1. Chapter 5 Atomic Structure & The Periodic Table
College Prep Chemistry
Mrs. Lips
WOC
The Atom
Ancient Ideas

2. The Atom: From Philosophical Idea to Scientific Theory
Democritus (Greek philospher; 450 BCE)
First suggested the idea that atoms
existed but had no experimental support
Atoms: the smallest particle of an element that retains the properties of that element.
(Greek: atomos = indivisible)
Although Democritus’ ideas agreed with later scientific theory, they were not useful in explaining chemical behavior. They also lacked experimental support because scientific testing was unknown at the time. The real nature of atoms and the connection between observable changes and events at the atomic level were not established for more than 2000 years after Democritus. The modern process of discovery of atoms began with John Dalton.

3. John Dalton (1766-1844) English school teacher
Performed experiments and came up with 5 hypotheses to explain his results (Dalton’s Atomic Theory)

4. Dalton’s Atomic Theory
1. All elements are composed of tiny indivisible particles called atoms.
2. Atoms of the same element are identical. The atoms of one element are different from those of another.

5. Dalton’s Atomic Theory (cont’d)
3. Atoms of different elements can combine with one another in simple whole number ratios.
H2O C12H22O NOT H2.5O¾
4. Chemical reactions occur when atoms are separated, joined or rearranged. Atoms are not changed into atoms of another!

6. Dalton’s Atomic Theory
5. Atoms can not be subdivided.
Today, we know they can be divided into subatomic particles

7. Dalton’s atom A simple sphere with no internal structure
The way compounds were held together was poorly understood

8. How small is an atom? Average atom size: Diameter = 1 x 10-8 cm
Mass = 1 x 10 –23 g
100,000,000 copper atoms in a row would = 1 cm in length!
Atomic Size

9. Changes to Dalton’s Atomic Theory
Most of Dalton’s Atomic Theory is accepted
One major revision includes the idea that atoms have smaller parts…
There are 3 parts to an atom….
1. electrons
2. protons
3. neutrons

10. The Electron (e-) Discovered by J.J. Thomson (1897) “Cathode Ray”
Negatively charged subatomic particles called electrons responsible for the ray were attracted to the positive end of a magnet
Plum Pudding Model
The debate was resolved in 1897 when J. J. Thomson measured the mass of cathode rays, showing they were made of particles, but were around 1800 times lighter than the lightest atom, hydrogen. Therefore they were not atoms, but a new particle, the first subatomic particle to be discovered, which he originally called "corpuscle" but was later named electron, after particles postulated by George Johnstone Stoney in He also showed they were identical with particles given off by photoelectric and radioactive materials.[4] It was quickly recognised that they are the particles that carry electric currents in metal wires, and carry the negative electric charge of the atom.

11. The Electron (e-) continued
Robert Milikan (1916)
Determined the charge and mass of an electron
1 unit of negative charge
mass about 1800 times smaller than the mass of a hydrogen atom

12. 2. The Proton (p+) Discovered by Goldstein (1886)
Positively charged subatomic particles called protons
1,840 times heavier than an electron

13. 3. The Neutron (no) discovered by James Chadwick (1932)
Subatomic particles with no charge called neutrons
Mass is nearly the same as a proton

14. Summary of Subatomic Particles
Symbol
Relative Charge
Approx. Relative Mass (amu)
Actual Mass (g)
Electron e- 1-
1/1840
9.11x10-28
Proton p+ 1+ 1
1.62x10-24
neutron no

15. Discovery of the Atomic Nucleus
Gold Foil Experiment
Ernest Rutherford
nucleus of the atom is a small dense positively charged region in the center of an atom (contains almost all the mass of the atom)
Their test used massive alpha particles (helium atoms that have lost their two electrons and have a double positive charge because of the two remaining protons. Rutherford directed a narrow beam of alpha particles at a very thin sheet of gold foil. According to the prevailing plum pudding theory of atomic structure, the alpha particles should have passed easily through the gold with only a slight defelection due to the positive charge though to be spread out in the gold atoms.
BBC Video
Gold Foil Expt

16. Rutherford Model of the Atom
Planetary model – electrons orbit the nucleus like the planets orbit the sun
Atom is mostly empty space
The nucleus is tiny compared to the rest of the atom (like a marble in a stadium)

17. Bohr Model of the Atom Niehls Bohr
Atom has a nucleus but electrons orbit in definite energy levels

19. Atomic Number
the number of protons in the nucleus of an atom of an element
Because Atoms are electrically neutral
Atomic # also = #e
Periodic Table
#6 – Carbon:
has 6 p+ and 6 e-
#1 – Hydrogen:
has 1 p+ and 1 e-

20. Mass Number Periodic Table
Mass Number –total # of protons and neutrons in an atom’s nucleus
# of neutrons = mass # atomic #
= (# p+ + # no) - (# p+)
Periodic Table
Carbon:
6 p+ and 6 e-
Also has 6 n0

21. Determining # of particles
Atomic # = number of protons
Number of protons = number of electrons (in a neutral atom)
Mass number = # of protons + # of neutrons

22. Atomic Notation
Mass number 16 O 8
Atomic number
8 p+ , 8 n0 , 8 e-

23. Practice… Element: ______________ Atomic #: ______________
Mass #: _______________
Protons: _______________
Neutrons: ______________
Electrons: ______________
More Practice

24. Isotopes
Definition – atoms that have the same number of protons but different numbers of neutrons
Identified by mass number
Atomic # (# of protons) stays the same
Ex) Carbon – has 3 isotopes
1) Carbon – 12
2) Carbon – 13
3) Carbon – 14

25. Isotopes of Carbon All have the same # of p+ 1) Carbon – 12
If not, it would be a different element
All have 6 protons
1) Carbon – 12
Has 6 neutrons
2) Carbon – 13
Has 7 neutrons
3) Carbon – 14
Has 8 neutrons

26. Atomic Mass
Definition = weighted average mass of the atoms in a naturally occurring sample of the element
Units = amu
Equal to 1/12 the mass of a carbon-12 atom
This is the mass that is on the periodic table
Based on the mass and abundance of each isotope

27. Atomic Mass
Formula:
Atomic mass =
Repeats for however many isotopes exist for that element…. +
Abundance x mass
Abundance x mass
ISOTOPE #1
ISOTOPE #2

28. Atomic Mass – Sample Problem
Chlorine has 2 isotopes: chlorine-35 has an abundance of 75.77% and chlorine-37 which is24.33% abundant. What is the atomic mass of chlorine?
Atomic mass = (35*0.7577) + (37*0.2433)
Atomic mass = 35.5 amu
Abundance x mass +

29. Isotopes of Hydrogen Hydrogen-1: 1 p+ and 0 no
Relative abundance = %
Commonly called normal “hydrogen”
Hydrogen-2: 1 p+ and 1 no
Relative abundance = 0.015%
Commonly called heavy hydrogen or “deuterium”
Hydrogen-3: 1 p+ and 2 no
Relative abundance = ~0.00%
Commonly called “tritium”

30. IONS Atoms with a charge (results from gain/loss of electron)
If an atom GAINS an electron = ANION
Negative charge
Ex: Cl-
If an atom LOSES an electron = CATION
Positive charge
Ex: Ca2+

31. Practice with ions Examples: 15 O 2- p+: ___ no: ___ e-: ___
89 Sr p+: ___ no: ___ e-: ___
107 Ag + p+: ___ no: ___ e-: ___

32. Ions p+ n0 e- 6 14 11 12 1+ S 33 2- Examples: Symbol Atomic # Mass #
Charge 6 14 11 12 1+ S 33 2-

33. CH5 EOC Questions Answer the following in your composition book
Questions found on page 129
#33, 36, 38-43, 45-46, 48-49, 53, 57