Chapter 3Atoms: The Building Blocks of

Matter

3.1 The Atom: From Philosophical Idea to Scientific

Theory– 4th century B.C. Greek

philosopher Democritus stated the universe was made of invisible units called atoms (atom “unable to be divided”).

– Believed movements in atoms caused changes observed in matter. What held these atoms together? Democritus did not know…. And no one did until…

• Late 1700’s– French Chemist, Antoine Lavoisier

established the Law of Conservation of Mass.Mass can NOT be created or destroyed; only change forms.

– Joseph Proust, later established the Law of definite proportionsthe compound always contains the same elements in the same proportions by mass.

• Law of multiple proportions- if 2 or more different compounds are composed of the same 2 elements, then the ratios of the masses of the 2nd element combined with a certain mass of the 1st element is always a ratio of small whole #s.

Dalton’s Atomic Theory– In 1803 John Dalton proposed an

atomic theory:• Each element is composed of

extremely small particles called atoms .

• Atoms of the same element are exactly alike .

• Every compound always has the same ratio and kinds of atoms .

• A chemical rxn is a rearrangement of atoms; they are not created or destroyed .

• Remember: an atom is the smallest part of an element that still has the element’s properties.

Today…• There are some exceptions to Dalton’s

theory; however they still are the basis for understanding Chemistry. – 1. all matter is composed of atoms – 2. atoms of any one element differ in properties

from atoms of another element• BUT… how can we be sure they exist?!?

3.2 The Structure of the Atoms• Atom- “ the smallest particle of an

element that retains the chemical properties of that element.”

The Structure of the Atom• We now know that atoms can be divided into

many different subatomic particles. For example:– Nucleus- the center of the atom– Protons- positively (+) charged subatomic particle.– Neutrons- neutral (not charged) subatomic

particle.– Electrons- negatively (-) charged subatomic

particle

Discovery of the Electron• The English Chemist, Michael

Faraday, in 1839, showed atoms contain electrical charge.

• American, Benjamin Franklin, experimented with electricity (think kite and key experiment). Franklin was able to determine the 2 charges an object has and it was he who coined the names “positive” and “negative” charge. Opposite charges attract and like charges repel. • But where do these charges

come from? What are they?

The Discovery of the Electron

Cathode Rays and Electrons• Cathode Ray Tube (CRT) - A battery is

connected to a tubing of partially evacuated glass. The glass is lined with fluorescent material, current flows to the ends of the tube. The end connected to the (-) terminal of the battery is called the cathode and the other is the anode (+). A stream of radiation flows from the cathode to the anode when the battery is turned on.

http://www.youtube.com/watch?v=XU8nMKkzbT8

• English physicist, J.J. Thompson (1856-1940), determined the ray WAS made (-) particles by allowing the ray to pass through a hole in the anode and then through a magnetic field. He determined that the negative particles emanate from the cathode, they had structure and he named them electrons.

Charge and Mass of the Electron

• American physicist, Robert Millikan (1868-1953), was able to measure the charge of an electron. He sprayed oil and used X-rays to give the oil a negative charged. Then measured how different magnetic charges changed the rate the oil fell.

He calculated the mass of the e- to be 9.11 X 10-19 grams.

Discovery of the Atomic Nucleus

• New Zealander, Ernest Rutherford passed the cathode ray of a radioactive substance between two charged plates. The ray split- part of the beam was deflected towards the (-) plate (alpha radiation), another was deflected towards the (+) plate (beta radiation) and a third passed straight through undisturbed (gamma radiation).

• Alpha particles 2+ charge• Beta particles 1- charge• Gamma rays have no charge• This experiment demonstrated that the atom

was much more complex than previously thought.

Composition of the Atomic Nucleus

• Thompson showed us that atoms had electrons, but that doesn’t explain why atoms are electrically neutral (they don’t have a charge). If the have electrons they must have some type of (+) charges, too.

• Thompson had suggested, earlier, that the atom was like plum pudding. The e-s were spaced evenly throughout the atoms (+) interior, the way the plums were distributed through the pudding. (You can also think of chocolate chip cookie dough- the dough is the positive interior while the chocolate chips are the e-s).

Alpha- Scattering Experiment• 1909 The Alpha-scattering Experiment by

Rutherford- A beam of high-speed alpha particles bombarded a thin sheet of Au foil. Most of the alpha particles went straight through the foil but a small portion DID deflect. And they would scatter in every direction possible. Why??

http://www.youtube.com/watch?v=XBqHkraf8iE

• Rutherford’s data suggested the “plum pudding model” was not true. He determined that there must be an (+) core in the atom. He named this the nucleus. The particles which went straight through suggested the atom was mostly empty space.

3.3 Counting Atoms• Atomic Numbers• Englishman, Henry Mosley (1887-1915)

discovered that each element had a unique amount of positive charge. This concept leads to the understanding of why atoms of different elements are unique. The identity of the atom comes from the number of protons contained in the nucleus.

• Atomic # (Z)- the # of protons in the nucleus of an atom. (Always a whole #)– Ex. O-8 = 8 protons in the nucleus– C-6 = 6 protons in the nucleus– Pb-82= 82 protons in the nucleus

• Now, individual atoms are electrically neutral… meaning they have an equal # of protons as electrons. – O has 8 protons (we know this by its atomic number) and 8 electrons. – C has 6 protons and 6 electrons– Pb has 82 protons and 82 electrons.

Ions: When p+ and e- are not equal.

• The atom can gain or lose electrons (NOT PROTONS- if it change

the # of protons you change the element/atom.) This is an ion. If you gain an electron you get a negative ion. If you lose an electron you have a positive ion.

• Charge on Ion = # of protons - # of electrons– Magnesium (Mg) can loose 2 electrons making it a positive ion- written

as:Mg2+ . O will gain 2 electrons making it a negative ion. It is written as such: O2- .

• Atoms that lose electrons, like Mg, are called cations. Atoms that gain electrons, like O, are called anions.

Isotopes• Just as we have the atomic # (# of protons) we

have atomic mass #:Atomic mass # (A) - the # of protons + neutrons in

an atom.

• Because the bulk of the atom’s mass is provided by the protons and neutrons, we only consider their masses when calculating the atomic mass #.– Ex: O has 8 protons and 8 neutrons, so A= 16.

• Dalton said that every atom of the same element is exactly alike… NOT SO. They do have the same # of protons; however, they do not necessarily have the same # of neutrons!

• This means that different atoms of the same element may/will have different masses.

• These are: isotopes (atoms having the same # of protons

but different # of neutrons).

– Some isotopes are more common than others. H is present on the earth and the sun. The most common isotope of H is protium, then deuterium and finally, tritium.

– Ex: there are 3 isotopes of H (why A = 1.00794).

• Protium (only has one proton in the nucleus) A=1• Deuterium (1 proton + 1 neutron) A=2• Tritium (1 proton + 2 neutrons) A=3

• Calculating the # of neutrons in an atom:General Formula:A – Z= # of neutrons.– Ex: calculate the # of neutrons in the isotopes of Cl (Cl-35 and Cl-37)

Relative Atomic Masses• The mass of a single element is extremely small

(in the one trillionth of a billionth range) and is very difficult to work with. So instead we express the mass of atoms in atomic mass units (amu). – One amu is equal to 1/12th of the mass of a carbon-

12 atom. This isotope has exactly 6 protons and 6 neutrons so the mass of each has to be about 1.0 amu.

• 1amu = 1/12 (mass of 126C atom) = 1.66x10-24 g

Average Atomic Masses of Elements

• The atomic mass of an element is often listed as the average atomic mass as found in nature. This is a weighted average of the isotopes for that particular element. The more commonly found isotopes have a greater effect on the averages mass than the more rare isotopes. – Table 3-4 on p.80 in your book

– Ex: Cl 24% Cl-37 and 76% Cl-35, thus the average atomic mass (35.45 amu) is much closer to 35 than 37.

Relating Mass to Numbers of Atoms

• If we can have an atomic mass measured in amu…. we can also have FORMULA MASS:

• Formula Mass- the sum of the atomic masses of all the atoms in the compound.

• Ex. CO2 C- 12.01O- 16.00 X 2

44.01 amu – formula massYOU TRY!1. H2O2. CH4Cl2

So… What’s a MOLE?!?• We have a problem… how do you

measure amu’s? How do you count atoms? They are so small… that you don’t. Instead, we use a number called the mole to help us.

• A mole (of any element) – the number of atoms equal to the number of atoms in exactly 12.0 g of carbon-12.

How does it work?• There is an actual # … it is 6.02 x 1023

atoms/mole. – If you have 1.0 g of H, you have 6.02 x 1023

atoms of H– If you have 16.0 g O, you have 6.02 x 1023 atoms

of O

How many grams of Pb would you need in order to have 6.02 x 1023 atoms of Pb?

Notice the amu- gram link?

• “The mass in grams of 1 mole of a substance is numerically equal to its atomic mass or formula mass in atomic mass units.”

• “Mole” is just the name of the number we use in chemistry to help count atoms.– Like…

• A couple-• A dozen –• A baker’s dozen-• A year-

Avogadro’s Number• The Mole is also known as Avogadro’s Number (N) in honor of Amadeo Avogadro, the Italian chemist.

How Big is Avogadro’s number? 602,000,000,000,000,000,000,000!

if it was a stack of paper… it would reach beyond our solar system!…basketballs… it could create a planet the size of the earth!…grains of rice… would cover the earth to a depth of 246 feet!

Molar Mass• “The mass in grams of 1 mole of a substance..” (M).• Molar Mass of…

– Ca 40 amu –> 40 g/mol– C 12 amu --> 12 g/mol– H2O2 34 amu --> 34 g/mol– NaCl– CaCl2

– C6H12O6

Gram-mole-Particle Conversions

• We can use the molar mass to convert moles to grams or grams to moles.

• Likewise, moles can be converted to # of particles via Avogadro’s #.