1. Atoms: the building blocks of matter
Chapter 3
Chemistry chapter 3

2. The atom
The atom – smallest piece of matter that has the properties of an element.
Made of
Protons
Neutrons
Electrons
Each specimen of a specific subatomic particle is the same
 If we split an atom, we no longer have a specific element
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3. Early atomic theory - Democritus
Greek philosopher about 400 B.C.
Gave us the word atom
Atomos - indivisible.
Thought
The world was made of empty space and particles called atoms.
There were different types of atoms for different types of materials.
Theory was not supported by experimental evidence.
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4. Early atomic theory – Aristotle
Aristotle did not believe in atoms
thought matter was continuous
He was very influential, so Democritus’s theory was not accepted for many centuries.
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5. 17th century People began to express doubts in Aristotle’s theory.
Isaac Newton and Robert Boyle published articles stating their belief in the atomic nature of elements, but they had no proof.
Their theory also had no ability to predict the unknown.
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6. Antoine Lavoisier – late 1700s
Law of conservation of mass
during a chemical change in a closed system, no mass is lost
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7. Joseph Proust – late 1700s to early 1800s
Law of definite proportions
specific substances always contain elements in the same ratio by mass
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8. Law of multiple proportions
Some elements form more than one compound with each other.
If two or more different compounds are composed of the same two elements, then the ratio of their masses always contains small whole numbers
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9. John Dalton – early 1800s
Studied experimental observations of chemical reactions
Proposed explanation of these three laws
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10. Dalton’s Hypothesis
All matter is composed of very small particles called atoms.
All atoms of an element are exactly alike; atoms of different elements are very different.
Atoms cannot be subdivided, created, or destroyed.
Atoms unite with other atoms in simple ratios to form compounds
In chemical reactions, atoms are combined, separated, or rearranged.
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11. Did Dalton’s theory work?
Conservation of mass
the atoms are simply rearranged because they cannot be created or destroyed
Laws of definite and multiple proportions
Only whole atoms can combine, giving small whole numbers in ratios
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12. Gas research J.L. Gay-Lussac
Under constant temperature and pressure
Volumes of reacting gases and gaseous products are in a ratio of small whole numbers.
Amadeo Avogadro explained Gay-Lussac’s work with Dalton’s theory.
Equal volumes of gases, under the same temperature and pressure, have the same number of molecules.
Helped Dalton’s theory get accepted
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13. Dalton’s theories
Atomic theory and law of multiple proportions have been tested and accepted as correct.
However, there some major exceptions to the rules.
Splitting atoms
Different atoms of the same element
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14. Discussion
Section review on page 69
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15. Cathode tubes Anode – positive electrode Cathode – negative electrode
When the tube is on, cathode rays appear that begin at the cathode and travel to the anode.
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16. Cathode rays and electrons
1897 – J.J Thomson tested cathode rays and discovered that they were electrons.
Rays turned a paddlewheel – they had mass
Rays deflected by a magnet just like current-carrying wire – they were negatively charged
He determined the ratio of the electron’s charge to its mass.
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17. Charge on an electron Robert Millikan’s famous oil drop experiment.
Tiny oil drops fell through a chamber
gravitational force offset by applying an opposing electrical force.
Charge on oil drops determined
This charge was always a whole number multiple of one small charge
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18. Charge on an electron
This small charge was the charge on one electron.
This is now the standard unit of negative charge (1-). It can be written e-.
e- can also represent an electron
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19. Mass of an electron
Using Thomson’s ratio and Millikan’s charge, determined to be 9.1 x kg
It was found that it’s mass is only 1/1837 the mass of the lightest atom known – the hydrogen atom.
Most of the mass must be somewhere else
Since atoms are neutral, there must be some positive charge
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20. Thomson’s plum pudding model
In this model, the raisins were the electrons and the pudding was the positive charge.
Sort of like chocolate chip cookie dough.
The chips are the electrons and the dough is the positive charge.
Explained the experiments that had been done so far.
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21. Testing the plum pudding model
See page 72
fired alpha particles at a very thin (a few atoms thick) sheet of gold foil.
They expected the particles to go right through because the spread out positive charge in the “pudding” wouldn’t be strong enough to deflect them.
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22. What happened
Most of the particles did go right through without being deflected at all.
Some were deflected at large angles.
Ernest Rutherford explained it:
the positive charge on the atom was concentrated at a small core – now called the nucleus.
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23. The atom as we now “know” it
The nucleus contains all of the positive charge and most of the mass.
The negatively charged electrons have very small mass and are located around the nucleus in the electron cloud.
Most of an atom is empty space.
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24. Protons same charge as an electron; opposite sign.
standard unit of positive charge (1+)
Much larger mass than the electron: x kg
The number of protons determines the atom’s identity.
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25. Neutrons Weren’t discovered until the 1930s.
Neutral – no charge – harder to detect
Slightly more mass than a proton: x kg
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26. Nuclear or Strong Force
The force that holds protons and neutrons together.
It is effective only for very short distances – about m.
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27. Dalton’s theory Dalton thought that atoms were indivisible
discovery of electrons, protons, and neutrons did not fit with his theory.
Led to major revisions in atomic theory
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28. Isotopes Thomson discovered what seemed to be two kinds of neon atoms.
Same chemical properties; different masses.
Atoms of the same element that differ in mass are called isotopes.
Have the same number of electrons and protons but different number of neutrons.
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29. Atomic number Number of protons in an atom
Atoms are electrically neutral,  the number of electrons must equal the number of protons.
The number of protons determines the identity of the atom and the number of neutrons determines the isotope.
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30. Modification of Dalton’s theory
All atoms of an element contain the same number of protons but can contain different numbers of neutrons.
So we have to use average mass of an atom.
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31. Nucleons Particles in the nucleus – protons and neutrons
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32. Mass number Total number of nucleons : protons plus neutrons
Number of neutrons = mass number minus atomic number
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33. Designating Isotopes Hyphen notation
Uranium-235
Carbon-14
Carbon-12
The number refers to the mass number
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34. Nuclide General term for any isotope of any element
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35. Atomic mass units
There must be a standard for all units of measurement.
A Carbon-12 atom with 6 protons and 6 neutrons was chosen as the standard
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36. Atomic mass unit Defined as 1/12 the mass of that carbon atom.
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37. Average atomic masses
Many elements have an average atomic mass close to the number of nucleons in their nuclei – near whole numbers.
Some don’t – look at Chlorine
The periodic table shows average atomic masses.
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38. Weighted averages
We then use a weighted average to find the average mass of an atom of a given element.
This is called the average atomic mass or just atomic mass.
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39. Finding a weighted average
A class of 25 students took a test. 10 of them got 80%. 12 got 90%. 3 got 100%. What was the average score?
Not 90% - probably less than that.
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40. You try
Neon has two isotopes. Neon-20 has a mass of amu and neon-22 has a mass of amu. In any sample of 100 neon atoms, 90 will be neon-20 and 10 will be neon-22. Calculate the average atomic mass of neon.
amu
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41. You try
Compute the average atomic mass of silver, if 51.83% of the silver atoms occurring in nature have mass amu and 48.17% of the atoms have mass amu.
107.9 amu
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42. The Mole SI unit for amount of substance Abbreviated mol
A counting unit
6.022 x 1023 particles
Avogadro’s number
Based on carbon-12, 12 g of C-12 contains 1 mol of atoms
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43. Molar mass The mass of 1 mol of a pure substance g/mol
Numerically equal to the atomic mass in amu
On the periodic table the number with a decimal is the atomic mass in amu AND the molar mass in g/mol
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44. conversions Grams to moles or moles to grams Use the molar mass
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45. Example What is the mass in grams of 5.60 mol of sulfur?
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46. Example How many moles of carbon are in a sample with a mass of 567 g?
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47. Example
How many atoms of lithium are in a sample with a mass of 76.2 g?
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48. You try
How many moles of rubidium are in x 1023 atoms of rubidium?
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49. You try
How many moles are in g of zinc?
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50. You try What is the mass of 1.20 x 1025 atoms of helium?
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