Unit 3 Atoms SubAtomic Particles Nuclear Chemistry The

Unit 3 Atoms, Sub-Atomic Particles & Nuclear Chemistry

The Particle Theory of Matter • In 400 B. C. Democritus, Democritus a Greek philosopher, first proposed the idea of a basic particle of matter that could not be divided any further. • He called this particle the atom, atom based on the Greek word atomos meaning indivisible. • This early theory was not backed up by experimental evidence and was ignored by the scientific community for nearly 2000 years.

Foundations of Atomic Theory • By the late 1700 s, experiments with chemical reactions led to the discovery of 3 basic laws: 1. The Law of Conservation of Mass • This law states that mass is neither created nor destroyed during ordinary chemical reactions or physical changes. • Formulated by Antoine Lavoisier in 1789.

Foundations of Atomic Theory The Law of Conservation of Mass

Foundations of Atomic Theory 2. The Law of Definite Proportions • States that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound. • Discovered by Joseph Proust in 1797. Visual Concept

Foundations of Atomic Theory 3. The Law of Multiple Proportions – • States that if different compounds are composed of the same 2 elements, then the ratio of the masses of the elements is always a ratio of small whole numbers. • Published by John Dalton in 1804.

Dalton’s Atomic Theory • In 1808, John Dalton proposed an explanation for the three laws. His atomic theory states: 1. All matter is composed of atoms. 2. Atoms of the same element are identical; atoms of different elements are different. 3. Atoms cannot be subdivided, created, or destroyed. 4. Atoms of different elements combine in simple wholenumber ratios to form chemical compounds. 5. In chemical reactions, atoms are combined, separated, or rearranged.

Corrections to Dalton’s Theory • Dalton turned Democritus’s idea into a scientific theory that could be tested by experiment. • But not all aspects of Dalton’s theory have proven to be correct. We now know that: – Atoms are divisible into even smaller particles. – A given element can have atoms with different masses.

Dalton’s Atomic Model • An atom is the smallest particle of an element that has all the properties of that element. • Atoms are too small to see…even through the most powerful microscope!! • Dalton thought atoms were solid balls of matter and were indivisible.

Discovery of the Electron • In 1897, Joseph John (JJ) Thomson showed that cathode rays are composed of identical negatively charged particles, which were named electrons • The electron was the first subatomic particle to be discovered.

Thomson’s Cathode Ray Tube Experiment Visual Concept

Charge and Mass of the Electron • In 1909, Robert Millikan measured the charge on the electron during his oil drop experiment • Using the charge-to-mass ratio, scientists were able to figure out the mass of the electron: about 1/2000 the mass of a hydrogen atom.

Millikan’s Oil Drop Experiment Visual Concept

Thomson’s Plum Pudding Model • After the work of Thomson and Millikan, the accepted model of the atom was called the plum pudding model • The atom was viewed as a ball of positivelycharged material with tiny negatively-charged electrons spread evenly throughout.

Discovery of the Atomic Nucleus • More detail of the atom’s structure was provided in 1911 by Ernest Rutherford and his associates Hans Geiger and Ernest Marsden. • The results of their gold foil experiment led to the discovery of a very densely packed bundle of matter with a positive electric charge. • Rutherford called this positive bundle of matter the nucleus

The Gold Foil Experiment Visual Concept

Rutherford’s Atomic Model • After Rutherford’s gold foil experiment, the accepted model of the atom looked like this: • A small, positively-charged nucleus with negative electrons surrounding it at some distance away. Most of the atom is empty space

Structure of the Atom • Rutherford proposed that the nucleus had particles with the same amount of charge as an electron but the opposite sign, called protons – Relative charge = +1 – Relative mass = 1 amu • For an atom to be neutral there must be equal numbers of protons and electrons. • Throughout the 1920’s scientists accepted an (incorrect) model of the atom composed of protons and electrons.

Some Problems • How could beryllium have 4 protons stuck together in the nucleus? – shouldn’t they repel each other? • If a beryllium atom has 4 protons, then it should weigh 4 amu; but it actually weighs 9. 01 amu! Where is the extra mass coming from? – Each proton weighs approximately 1 amu. – The electron’s mass is only about 0. 00055 amu and Be has only 4 electrons, so they don’t account for the extra 5 amu of mass.

There Must Be Something Else There! • These questions were answered in 1932 by James Chadwick (a student of Rutherford’s), who discovered another particle in the nucleus, which he called a neutron – Charge = 0 (no charge). – Relative mass = 1 amu.

Subatomic Particles • The nucleus is made up of at least one positively charged particle called a proton and usually one or more neutral particles called neutrons • Protons, neutrons, and electrons are often referred to as subatomic particles

Atomic Number • The atomic number (Z) of an element is the number of protons of each atom of that element. • Atoms of the same element all have the same number of protons.

Mass Number • The mass number is the total number of protons and neutrons in the nucleus of an atom. • Atoms of the same element can have different mass numbers.

Isotopes • Isotopes are atoms of the same element that have different masses. • Isotopes have the same number of protons and electrons but different numbers of neutrons. • Most of the elements consist of mixtures of isotopes.

Designating Isotopes • Hyphen notation: notation The mass number is written with a hyphen after the name of the element. uranium-235 • Nuclear symbol: symbol The superscript indicates the mass number and the subscript indicates the atomic number. Mass number Atomic number

Calculating Neutrons • The number of neutrons is found by subtracting the atomic number from the mass number − atomic number = number of neutrons • Nuclide is a general term for a specific isotope of an element.

Calculating Subatomic Particles Sample Problem How many protons, electrons, and neutrons are there in an atom of chlorine-37? Solution: Number of protons = atomic number (on periodic table) 17 Number of electrons = number of protons 17 Number of neutrons = mass number - protons 20

The Atomic Mass Unit • The standard used by scientists to compare units of atomic mass is the carbon-12 atom. • One atomic mass unit, unit or 1 amu, amu is exactly 1/12 the mass of a carbon-12 atom. • The atomic mass of any atom is determined by comparing it with the mass of the carbon-12 atom.

Average Atomic Mass • Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element. • The average atomic mass of an element depends on both the mass and the relative abundance of each of the element’s isotopes.

Calculating a Weighted Average If you have the following grades, what would your marking period average be? Category Percent of Grade Tests 45 ÷ 100 = 0. 45 x 72 = 32. 4 Quizzes 10 ÷ 100 = 0. 10 x 78 = 7. 8 Labs 25 ÷ 100 = 0. 25 x 84 = 21. 0 HW/CW 10 ÷ 100 = 0. 10 x 98 = Project 10 ÷ 100 = 0. 10 x 94 = 9. 4 + 9. 8 • First, change percents to decimals. 80. 4 • Next, multiply each grade by its decimal percent. • Finally, add up all the products.

Calculating Average Atomic Mass Sample Problem 1 Copper consists of 69. 15% copper-63, which has an atomic mass of 62. 929 601 amu, and 30. 85% copper-65, which has an atomic mass of 64. 927 794 amu. What is the Average Atomic Mass of Copper? Solution • Change percents to decimals. • Multiply the atomic mass of each isotope by its relative abundance. • Add up all of the products. Relative Abundance Mass Cu-63 0. 6915 x 62. 93 = 43. 52 Cu-65 0. 3085 x 64. 93 = 20. 03 + 63. 55

Calculating Average Atomic Mass Sample Problem 2 A student believed that she had discovered a new element and named it mythium. Analysis found it contained two isotopes. The composition of the isotopes was 19. 9% of atomic mass 10. 013 and 80. 1% of atomic mass 11. 009. What is the average atomic mass, and do you think mythium was a new element? Solution: Round off to: • Average Atomic Mass: 10. 8 (. 199 x 10. 013) + (. 801 x 11. 009) = 10. 811 • Because the atomic mass is the same as the atomic mass of boron, boron mythium was not a new element.

Forces in the Atom • Electrons and protons attract because of opposite electrical charges, charges but protons and protons repel since they have the same charge. • The nucleus is held together by a mysterious force called the strong nuclear force which only exists between nucleons (protons and neutrons) which are very close together.

Valley of Stability for Z = 1 20, stable N/Z ≈ 1 for Z = 20 40, stable N/Z approaches 1. 25 for Z = 40 80, stable N/Z approaches 1. 5 for Z > 83, there are no stable nuclei

Naturally Radioactive Elements • All of the elements beyond atomic number 83 are unstable and thus radioactive.

Nuclear Reactions • Large, unstable nuclei spontaneously break apart to form smaller, more stable nuclei. • A nuclear reaction is a reaction that affects the nucleus of an atom. Example: • A transmutation is a change in the identity of a nucleus as a result of a change in the number of its protons.

Nuclear Reactions Sample Problem Identify the products that balance the following nuclear reactions: a. b. Solution: 208 a. Atomic Mass: 212 = 4 + _____ 82 Atomic Number: 84 = 2 + _____ 0 = 22 b. Atomic Mass: 22 + ____ -1 = 10 Atomic Number: 11 + ____

Radioactive Decay • Radioactive decay is the spontaneous disintegration of a nucleus into a lighter nucleus, accompanied by nuclear radiation. • Nuclear radiation is particles and/or electromagnetic radiation emitted from the nucleus during radioactive decay.

Types of Radioactive Decay Alpha Emission • An alpha particle (α) is two protons and two neutrons bound together and is emitted from the nucleus during some kinds of radioactive decay. • • The atomic number decreases by two and the mass number decreases by 4.

Types of Radioactive Decay (continued) Beta Emission • A beta particle (β) is an electron emitted from the nucleus during some kinds of radioactive decay (a neutron can be converted into a proton and an electron. ) • • The atomic number increases by one and the mass number stays the same.

Types of Radioactive Decay (continued) Positron Emission • A positron (β+) is a particle that has the same mass as an electron, but has a positive charge (to decrease the number of protons, a proton can be converted into a neutron by emitting a positron. ) • • The atomic number decreases by one and the mass number stays the same.

Types of Radioactive Decay (continued) Electron Capture • In electron capture, an inner orbital electron is captured by the nucleus of its own atom. (An inner orbital electron combines with a proton to form a neutron. ) • • The atomic number decreases by one and the mass number stays the same.

Types of Radioactive Decay (continued) Gamma Emission • Gamma rays ( ) are high-energy electromagnetic waves emitted from an unstable nucleus. • Atomic number and mass number both stay the same because gamma rays have no charge and no mass. • They are pure energy, and very dangerous to living things.

Comparing Alpha, Beta and Gamma • Alpha particles are big and slow. They can’t penetrate skin or paper. • Beta particles have about 100 x the penetrating power of alpha. They can be stopped by clothing, wood, or aluminum foil. • Gamma rays have the greatest penetrating ability. They can only be stopped by a thick layer of lead or concrete. They cause a lot of damage to living cells. Visual Concept

Half-Life • Half-life is the time required for half the atoms of a radioactive nuclide to decay. • Each radioactive nuclide has its own half-life. • More-stable nuclides decay slowly and have longer half-lives.

Half-Life Sample Problem 1 The half-life of radon-222 is 4 days. After what time will ¼ of a given amount of radon remain? Solution: Determine the number of half-lives that it takes to cut a sample to ¼ of the original amount. Then multiply that number by the half-life. 2 x 4 days = 8 days

Half-Life Sample Problem 2 Phosphorus-32 has a half-life of 14. 3 days. How many milligrams of phosphorus-32 remain after 57. 2 days if you start with 4. 0 mg of the isotope? Solution: Determine the number of half-lives in 57. 2 days ÷ 14. 3 days = 4 half-lives For each half-life, multiply the original amount by ½: 4. 0 mg x ½ x ½ = 0. 25 mg

Uses of Radiation • Radioactive dating – scientists can determine the approximate age of an object based on the amount of certain radioactive nuclides present.

Uses of Radiation (continued) • Radioactive tracers are radioactive atoms that are incorporated into substances so that movement of the substances can be followed by radiation detectors. –Radioactive tracers can be used by doctors to diagnose diseases. –Radioactive tracers are also used in agriculture to determine the effectiveness of fertilizers.

Uses of Radiation (continued) • Irradiated Food – nuclear radiation is used to prolong the shelf life of food.

Uses of Radiation (continued) • Nuclear Power Plants – use energy as heat from nuclear reactors to produce electrical energy. They have five main components: 1. Shielding – radiation-absorbing material used to decrease exposure to radiation from nuclear reactors. 2. Fuel – usually Uranium-235. 3. Coolant – usually water, it absorbs excess heat energy. 4. Control rods – neutron-absorbing rods that limit the number of free neutrons 5. Moderator – used to slow down the fast neutrons produced by fission.

Nuclear Power Plant

Whole-Body Radiation Exposure • Rem – The unit used to measure the biological effects of absorbed radiation in humans.

Radiation Detection • Film badges - use exposure of film to measure the approximate exposure of people working with radiation. • Geiger-Müller counters - instruments that detect radiation by counting electric pulses carried by gas ionized by radiation. • Scintillation counters - instruments that convert scintillating light to an electric signal for detecting radiation.

Fission vs. Fusion • Nuclear Fission – very heavy nuclei split into smaller, more stable nuclei. – Can occur spontaneously or when nuclei are bombarded by particles. – Controlled fission chain reactions are used in nuclear power plants. Uncontrolled fission chain reactions are used in nuclear bombs.

Fission vs. Fusion (continued) • Nuclear Fusion – low-mass nuclei combine to form a heavier, more stable nucleus. – Fusion releases even more energy per gram of fuel than fission. – Scientists are not yet able to control fusion reactions, so we can’t use them in power plants. – Fusion is the primary process that fuels our sun and the stars.

Nuclear Waste • Radioactive waste produced in nuclear reactors can take hundreds of thousands of years to decay. • Disposal of nuclear waste is done with the intention of never retrieving it. • There are 77 disposal sites around the country. A new one (Yucca Mountain) is being developed for the permanent disposal of much of our nuclear waste beginning in 2017. Visual Concept