1. Chemistry: Chapter 4
Atomic Structure

2. The Atom
An atom is the smallest particle of an element that retains the identity of that element.
If we repeatedly cut a piece of aluminum, the smallest possible piece is an atom of aluminum

3. Democritus (460 – 370 BCE) – Greek Philosopher
Proposed that all matter was made of tiny indivisible particles.
He called these particles atomos (meaning indivisible).
We call them atoms.
the “laughing philosopher”

4. Aristotle (384 – 322 BCE) – Greek Philosopher
Did not agree with Democritus.
Believed matter was made up of only one substance called “hyle” (greek for wood, or materials)
It wasn’t until the 1700’s when his ideas were reexamined.

5. Robert Boyle (1627 – 1691) Isaac Newton (1642 – 1727)
Published articles stating their belief in the atomic nature of elements
They had no proof
only 13 elements were known: antimony, arsenic, bismuth, carbon, copper, gold, iron, lead, mercury, silver, sulfur, tin, and zinc.
Robert Boyle ↑
← Isaac Newton

6. Antoine Lavoisier (1743 – 1794) – French Chemist
The “Father of Modern Chemistry”
Developed the “Law of Conservation of Matter.”
Matter is neither created nor destroyed.
Lavoisier published a list of elements, adding another 11: chlorine, cobalt, hydrogen, manganese, molybdenum, nickel, nitrogen, oxygen, phosphorus, platinum, and tungsten.
Pioneered Stoichiometry
Guillotined on May 8th in Paris, 1794

7. Joseph Proust (1754 – 1826) – French Chemist
Developed The Law of Definite Proportions
Compounds always contain elements in the same proportion by mass.

8. Law of Definite Proportions
H20 (by mass is always)
88.9% Oxygen, 11.1% Hydrogen
If we had an 80g sample of H20 how much is O?
.889 x 80 = 71g
How much is H?
.111 x 80 = 9g

9. John Dalton (1776 – 1844) – English Scientist
Proposed the atomic theory of matter, which is the basis for present atomic theory
John Dalton, English schoolteacher

10. Atomic Theory of Matter
Each element is composed of extremely small particles called atoms.
All atoms of a given element are identical, but differ from those of any other element.
Which element is this?

11. Atomic Theory of Matter
When elements unite to form compounds, they do so in a ratio of small whole numbers. This is called the Law of Multiple Proportions.
Ex: C and O can combine to form CO or CO2, but not CO1.8.

12. All matter is composed of tiny solid particles
Dalton’s Model of an Atom
All matter is composed of tiny solid particles

13. Joseph Louis Gay-Lussac (1778 – 1850) – French Chemist
Observed that working with gas reactions at constant volume: temperature and pressure are directly related.
He named the discovery of this relationship Charles Law, which is represented by:
P 1 T 1 = P 2 T 2

14. Amadeo Avogadro (1776 – 1856) – Italian Physicist
Explained Gay-Lussac’s work using Dalton’s theory: Equal volumes of gases at the same temperature and pressure have the same number of gas molecules.

15. Michael Faraday (1791 – 1867) – English Physicist
Suggested that atomic structure was related to electricity.
Atoms contain particles that have electrical charges.
Positive (+)
Negative (-)
Opposite charges attract
Like charges repel

16. William Crookes (1832 – 1919) – English Physicist
Developed the cathode ray tube to find evidence for the existence of particles within the atom.

17. J.J. Thomson (1856 – 1940) – English Physicist
Used a cathode ray tube to identify negatively charged particles, called electrons.
Determined the ratio of an electron’s charge to its mass.
Developed the “plum pudding” model of an atom.

18. Cathode Ray Tube Experiment

19. Plum Pudding Model
Atoms are composed of a solid sphere with charged particles embedded throughout

20. Robert Millikan(1868 – 1953) – American Physicist
US Physicist
Used the oil drop experiment to prove the electron has a negative charge
Was able to determine the charge of the electron

21. Millikan’s Oil Drop Experiment

22. Ernest Rutherford (1871 – 1937) – New Zealand Physicist
Used the gold foil experiment to prove the atom is mostly empty space.
Rutherford concluded that all of an atom’s positive charge, and most of its mass is located in the center, called the nucleus.

23. Gold Foil Experiment

24. Gold Foil Experiment 98% of the particles passed straight through
2% of the particles deflected off at varying angles
0.01% of the particles bounced straight back

25. Planetary Model of an atom
Rutherford’s
Planetary Model of an atom
Positive charge and majority of mass located in the nucleus.
Negatively charged electrons orbit the nucleus like planets.

26. James Chadwick (1891 – 1974) – English Physicist
Found high energy particles with no charge with the same mass as the proton called neutrons.

27. Max Planck (1858 – 1947) – German Physicist
Planck demonstrated that the electron in orbit around the nucleus accelerates because it changes direction.
Because the electron has charge, acceleration means a changing electric field, which means photons should be emitted.
But, then the electron would lose energy and fall into the nucleus.
Niels Bohr (left) and Max Planck

28. Planck’s Dilemma – Why don’t the electrons fall into the nucleus?
Planck determined that energy, at the sub-atomic level, can only be transferred in small units, called quanta.
Quanta are not divisible.
There are no “in between” energy levels.
Quantization limits the energy to be transferred to photons and resolves the dilemma of the electron falling into the nucleus.

29. Components of the Atom
The smallest particle of an element that has the properties of that element.
Make up of nucleus consists of protons and neutrons
Surrounded by an electron cloud

30. Protons (p+) Found in the nucleus Positive (+) charge
Composed of 3 quarks
(2 up, 1 down)
Mass= x kg
Atomic mass  1 amu (µ)
The number of protons in an atom refers to the atomic number (Z)

31. Neutrons (n0) Found in the nucleus Neutral (0) charge
Composed of 3 quarks
(1 up, 2 down)
Mass= x kg
Atomic mass  1 amu (µ)
Isotopes: atoms of the same element that have a different number of neutrons.

32. Electrons (e–) Found in electron clouds surrounding the nucleus.
Negative (–) charge
Elementary Particle
Mass = x kg
1800 times smaller than protons & neutrons
Mass  0 amu (µ)

33. Electrons (e–)
Have negligible mass (when compared to protons and neutrons)
Orbit the nucleus at very high speed in energy levels (electron clouds).

34. Atomic Number = Protons
The atomic number of an element is the number of protons an element has.
Located above the symbol of the element
The number of protons determines the identity of the element.
Each element has a different atomic number

35. How many electrons are in an atom?
For an atom to have an overall neutral charge the number of electrons must equal the number of protons.
p+ = e–
What element is this?

36. Mass number Mass = Protons + Neutrons
The Mass number of an atom is the average of the mass of the isotopes of that element
Located below the symbol of the element
Atomic mass is measured in amu’s, (atomic mass units)
Based on Carbon having a mass of 12
Mass = Protons + Neutrons

37. How many neutrons are in this isotope of platinum?
mass = p+ + n0
196 = 78 + n0
n0 = 196 – 78
Platinum has 118 Neutrons
Find the number of neutrons in:
Hydrogen - 1 Carbon - 13
Helium Potassium - 40
Boron - 11 Gold - 198
196

38. mass = p+ + n0 Hydrogen–1 ( 1 1 H) hydrogen–1 has 0 neutrons
Helium–4 ( 4 2 He) helium–4 has 2 neutrons
Boron–11 ( 11 5 B) boron–11 has 6 neutrons
Carbon–13 ( 13 6 C) carbon–13 has 7 neutrons
Potassium–40 ( K) potassium–40 has 21 neutrons
Gold–198 ( Au) gold–198 has 119 neutrons

39. Average Atomic Mass Number
The average mass of all of the isotopes of an element.
Also known as the average atomic mass number, or atomic weight.
Isotopes: atoms of the same element with different masses.

40. Average Atomic Mass 20 10 Ne has a mass of 19.992 amu
90.00% of a sample of neon is neon–20
10.00% of a sample of neon is neon–22
Calculate the average atomic mass of neon
.90 x =
.10 x =
average mass = amu

41. Ions An atom that has gained or lost an electron.
It acquires a net electrical charge.
If an atom loses an electron (oxidation) it has more protons than electrons and has a net positive charge. (cation)

42. Ions
If an atom gains an electron (reduction) it has more electrons than protons and has a net negative charge (anion)

43. Ionic Charges Charge of ion = p+ – e–
What is the charge of a magnesium atom that loses 2 electrons?
Number of p+ 12
Number of e– 10
charge of ion +2
Mg2+ or Mg+2
Charge is written to the upper right of the symbol.

44. Representations of atoms
General form: (Elemental Notation)
X = Element Symbol
A = Atomic Mass (p+ + n0)
Z = Atomic Number (p+)
Ionic Charge
A Z X charge

45. What is the atomic structure?
Determine the number of:
p+ =
n0 =
e– =
Na +

46. What is the atomic structure?
Determine the number of:
p+ = 11
n0 = 12
e– = 10
Na +

47. What is the atomic structure?
Determine the number of:
p+ =
n0 =
e– =
I −

48. 127 53 I − p+ = 53 n0 = 74 e– = 54 What is the atomic structure?
Determine the number of:
p+ = 53
n0 = 74
e– = 54
I −

49. Put into elemental notation?
Atomic # = 29
Atomic Mass = 64
Ionic charge = +2

50. 64 29 Cu 2+ How many electrons? Atomic # = 29 Atomic Mass = 64
Ionic charge = +2
Number of e– =
Cu 2+

51. Put into elemental notation?
37 p+
48 n0
36 e–

52. Put into elemental notation
Rb +

53. Alpha Decay (α) 2 protons and 2 neutrons bound together
emitted from the nucleus during radioactive decay
Occurs when nucleus is unstable
4 2 He2+

54. Beta Decay (β) An electron is emitted from the nucleus
This occurs when the neutron to proton ratio is too large
A neutron decays into a proton, an electron, and a antineutrino

55. Gamma Decay (γ)
High energy electromagnetic waves (no mass) emitted from the nucleus as it loses energy
Very penetrating form of radiation
High energy radiation, stopped by lead plates.