particles contained inside atoms, now called electrons, which are part of all atoms and carry a charge of -1. No matter what metal he used he got the same particles The mass was always 2000 times smaller than the mass of hydrogen atoms (a proton)

12. Thomson’s model
First evidence that atoms are made of even smaller particles
Atom is neutral- negative charges scattered throughout an atom filled with a positively charged mass of matter

13. Ernest Rutherford’s Gold Foil Experiment
Atom was believed to have its positive charge spread throughout.
Rutherford shot alpha particles (large 2 + atoms) at a very thin sheet of gold foil.
If the current model of the atom was correct the alpha particles should pass though gold the mass and charge being too small to deflect the alpha particles.

14. Rutherford’s atomic theory
Most alpha particles actually passed straight through a small fraction bounced off the gold foil at large angles
Atoms are mostly empty space
Positive charge is concentrated in the nucleus, not evenly distributed, which contains the protons and neutrons and has a positive charge
Positive charge varies among elements
Each nucleus must contain at least one proton, each proton is assigned a charge of +1

15. Rutherford’s atomic model
All of an atoms positive charge is concentrated in the dense nucleus Electrons are outside the nucleus

16. James chadwick
1932 James Chadwick, English physicist did experiments that proved the neutrons existed Concluded that they were neutral because they were not effected by a charged particle Neutrons are contained in the nucleus and have a mass equal to that of a proton
Compare the mass, location and charge of protons, neutrons, and electrons.

17. Rutherford's atomic model couldn’t explain chemical properties of elements
required knowledge of electron behavior
Niels Bohr ( )
Focused on Electrons
Agreed with Rutherford that the nucleus of and atom was surrounded by a large volume of space

18. Electrons are only found in specific circular paths- orbits around the nucleus
Each electron orbit has a fixed energy or energy level
An electron cannot exist between energy levels
The energy level closest to the nucleus is the lowest
An electron in an atom can move from one energy level to the next when it gains or loses energy
Energy lost can be in the form of light
The Bohr Model

19. The Quantum Mechanical Model
Rutherford-Bohr described the path of an electron as a large object would behave which was inconsistent with theoretical calculations and experimental results
Electrons move in a much less predictable way then planets in a solar system.
Erwin Schrodinger
Devised and solved a mathematical equation describing the behavior of the electron in hydrogen atom

20. Quantum Mechanical (Electron Cloud) Model
Electron Cloud- visual model of the most likely locations for electrons in an atom
Based on the probability of finding an electron with in a certain volume of space surrounding the nucleus is described as a fuzzy cloud where the electron is 90% of the time
More dense- probability high
Less dense- probability low

21. Atomic orbitals
The electron cloud represents all the orbitals in an atom
An orbital is a region of space around the nucleus where an electron is likely to be found
Each orbital can hold 2 electrons
The lowest energy level has one orbital, the second has four, the third 9 and the fourth 16
Electron configuration- the arrangement of electrons in the orbitals of an atom
The most stable electron configuration is the one in which the electrons are in orbitals with the lowest possible energies this is called the ground state

22. Distinguishing Among Atoms
How are atoms of Hydrogen different than atoms of oxygen? Element of different atoms are contain different numbers of protons

23. Atomic Number- the number of protons in the nucleus of an atom of a given element
Identifies an Element
Atoms are electrically neutral so the number of electrons are also equal to the atomic number

24. Au Elements can be represented in the following shorthand notation:
Mass Number- the number of protons and neutrons in the nucleus
Example- carbon has 6 protons and 6 neutrons so the mass number is 12
Most the mass of an atom is concentrated in the nucleus
The number of neutrons in an atom is the difference between the mass number and the atomic number.
# neutrons = mass # - atomic #
symbol
Mass # Au
197 79
Atomic #
Or Gold-197

25. Isotopes- atoms that have the same number of protons but a different number of neutrons
Different mass numbers
Chemically alike- same protons and electrons which are responsible for chemical behavior

26. Atomic Mass
The mass of atoms are givin in comparison to carbon-12
An atomic mass unit (amu) = 1/12 the mass of a carbon-12 atom
Carbon 12 amu
Flourine- actual mass x 10 –23 g, atomic mass amu
1 proton or neutron- 1 amu
Atomic mass- weighted average mass and relative abundance of isotopes as they occur in nature
Example- Hydrogen –1, 99 %,
Hydrogen- 2, 1% heaver
Hydrogen- 3
Atomic mass