1. Chapter 0 A Very Brief History of Chemistry
Chemistry: The Molecular Nature of Matter, 7E
Jespersen/Hyslop

2. Chapter in Context
Scope and purpose of the chemical sciences – Chemistry’s big ideas
Formation of the elements – supernovas
Distribution of substances around the world – elements and the Earth
Atomic theory – explanatory and predictive power
The structure of the atom – key experiments

3. The Four Big Ideas Atomic theory – John Dalton, 1813
Described atoms and how they interact with one another
2. Careful laboratory observation
Can lead to understanding the atomic world
3. Energy changes and probability
Lead to predictions about chemical interactions
Geometric shapes of molecules are important Affect properties, reactivity, and function (e.g. DNA, RNA, and proteins)

4. Supernovas and the Elements
The Big Bang:
14 billion years ago
Explosion of energy and subatomic particles
Extreme temperature, pressure, and density
As the Big Bang cooled
Initially only quarks exist
After 1 second, quarks form protons and neutrons
After 3 minutes, nucleosynthesis begins of light nuclei (e.g., helium, lithium)
After further cooling, electrons join nuclei to form atoms.

5. Supernovas and the Elements
Universe was 91% hydrogen, 8% helium, 1% other light atoms
Uneven distribution of matter resulted in star formation
Formation of elements occurred in the stars
Small atoms combined, due to high pressure at the center, to create slightly heavier elements
New elements concentrated in the stars’ centers
Heavier elements then combined into new, even heavier elements
Cycle kept repeating

6. Supernovas and the Elements
Iron is the heaviest element created in stars
Causes the nuclear reactions to stop and the star to cool and collapse in on itself
allows for even heavier elements to form
Eventually, the star disintegrates
Called a supernova
Spews its content into space
Remnants rejoin to form new stars
Cycle begins again
Some of the debris combines to form moons, planets, and asteroids

7. Distribution of the Elements
Earth formed 4.5 billion years ago
Result of gravitational forces
Earth heated up
Iron and nickel melted
Migrated to the core
Outer core is superheated lava
Mantel is superheated rock
Crust is the surface
10 miles thick
Contains the familiar elements (gold, silicon, carbon, etc.)

8. Atomic Theory Most significant theoretical model of nature Atoms
Tiny submicroscopic particles
Make up all chemical substances
Make up everything in macroscopic world
Smallest particle that has all properties of given element

9. Atomic Theory Three important ideas Law of Definite Proportions
In a given compound, the elements are always combined in the same proportion by mass
Always find 1 g H to 8 g O in water
Law of Conservation of Mass
No detectable gain or loss of mass occurs in chemical reactions. Mass is conserved.
A closed vessel with 16 g O and 2 g H will weigh 18 g after water is formed from them
Dalton’s atomic theory

10. Atomic Theory
Law of definite proportions and law of conservation of mass
Based on laboratory observations of mass and volume
Discussed in detail in a later chapter

11. Dalton’s Atomic Theory
John Dalton
Developed underlying theory to explain
Law of Conservation of Mass
Law of Definite Proportions
Reasoned that if atoms exist, they have certain properties

12. Dalton’s Atomic Theory (cont.)
Matter consists of tiny particles called atoms
Atoms are indestructible
In chemical reactions, atoms rearrange but do not break apart
In any sample of a pure element, all atoms are identical in mass and other properties
Atoms of different elements differ in mass and other properties
In a given compound, constituent atoms are always present in same fixed numerical ratio

13. Proof of Atoms Early 1980’s, use Scanning Tunneling Microscope (STM)
Surface can be scanned for topographical information
Image for all matter shows spherical regions of matter
Proof of atoms
STM of palladium

14. Discovery of Subatomic Particles
Late 1800s and early 1900s
Cathode ray tube experiments showed that atoms are made up of subatomic particles
Discovered negatively charged particles moving from the cathode to the anode
Cathode – negative electrode
Anode – positive electrode

15. Discovery of Electron JJ Thomson (1897) Modified cathode ray tube
Made quantitative measurements on cathode rays
Discovered negatively charged particles
Electrons (e –)
Determined charge to mass ratio (e/m) of these particles
e/m = –1.76 x 108 coulombs/gram

16. Millikan Oil Drop Experiment
Determining charge on Electron
Calculated charge on electron
e – = –1.60 × 10–19 Coulombs
Combined with Thomson’s experiment to get mass of electron
m = 9.09 × 10–28 g

17. Discovery of Atomic Nucleus
Rutherford’s Alpha Scattering Experiment
Most alpha () rays passed right through gold
A few were deflected off at an angle
1 in 8000 bounced back towards alpha ray source
Gave us current model of nuclear atom

18. Discovery of Proton Discovered in 1918 in Ernest Rutherford’s lab
Detected using a mass spectrometer
Hydrogen had mass 1800 times the electron mass
Masses of other gases whole number multiples of mass of hydrogen
Proton
Smallest positively charged particle

19. Rutherford’s Nuclear Atom
Demonstrated that nucleus:
has almost all of mass in atom
has all of positive charge
is located in very small volume at center of atom
Very tiny, extremely dense core of atom
Where protons ( ) and neutrons ( ) are located

20. Discovery of Neutron First postulated by Rutherford and coworkers
Estimated number of positive charges on nucleus based on experimental data
Nuclear mass based on this number of protons always far short of actual mass
About ½ actual mass
Therefore, must be another type of particle
Has mass about same as proton
Electrically neutral
Discovered in 1932 by James Chadwick

21. Your Turn!
Which scientist is correctly matched with the discovery for which he is known?
Thomson, neutron
Rutherford, electron
Chadwick, neutron
Milliken, nucleus
Dalton, charge on the electron

22. Atomic Structure
Electrons ( , or e –)
Very low mass
Occupy most of atom’s space
Balance of attractive and repulsive forces controls atom size
Attraction between protons ( ) and
electrons ( ) holds electrons around nucleus
Repulsion between electrons helps them spread out over volume of atom
In neutral atom
Number of electrons must equal number of protons
Diameter of atom ~10,000 × diameter of nucleus

23. Properties of Subatomic Particles
Three kinds of subatomic particles of principal interest to chemists
Particle
Mass (g)
Electrical Charge
Symbol
Electron
9.109  10–28 –1
Proton
1.673  10–24 +1
Neutron
1.675  10–24

24. Atomic Notation Atomic number (Z) Isotopes
Number of protons that atom has in nucleus
Unique to each type of element
Element is substance whose atoms all contain identical number of protons
Z = number of protons
Isotopes
Atoms of same element with different masses
Same number of protons ( )
Different number of neutrons ( )

25. Atomic Notation Isotope Mass number (A) Atomic Symbols
A = (number of protons) + (number of neutrons)
A = Z + N
For charge neutrality, number of electrons and protons must be equal
Atomic Symbols
Summarize information about subatomic particles
Every isotope defined by two numbers Z and A
Symbolized by
Ex. What is the atomic symbol for helium?
He has 2 e–, 2 n and 2 p Z = 2, A = 4

26. Isotopes Most elements are mixtures of two or more stable isotopes
Each isotope has slightly different mass
Chemically, isotopes have virtually identical chemical properties
Relative proportions of different isotopes are essentially constant
Isotopes distinguished by mass number (A):
e.g.,
Three isotopes of hydrogen (H)
Four isotopes of iron (Fe)

27. Example: What is the isotopic symbol for Uranium- 235?
Number of protons ( ) = = number of electrons in neutral atom
Number of neutrons ( ) = 143
Atomic number (Z ) = 92
Mass number (A) = = 235
Chemical symbol = U
Summary for uranium-235:

28. Your Turn!
An atom of has ___ protons, ___ neutrons, and ___ electrons.
82, 206, 124
124, 206, 124
124, 124, 124
82, 124, 82
82, 124, 124

29. Your Turn!
What is the correct symbol for an element that has 27 protons, 33 neutrons, and 27 electrons.

30. Learning Check: 78 53 53 46 35 35 Fill in the blanks: 131I 81Br
symbol neutrons protons electrons
131I
81Br 78 53 53 46 35 35

31. Carbon-12 Atomic Mass Scale
Need uniform mass scale for atoms
Atomic mass units (symbol u)
Based on carbon:
1 atom of carbon-12 = 12 u (exactly)
1 u = 1/12 mass 1 atom of carbon-12 (exactly)
Why was 12C selected?
Common
Most abundant isotope of carbon
All atomic masses of all other elements ~ whole numbers
Lightest element, H, has mass ~1 u

32. Calculating Atomic Mass
Generally, elements are mixtures of isotopes
e.g. Hydrogen
Isotope Mass % Abundance
1H u
2H u
How do we define atomic mass?
Average of masses of all stable isotopes of given element
How do we calculate average atomic mass?
Weighted average
Use isotopic abundances and isotopic masses

33. Learning Check
Naturally occurring magnesium is a mixture of 3 isotopes; 78.99% of the atoms are 24Mg (atomic mass, u), 10.00% of 25Mg (atomic mass, u), and 11.01% of 26Mg (atomic mass, u). From these data calculate the average atomic mass of magnesium.
x u = u Mg
x u = u 25Mg
x u = u 26Mg
Total mass of average atom = u rounds up to u

34. Your Turn!
A naturally occurring element consists of two isotopes. The data on the isotopes:
isotope # u %
isotope # u %
Calculate the average atomic mass of this element. u u u u u
× u = u
× u = u u

35. Your Turn!
Magnesium (average atomic mass u) consists of three isotopes.
isotope # u %
isotope # u %
isotope #3 _______ _______
Calculate the abundance and atomic mass of #3. u u u u u
Abundance: – –
= % %
u = ( × u) ( × u) +
( × X u)
Solving for X yields u u