1. Atoms: The Building Blocks of Matter
MC – Ch. 3

2. 3.1 The Atom: From Philosophical Idea to Scientific Theory

3. (ideas based on observations) Named them ‘atoms’
Democritus
400 BCE – Greek philosopher Democritus Believed that matter was made of small indivisible particles and empty space
(ideas based on observations)
Named them ‘atoms’
Aristotle
who did not believe in particles, but who thought
matter was continuous with no
indivisible particles
Aristotle “won” the debate with Democritus. His idea was accepted until the 1700’s

4. Experimental evidence for atomic theory
Antoine Lavoisier – French
chemist in the 1780’s, observed
that mass before and after a
reaction remained the same
Law of conservation of mass: mass is neither created nor destroyed during ordinary chemical reactions or physical changes.

5. Joseph Proust – French chemist ~ 1800
Law of Definite Proportions: any sample of a chemical compound contains the same elements in exactly the same proportions by mass
Example – H2O is always 11% hydrogen and 89% oxygen

6. Law of Multiple Proportions: If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole
numbers.
John Dalton
Began teaching at the age of 12
Studied colorblindness – also called Daltonism
Main focus was meteorology

7. Modern Atomic Theory
John Dalton (English, ) – school teacher and chemist
Revised Democritus’ ideas based on scientific research
Proposed atomic theory in 1808
Best explanation of research
Began teaching at the age of 12
Studied colorblindness – also called Daltonism
Main focus was meteorology

8. Dalton’s Atomic Theory
1. All element are composed of small, indivisible particles called atoms.
2. Atoms of the same element are identical. Atoms of different elements are different.
3. Atoms combine in simple, whole-number ratios to form compounds
4. Chemical reactions occur, but atoms are never changed into atoms of another element

9. Modifications to Dalton’s Theory
We now know that atoms can be divided into subatomic particles. (P, N, E)
We also know that atoms of an element can have slightly different masses. (isotopes)

10. Jons Jakob Berzelius Swedish Chemist – 1811
Established modern system of chemical symbols for elements
Berzelius's work with atomic weights and his theory of electrochemical dualism led to his development of a modern system of chemical formula notation that could portray the composition of any compound both qualitatively (by showing its electrochemically opposing ingredients) and quantitatively (by showing the proportions in which the ingredients were united). His system abbreviated the Latin names of the elements with one or two letters and applied superscripts to designate the number of atoms of each element present in both the acidic and basic ingredients
Berzelius is credited with identifying the chemical elements silicon, selenium, thorium, and cerium. Students working in Berzelius's laboratory also discovered lithium and vanadium.
Berzelius is credited with identifying the chemical elements silicon, selenium, 
thorium, and cerium.

11. Section 4.2 Subatomic Particles and the Nuclear Atom
JJ Thompson – English physicist, late 1800’s/early 1900’s
Using Crookes’ tubes (cathode ray tubes), he discovered the electron

12. Thompson Continued Determined: Plum Pudding Model of Atoms
1. electrons are negatively charged particles
2. all electrons are identical
3. the charge to mass ratio using deflection of rays
Plum Pudding Model of Atoms
the atom is a sphere of positive
matter in which electrons are
embedded.

13. Robert Millikan-American physicist 1909
Oil Drop experiment
Determined the exact charge of an electron,
now called -1
Calculated the mass of an electron
9.1 x g = 1/1840 of
the mass of a hydrogen
atom (amu)

14. Eugen Goldstein – German physicist – 1886
Discovered canal rays – also called anode rays.
This led to the development of mass spectrometry

15. Wilhelm Wien German physicist – 1898
Discovered the proton. It has a +1 charge (equal and opposite to the electron)
It has a mass of 1 amu
While studying streams of ionized gas, Wien, in 1898, identified a positive particle equal in mass to the hydrogen atom. Wien, with this work, laid the foundation of mass spectrometry. J. J. Thomson refined Wien's apparatus and conducted further experiments in 1913 then, after work by Ernest Rutherford in 1919, Wien's particle was accepted and named the proton

16. The Nuclear Atom
Ernest Rutherford (from New Zealand, , working in England,)
Performed the Gold Foil Experiment with the help of Geiger and Marsden.
Alpha particles were “shot” at a thin sheet of gold foil.

17. Results of Gold Foil:
1. Most particles (93%) went straight through the foil
2. Some particles were slightly deflected
3. About 1 in 8000 particles came straight back

18. Conclusions from Gold Foil
1. The atom is mainly empty space
2. The atom contains a small, dense positively charged core called the nucleus
**Rutherford also predicted the existence of neutrons.
- the mass of the protons in the nucleus was only about ½ of the known mass of the nucleus. He predicted a neutral particle must exist in the nucleus.

19. Chadwick (English physicist, 1932)
Identified the neutron (charge of 0) as the other particle in the nucleus
Mass nearly equals proton (1amu)
Worked with Geiger and Rutherford
He was the head of the British team that worked on the Manhattan Project during the Second World War.
He was the head of the British team that worked on the Manhattan Project during the Second World War.

20. Summary of Atomic Model
Consists of electrons, protons, and neutrons
Most of volume consists of electrons moving through empty space around nucleus
Nucleus consists of protons and neutrons
Contains 99.97% of atom’s mass
Number of protons equals number of electrons, hence the atom is neutral

21. Parts of the Atom Protons - Positively charged In the nucleus
Contribute to the atomic mass
Determine the atomic number
Never gained or lost

22. Parts of the Atom Electrons – Negatively charged Outside the nucleus
In different energy levels, each row on the periodic table is a new energy level
Do not contribute to atomic mass
A proton’s mass is 1800 x’s greater than that of an electron
Gained, lost or shared in bonding

23. Parts of the Atom Neutrons – Neutral charge (no charge) In the nucleus
Contribute to atomic mass
Never gained or lost
Different isotopes of the same element have a different number of neutrons

24. Section 3:3 – Counting Atoms
Atomic Number -
Equal to the number of protons
Identifies the element
Is the whole number on the periodic table
(also equal to the # of electrons in the elemental state)

25. Mass NUmber Mass Number (also called atomic mass number) –
Equal to the atomic number + the number of neutrons (p + n)
Is a whole number, is not shown on periodic table

26. Average Atomic Mass Average Atomic Mass –
A weighted average of the masses of all the isotopes for a given element
Ex. H-1.01, Cl-35.45

27. Isotopes
Elements that contain the same number of protons but a different number of neutrons

28. Practice 28 protons, 30 neutrons B. Tungsten – 186 C. Pb, 124 neutrons
Determine the following
for each atom:
Element name
Element symbol
Atomic number
Atomic mass number
# of protons
# of electrons
# of neutrons

29. 28 protons, 30 neutrons
Nickel Ni 28 58 30

30. Tungsten – 186
1. Tungsten
2. W
3. 74
5. 74
6. 74

31. Pb, 124 neutrons
Lead Pb 82
206
124