Atoms: The Building Blocks of Matter

CHAPTER 3

Origins of the Atom• Democritus: Greek philosopher

(460 BC - 370 BC)

• Coined the term atom from the Greek word “atomos”

• Democritus believes that atoms were indivisible and indestructible.

Let’s Get Ready to Rumble• The idea of the atom was met

with great skepticism, especially among great thinkers.

• The most vocal critic of Democritus’s idea was Aristotle.

• Aristotle’s chief argument was that there was no proof of the existence of atoms. Democritus’s claim was based purely on philosophical argument.

Aristotle’s Theory of the Atom

• Aristotle’s theory centers around the idea that everything is made up of only 4 elements: earth, wind, fire, and water.

• THERE ARE NO INDIVISIBLE PARTICLES!

Bringing Atoms Back!• During the 17th century,

scientists quietly revive the idea of the atom.

• One of the chief supporters include Sir Isaac Newton.

• At the time, most scientists are concerned with trying to explain the properties of gases.

Detour: The Three Basic Laws of Chemistry

• In their quest to discover proof of atoms, scientists began to propose three basic laws that explain all of the behavior in chemistry.

• Law of Conservation of Mass (or Matter)

• Law of Definite Proportions (or Constant Composition)

• Law of Multiple Proportions

Law of Conservation of Mass (or Matter)

• 1789: French chemist Antoine Lavoisier discovers that during an experiment involving red mercury oxide that the mass of the oxide before heating was equal to the mass of the newly formed mercury metal and oxygen gas.

• Matter cannot be created or destroyed.

Law of Definite Proportions (or Constant Composition)

• 1797-1804: French chemist Joseph Proust proposes the Law of Definite Proportions based on results from experiments using copper carbonate.

• Proust finds that all samples of copper carbonate had the same fixed composition.

• A chemical compound contains the same elements in the same proportions regardless of sample source or size.

Law of Multiple Proportions• 1803: John Dalton creates an explanation for the

Law of Conservation of Mass and the Law of Definite Proportions.

• As a result, Dalton creates the Law of Multiple Proportions.

• If 2 or more different compounds are made up of the same elements, then the ratio of the masses of elements is always a small, whole number.

Johnny D and the AT• 1808: English schoolteacher John Dalton proposes

his explanation of 2 of the 3 basic laws of chemistry.

• In his explanation, Dalton proposes proof of the atoms existence.

• Dalton’s Atomic Theory has 5 main points.

• You will have a quiz over Dalton’s Atomic Theory on Thursday!

Dalton’s Atomic Theory• 1. All matter is composed of

atoms.

• 2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.

• 3. Atoms cannot be subdivided, created, or destroyed.

Dalton’s Atomic Theory

• 4. Atoms of different elements combine in simple, whole-number ratios to form chemical compounds.

• 5. In chemical reactions, atoms are combined, separated, or rearranged.

Does Dalton’s Theory Still Hold?

• Not all portions of Dalton’s Atomic Theory are still vallid.

• Now we know that atoms can be subdivided into smaller subatomic particles such as electrons, protons, and neutrons.

• We also know that a given element can have atoms with different masses.

ATOMIC STRUCTURE

Discovery of the Electron• 1897: J.J. Thomson uses

a cathode ray tube to deduce the presence of a negatively charged particle: the electron.

How Heavy Is an Electron?• 1916: American scientist

Robert Millikan determines the mass of an electron to be 1/1837 the mass of a hydrogen atom.

• In addition, Millikan discovers that an electron has a negative one unit charge.

Is There Anything Else…• Many conclusions were made after the discovery of

the electron.

• 1. All elements must contain identically charged electrons.

• 2. There has to be positively charged particles if the atom is neutral.

• 3. There have to be heavier particles since electrons have very little mass.

On Another Note• 1886: German physicist

Eugen Goldstein discovers the proton, a positively charged particle, using anode rays.

• 1932: English scientist James Chadwick discovers the neutron, a particle with no charge but with a mass slightly larger than a proton.

Thomson’s Atomic Model• Thomson believed that the

electrons were like plums embedded in a positively charged pudding, thus calling his model the Plum Pudding model.

• This model has also been called the Blueberry Muffin Model, Chocolate Chip Cookie Model, or the Pepperoni Pizza Model.

Ernest Rutherford’s Gold Foil Experiment

• 1911: Ernest Rutherford, Hans Geiger, and Ernest Marsden fire alpha particles at a thin piece of gold foil.

Rutherford’s Results

• Most of the alpha particles passed right through; whereas, a few were either deflected or greatly deflected.

• Rutherford concluded that the nucleus is small, dense, and positively charged.

The Rutherford Atomic Model• Based on his experiment, we now

know…

• The atom is mostly empty space.

• All of the atom’s mass and positive charge is in the nucleus.

• The nucleus is composed of protons and neutrons.

• The majority of the atom’s volume is the electron cloud.

Counting Subatomic Particles

• Now that scientists have discovered that atoms can be subdivided into subatomic particles, there was a new problem.

• How do we count subatomic particles?

• We use terms like atomic number and mass number to do so.

Atomic Number• Atoms are composed of identical protons,

neutrons, and electrons.

• How are atoms of one element different from those of another element?

• Each element contains a particular number of protons.

• The atomic number of an element is the number of protons in the nucleus.

• # protons in an atom = # electrons (if the atom is neutral!)

Mass Number• The mass number of an element is

the average atomic mass of an element rounded to a whole number.

• This number is equal to the number of protons and neutrons in the nucleus.

• Mass number = p+ + n0Mass number = 195

Element # p # e # n Mass no.

Carbon 6

Nitrogen 7 14

Sodium 11 12

Uranium 92 238

Radon 136

Nuclear Symbols• Contain the symbol of the element, the mass

number, and the atomic number.

Nuclear Symbols• Find each of these

• a) number of protons

• b) number of neutrons

• c) number of electrons

• d) atomic number

• e) mass number

Nuclear Symbols• If an element has an atomic number of 34 and a

mass number of 78, what is the…

• a) number of protons

• b) number of neutrons

• c) number of electrons

• d) complete nuclear symbol

Nuclear Symbols• If an element has 91 protons and 140 neutrons,

what is the…

• a) atomic number

• b) mass number

• c) number of electrons

• d) complete nuclear symbol

Isotopes• Dalton was wrong about all atoms of

elements of the same type being identical.

• Atoms of the same element can have different numbers of neutrons.

• Thus, different mass numbers!

• These atoms are called isotopes.

• 1912: English radiochemist Frederick Soddy proposes the idea of isotopes.

• 1921: Soddy wins Nobel Prize in Chemistry for this work

Frederick Soddy 1877 - 1956

Naming Isotopes

• When referencing isotopes of an element, it is important to indicate which mass number the particular isotope has.

• We typically name isotopes using their element name along with their mass number.

• Ex. carbon-12, carbon-14, hydrogen-1, hydrogen-2

Isotope Protons Neutrons Electrons Mass number

8 10

33 42

15 31

29 63

6 8

Average Atomic Mass• Elements occur in nature as a mixture of

isotopes.

• The percentage of each isotope in the naturally occurring element on Earth is nearly always the same, no matter where the element is found.

• This percentage is taken into account when determining the average atomic mass of a particular element.

• The average atomic mass is the weighted average of the atomic masses of all naturally occurring isotopes of an element.

Calculating Average Atomic Mass• To calculate the average atomic mass of an

element, two pieces of information will be needed: percent abundance of each isotope and the atomic mass of each isotope.

AAM = (mass1 x %A1) + (mass2 x %A2) + …

*Percent abundance (%A) must be in decimal form in order to conduct calculation.

Calculating AAM Ex. 1• Oxygen has three naturally occurring isotopes,

oxygen-16, oxygen-17, and oxygen-18. Oxygen-16 has a percent abundance of 99.762%, oxygen-17 has an abundance of 0.038%, and oxygen-18 has an abundance of 0.200%. The atomic masses of the three isotopes are 15.995 amu, 16.999 amu, and 17.999 amu, respectively. Calculate the average atomic mass of oxygen.

Calculating AAM Ex. 2• There are three naturally occurring isotopes of

neon. Their percent abundances and atomic masses are: neon-20, 90.51%,19.99244 amu; neon-21, 0.27%, 20.99395 amu; neon-22, 9.22%, 21.99138 amu. Calculate the weighted average atomic mass of neon.

Calculating AAM Ex.3• Naturally occurring strontium consists of the

following isotopes.

• Calculate the weighted average atomic mass of strontium.

Isotope Atomic mass, amu Percent abundance

Strontium-84 83.913 0.56

Strontium-86 85.909 9.86

Strontium-87 86.909 7.00

Strontium-88 87.906 82.58

Calculating AAM Ex.4• The two naturally occurring isotopes of nitrogen are

nitrogen-14, with an atomic mass of 14.003074 amu, and nitrogen-15, with an atomic mass of 15.000108 amu. What are the percent abundances of these isotopes?

Relating Mass to Number of Atoms

• How can we determine the number of atoms in a particular number of grams of a substance?

• We can use three concepts: the mole, Avogadro’s number, and molar mass. These three concepts allow us to relate atoms and mass.

The Mole

• Remember, that the mole is the SI base unit for measuring amount of substance.

• In particular, one mole is equal to the number of particles as there are atoms in exactly 12 grams of carbon-12.

• The mole is simply a counting unit like a dozen.

Avogadro’s Number (il numero d’Avogadro)

• Italian chemist Count Amedeo Avogadro devised a way to count the number of representative particles of a substance.

• Avogadro’s number is the number of representative particles in exactly 1 mole of a pure substance.

Avogadro’s number = 6.022 x 1023 particles

Molar Mass• We can also define a mole in terms of the

amount of substance that contains Avogadro’s number of particles.

• The mass of one mole of a pure substance is called the molar mass of that substance. Molar mass is measured in grams/mole.

• The molar mass of an element is numerically equal to the atomic mass of the element.

Molar Conversions• We can convert between particles, moles, and grams.

molesatoms grams

x molar massx NA

÷ molar mass

÷ NA