Unit 2 – Atomic Structure & Nuclear Chemistry Part I – Atomic Theory, Subatomic Particles, and Average Atomic Mass

Part I Key Terms Atomic mass - The mass of an atom of a chemical element expressed in atomic mass units. It is approximately equivalent to the number of protons and neutrons in the atom (the mass number) Average atomic mass – Weighted average of all atoms of a particular element and is dependent on the mass of isotopes for an element and the relative population of each isotope Bohr model - Devised by Niels Bohr, depicts the atom as a small, positively charged nucleus surrounded by electrons that travel in circular orbits around the nucleus Dalton’s Postulates - States that matter is composed of extremely small particles called atoms; atoms are invisible and indestructable; atoms of a given element are identical in size, mass, and chemical properties; atoms of a specific element are different from those of another element; different atoms combine in simple whole-number ratios to form compounds; in a chemical reaction, atoms are separated, combined, or rearranged Isotope -Atoms of the same element with different numbers of neutrons

Part I Key Terms (cont.) Isotope notation - Subscripts and superscripts can be added to an element’s symbol to specify a particular isotope of the element and provide other important information. The atomic number is written as a subscript on the left of the element symbol, the mass number is written as a superscript on the left of the element symbol Mass number - The total number of protons and neutrons in a nucleus. Subatomic particles - The three kinds of particles that make up atoms: protons, neutrons, and electrons Theory - An explanation supported by many experiments; is still subject to new experimental data, can be modified, and is considered valid if it can be used to make predictions that are proven true

Early Development of Atomic Theory Major Contributors to Understanding Atomic Structure Democritus – ancient Greek philosopher that originally stated all matter consists of atoms 1605: Francis Bacon – published the scientific method 1803: John Dalton – Postulates of Atomic Theory 1897: J.J. Thomson – Discovery of the negatively charged electron and the mass to charge ratio of the electron 1908 Robert Millikan – Determines the charge of the electron 1911: Ernest Rutherford – Discovers positively charged nucleus 1913: Niels Bohr – Theorizes structure of the electron cloud with energy levels and planetary orbits of electrons 1932: James Chadwick – Discovers neutrons

Atomic Theory – John Dalton English physicist Experimented extensively with multiple gases and gaseous compounds Contributions – Five Postulates of Atomic Theory  1. All matter consists of tiny particles called atoms  2. Atoms are indestructible and unchangeable.  3. Elements are characterized by the mass of their atoms.  4. When elements react, their atoms combine in simple, whole number ratios.  5. When elements react, their atoms sometimes combine in more than one simple whole, number ratio.

Dalton’s Model of an Atom He made no prediction about the construction of atoms believing them to be solid spheres. Conclusions made based on his experiments and postulates: Law of the Conservation of Mass – when chemical reactions occur, the atoms are only rearranged and there is no difference in mass following a chemical reaction Law of Definite Proportions – elements combine in simple, low number ratios to form compounds (examples – H20, CO2) Law of Multiple Proportions –elements combine in different simple, low number ratios to form different compounds (examples – H20 and H202; CO and CO2)

Atomic Theory – J.J. Thomson Discovered the negatively charged electron and the mass to charge ratio of the electron Used cathode ray tube Beam of electrons deflected toward positive plate indicated the electron has negative charge Amount of deflection indicates the mass to charge ratio

Thomson’s Experiment Image used courtesy of http://www.chemteam.info/AtomicStructure/Disc-of-Electron-Images.html

Thomson’s Model of the Atom Plum Pudding Model Image used courtesy of http://www.kutl.kyushu-u.ac.jp/seminar/MicroWorld1_E/Part2_E/P24_E/Thomson_model_E.htm

Atomic Theory – Ernest Rutherford Discovered positively charged nucleus Used gold foil & detector ring Fired alpha particles at foil which are positively charged Most went through – atom mostly empty space Some deflected – nucleus positively charged Some bounced back – solid mass indicates nuclear core

Rutherford’s Experiment

Rutherford’s Model of the Atom Nuclear Atomic Model Image used courtesy of http://www.bbc.co.uk/manchester/content/articles/2008/09/10/100908_rutherford_physics_feature.shtml

Atom Theory – Niels Bohr Discovered electrons reside in energy levels with discrete amounts of energy Mathematic modeling Needed to explain why negatively charged electrons do not get absorbed into positively charged nucleus Used information from Balmer, Lyman, & Paschen series Emission spectra for Hydrogen explained by Rydberg equation

Bohr’s Model of the Atom Electron Shell Model Image used courtesy of http://www.blurtit.com/q982327.html

2 Regions of the Atom Nucleus Electron Cloud Contains the protons and neutrons Accounts for virtually all of the mass, but only a very small portion of the volume of the atom. Has a positive charge equal to the number of protons. Electron Cloud Contains the electrons in orbitals Has virtually no mass, but accounts for virtually all of the volume Has a negative charge equal to the number of electrons.

Subatomic Particles Electrons Protons Neutrons Charge = -1 Mass ≈ 0 amu Location: in orbitals in the electron cloud (outside the nucleus) Protons Charge = +1 Mass = 1 amu Location: Inside the nucleus Neutrons Charge = 0

Properties of the Atom Mass Charge Atomic Number Measured in Atomic Mass Units (amu) Equal to the sum of the number of protons and neutrons Represented by the Mass Number Charge Neutral unless electrons gained or lost (ionized) Number of electrons and protons is equal and, therefore balance out Atomic Number Equal to the number of protons Define the element and its chemical properties

Example assuming neutral atom of Fluorine Symbology Example assuming neutral atom of Fluorine Atomic number: 9 Mass Number: 19 Protons: 9 Neutrons: 10 (mass number – atomic number) Electrons: 9 F 19 9

Isotopes Atoms of the same element with different mass due to different number of neutrons

Isotope Atomic Mass (amu) Average Atomic Mass Weighted average of all atoms of a particular element Dependent on the mass of isotopes for an element and the relative population of each isotope % mass oxygen-16: (15.99491) (.99759) = 15.9564 % mass oxygen-17: (16.99913) (.00037) = 0.0063 % mass oxygen-18: (17.99916) (.00204) = 0.0367 Average Atomic Mass of Oxygen = 15.9994 Isotope Isotope Atomic Mass (amu) Population (%) Oxygen-16 15.99491 99.7590 Oxygen-17 16.99913 0. 037 Oxygen-18 17.99916 0.20400

Naming Isotopes Name of the element followed by the mass number of the isotope Carbon – 12 = the name of the carbon atom with a mass number of 12 (6 protons and 6 neutrons) Carbon – 14 = the name of the carbon atom with a mass number of 14 (6 protons and 8 neutrons) Fluorine – 19 = the name of the Fluorine atom with a mass number of 19 (9 protons and 10 neutrons)

Number of electrons (2n2) Energy Levels Energy levels correspond to the energy of individual electrons. Each energy level has a discrete numerical value. Different energy levels correspond to different numbers of electrons using the formula 2n2 where “n” is the energy level Energy Level Number of electrons (2n2) 1 2(12) = 2 2 2(22)= 8 3 2(32)= 18 4 2(42)= 32 n 2n2

Quantum Mechanical Model of Atomic Structure 1900: Max Planck – Develops law correlating energy to frequency of light 1905: Albert Einstein – Postulates dual nature of light as both energy and particles 1924: Louis de Broglie – Applies dual nature of light to all matter 1927: Werner Heisenberg – Develops Uncertainty Principle stating that it is impossible to observe both the location and momentum of an electron simultaneously 1933: Erwin Schrodinger – Refines the use of the equation named after him to develop the concept of electron orbitals to replace the planetary motion of the electron

Orbitals Impossible to determine the location of any single electron Orbitals are the regions of space in which electrons can most probably be found Four types of orbitals s – spherically shaped p – dumbbell shaped d – cloverleaf shaped f – shape has not been determined Each additional energy level incorporates one additional orbital type Each type of orbital can only hold a specific number of electrons

Total # of electrons per orbital type Orbital Types Orbital Type General Shape Orbital Sublevels # of electrons per sublevel Total # of electrons per orbital type s Spherical 1 2 p Dumbbell 3 6 d Clover leaf 5 10 f unknown 7 14

Electron Configuration Energy Level Orbital Type Orbital Sublevel # of orbitals per energy level (n2) # of electrons per orbital type # of electrons per energy level (2n2) 1 s 2 p 3 4 6 8 d 5 9 10 18 f 7 16 14 32

Electron Configuration Notation Find the element on the periodic table Follow through each element block in order by stating the energy level, the orbital type, and the number of electrons per orbital type until you arrive at the element. 1s 2s   2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

Samples of e- Configuration Element Electron Configuration H 1s1 He 1s2 Li 1s2 2s1 C 1s2 2s2 2p2 K 1s2 2s2 2p6 3s2 3p6 4s1 V 1s2 2s2 2p6 3s2 3p6 4s2 3d3 Br 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 (Note the overlap) Pb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2

Noble Gas Electron Configuration Notation Find element on the Periodic Table of Elements Example: Pb for Lead Move backward to the Noble Gas immediately preceding the element Example: Xenon Write symbol of the Nobel Gas in brackets Example: [Xe] Continue writing Electron Configuration Notation from the Noble Gas Example: [Xe] 6s2 4f14 5d10 6p2

Valence Electrons The electrons in the highest (outermost) s and p orbitals of an atom The electrons available to be transferred or shared to create chemical bonds to form compounds Often found in incompletely filled energy levels

Valence Electrons Shortcut to finding valence electrons for main group elements Family 1A (1) 1 valence electron Family 2A (2) 2 valence electrons Family 3A (13) 3 valence electrons Family 4A (14) 4 valence electrons Family 5A (15) 5 valence electrons Family 6A (16) 6 valence electrons Family 7A (17) 7 valence electrons Family 8A (18) 8 valence electrons Family 3-12 have multiple possibilities and shortcuts do not work

Electron Dot Notation Electron configuration notation using only the valence electrons of an atom. The valence electrons are indicated by dots placed around the element’s symbol. Used to represent up to eight valence electrons for an atom. One dot is placed on each side before a second dot is placed on any side. Valance Electrons: Sodium Magnesium Chlorine Neon 1 2 7 8 Electron Dot Notation:   • • •• •• Na Mg : Cl : : Ne : • • •• Oxidation Numbers: +1 +2 -1 0

Part II Key Terms Alpha particle: A helium nucleus emitted by some radioactive substances Beta particle: An energetic electron or positron produced as the result of a nuclear reaction or nuclear decay Beta radiation: Radioactive decay in which an electron is emitted Electron Configuration Notation -Consists of an element’s symbol, representing the atomic nucleus and inner-level electrons, that is surrounded by dots, representing the atom’s valence electrons. Emission spectrum: The range of all possible wave frequencies of electromagnetic radiation, waves created by the systematic interactions of oscillating electric and magnetic fields Energy Levels - A certain volume of space around the nucleus in which an electron is likely to be found. Energy levels start at level 1 and go to infinity. Excited state: The state of an atom when one of its electrons is in a higher energy orbital than the ground state.

Part II Key Terms (cont.) Gamma radiation: Electromagnetic radiation emitted during radioactive decay and having an extremely short wavelength Ground state: The lowest energy state of an atom or other particle Nuclear fission: Splitting of the nucleus into smaller nuclei Nuclear fusion: Combining nuclei of light elements into a larger nucleus Nucleon: a constituent (proton or neutron) of an atomic nucleus Planck’s constant: As frequency increases, the energy of the wave increases Radioactive decay: Spontaneous release of radiation to produce a more stable nucleus Radioactive isotope: An isotope (an atomic form of a chemical element) that is unstable; the nucleus decays spontaneously, giving off detectable particles and energy

Electromagnetic (EM) Spectrum The EM Spectrum is the range of all possible wave frequencies of electromagnetic radiation, waves created by the systematic interactions of oscillating electric and magnetic fields The general term for all electromagnetic radiation is light The range of the EM Spectrum is from very low frequency known as radio waves to very high frequency known as gamma radiation The visible spectrum of light is in the center portion of this EM Spectrum All EM Spectrum travels at the same speed in a vacuum – this speed is known as the speed of light, 3.00 x 108 m/s

EM Spectrum Image used courtesy of http://9-4fordham.wikispaces.com/Electro+Magnetic+Spectrum+and+light

Speed of Light and Frequency Since the speed of all EM radiation is the same, there is a clear mathematical relationship between the frequency of the light and its wavelength All waves travel at a speed that is equal to the product of its frequency (the reciprocal of time) and its wavelength (distance) c = f λ The speed of EM radiation is fixed at 3.00 x 108 m/s Therefore: 3.00 x 108 m/s = f λ Speed of light = frequency x wavelength As frequency increases, wavelength decreases. As wavelength increases, frequency decreases Example: If frequency doubles, wavelength is cut in half

As f ↑, λ↓: Calculations If the wavelength of a radio wave is 15 meter, what is its frequency? 3.00 x 108 m/s = f (10 m) (3.00 x 108 m/s) / 15 m = f 2.0 x107 s-1 = f Frequency = 2.0 x107 Hertz If the frequency of gamma radiation is 6.25 x 1022 Hertz, what is its wavelength? 3.00 x 108 m/s = (6.25 x 1022 s-1) λ (3.00 x 108 m/s) / (6.25 x 1022 s-1) = λ 4.80 x10-15 m = f Wavelength = 4.80 x10-15 m

As frequency increases, the energy of the wave increases Planck’s Law Max Planck determined in 1900 there was a mathematical relationship between the energy of EM radiation and the frequency of that radiation: As frequency increases, the energy of the wave increases E = h f Energy = Planck’s constant x frequency E = (6.63 x 10-34 Joule seconds) f

Planck’s Law Calculations Example: If the wavelength of green light is 5.21 x 10-7 meters, what is the energy of this light? 3.00 x 108 m/s = f (5.21 x 10-7 m) (3.00 x 108 m/s) / 5.21 x 10-7 m = f 5.76 x1014 s-1 = f Frequency = 5.76 x1014 Hertz E = (6.63 x 10-34 Joule seconds) (5.76 x1014 s-1) E = 3.82 x10-19 Joules

Implication of Planck’s Law In order to move an electron to a higher energy level, excite an electron, energy must be absorbed to move the electron Since electrons exist in fixed energy levels with a specific amount of energy, the amount of energy needed is a finite amount equal to the difference in the energy associated with the ground state of the electron and the energy associated with the level to which the electron is excited If the energy related to the excited electron is removed, the electron will return to its ground state and the energy released is equal to the energy absorbed to excite it The energy released is released as light The overall result is that every element has a unique spectra of light associated with it and the spectra can be used to identify the element

Nuclear Reactions All nuclear reactions are based on Einstein’s Theory of Relativity At speeds approaching the speed of light, energy and mass are interchangeable E = mc2 Energy = mass x (speed of light)2 Mass can be converted to energy and vice versa

Mass Defect There is a difference between the mass of an atom and the various particles that make up the atom This difference is called the mass defect of the atom This mass defect is the binding energy of the atom In nuclear reactions, the binding energy is released as energy (heat, light, or gamma radiation) and/or particles with measureable mass

Types of Nuclear Reactions Fission – Splitting of the nucleus into smaller nuclei Fusion – Combining nuclei of light elements into a larger nucleus Radioactive Decay – Spontaneous release of radiation to produce a more stable nucleus

Fission Nucleus splits into smaller nuclei when struck by a neutron of sufficient energy Tremendous release of energy When controlled can produce huge amounts of power in nuclear reactors Naturally occurs in uranium and other ores in spontaneous fission Clean source of energy with no carbon footprint Produces radioactive nuclear waste with long term environmental and health considerations

Fission Process

Fission and Nuclear Reactors

Fusion Lighter nuclei (such as hydrogen) combined to form heavier nuclei Tremendous release of energy 2H + 3H  4He + 1n + energy Deuterium Tritium Helium (occurs naturally in water) Powers the sun and stars No practical application to produce usable energy at this time

Fusion Process

Radioactive Decay Spontaneous release of radiation by unstable nuclei in order to increase stability Radiation can be either energy alone (gamma) or energy accompanied by release of a particle (all of the other forms of decay)

Forms of Radioactive Decay Alpha decay – release of alpha particle and energy Beta decay – release of beta particle and energy Gamma Emission – release of electromagnetic radiation (energy) Positron Emission – release of a positron and energy Electron Capture – absorption of and electron and release of energy Neutron Emission – release of a free neutron and energy

Alpha Decay Typically found in heavier nuclei and the means to achieve stability is to reduce mass Nuclei shed mass in the form of a helium nucleus to become more stable Helium nucleus that is released is ionized and called and Alpha Particle

Alpha Decay (cont.) Alpha Particle is positively charged (no electrons present) Alpha Particles are very massive, but travel slower (low penetrating power) Can cause significant tissue damage if not shielded Shielding can be accomplished with clothing or paper

Alpha Decay Process

Beta Decay Common in nuclei of any size where instability is caused by the number of neutrons Neutron decays into a proton and an electron Proton remains in the nucleus The electron leaves the atom and is called a Beta Particle

Beta Decay (cont.) Beta Particle is negatively charged Mass of the nucleus is unchanged Beta particles have very low mass but are travelling at very high speed Beta particles can penetrate through the skin and cause deep tissue damage

Beta Decay Process

Gamma Radiation Nucleus becomes more stable through the release of electromagnetic energy No change in mass No change in the element The Gamma radiation can be reduced by shielding, but Gamma radiation cannot be stopped Usually found with another type of decay, but not always

Radioactivity Decay Comparison Radioactive Decay Type Mass Charge Penetrating Power Transmutation Alpha 4 amu Positive Low New Element Formed Beta 0 amu Negative High Gamma None (no particle) Extremely High No

Nuclear Reaction Mass Conservation All nuclear reactions must conserve the overall mass of the particles involved in the reaction Two properties must be the same on both sides of a nuclear equation Total Mass Number – the sum of the mass numbers of all particles must be the same on both sides of the reaction Total Atomic Number – the sum of the atomic numbers of all particles must be the same on both sides of the reaction