Part 01 - Notes: Isotopes and Atomic Mass

Part 01 - Notes: Isotopes and Atomic Mass & PT Objectives: Identify, define, and explain: atomic number, zero net charge, atomic mass, mass number, isotope, nuclear charge, atomic mass unit, proton, neutron, electron, nucleus, atomic mass, mass number, amu, and weighted average. Describe the structure of an atom, including p +, n 0, and e with respect to the nucleus and an atom s relative size. Explain how atomic number identifies an element. Write isotopes in correct notation and use notation to determine information about p +, n 0, and e of an isotope. Differentiate between mass number and atomic mass with regard to the determination and the whole number versus decimal. Calculate the average atomic masses of an element from isotope data. Determine which isotope (of a pair) is most abundant based on the molar mass of the element. Text Reference: Chang and Goldsby (Chemistry: The Essential Concepts, McGraw-Hill, 2014) pp. 36-39 and 61 62. Text Vocabulary: Electron (p32) a subatomic particle that has very low mass and carries a single electric charge Nucleus (p34) the central core of an atom Proton (p34) subatomic particle having a single (+) electric charge; the mass of the proton is about 1840 that of an electron Neutron (p35) a subatomic particle that has no net electric charge; its mass is sightly greater than a proton Atomic number (Z) (p36) the number of protons in the nucleus of an atom Isotope (p36) atoms having the same atomic number but different mass numbers Mass number (A) (p36) the total number of neutrons and protons present in the nucleus of an atom Families (p38) the elements in a vertical column of the periodic table Groups (p38) the elements in a vertical column of the periodic table Metal (p38) an element that is a good conductor of heat and electricity and has the tendency to form positive ions in an ionic compound Metalloid (p38) an element with properties intermediate between those of metals and nonmetals Nonmetal (p38) elements that are generally poor conductors of heat and electricity Periodic table (p38) a tabular arrangement of the elements by similarities in properties and by increasing atomic number Periods (p38) horizontal row of the periodic table Alkali metals (p39) the group 1A elements (Li, Na, K, Rb, Cs, and Fr) Alkaline earth metals (p39) group 2A elements (Be, Mg, Ca, Sr, Ba, and Ra) Halogens (p39) the nonmetallic elements in group 7A (group 17) (F, Cl, Br, I, At) Noble gases (p39) nonmetallic elements in group 8A (group 18) (He, Ne, Ar, Kr, Xe, and Rn) Rare gases (p39) see noble gases Atomic mass (p61) the mass of an atom in atomic mass units (amu or u) Atomic mass unit (amu) (p61) a mass exactly equal to one-twelfth the mass of one carbon-12 atom ATOMS & ISOTOPES - Quick Check... 1. The number of protons is called what? 2. What gives the atom its identity? 3. Zero net charge means... 4. Atoms have the same number of what? 5. Which two subatomic particles have roughly the same mass? 6. Which subatomic particle has a mass that is MUCH less that the other two? 7. What makes up most of an atom? 8. What does mass number equal? 9. You have atoms of two different isotopes of tungsten. (a) (b) What will be the same in both isotopes? What is different in the isotopes? page 1 of 2 Revised - 2018-2019

Part 01 - Notes: Isotopes and Atomic Mass & PT 10. How do you find the number of neutrons? 11. Consider: 238 92U (a) atomic number? (d) electrons? (b) mass number? (e) neutrons? (c) protons? (f) also written as... 12. Difference between atom and ion? 13. What does nuclear charge refer to? Atomic Mass What is a fraction? How do you find a fraction? Formula for determining average atomic mass: Example 1: I have a sample of element Andersonium. It consists 18 atoms of isotope An-325, 45 atoms of isotope An-326, and 25 atoms of isotope An-328. The masses of the isotopes are as follows: An-325 is 324.9987 amu, An-326 is 325.9897 amu, and An-328 is 328.0102 amu. Calculate the abundances of each isotope. Example 2: There is a sample of the element chemium (Ch). The element chemium consists of 32.94% of the isotope Ch-340 (with a mass off 339.8998 amu), 43.53% of the isotope Ch-342 (with a mass of 342.0180 amu), and 23.53% of the isotope Ch-343 (with a mass of 343.0074 amu). What is the atomic mass of chemium? Example 3: There is a sample of terrerium. There are two isotopes of terrerium, Tr-463 with a mass of 462.9897 amu and Tr-465 with a mass of 465.0105 amu. If the atomic mass of terrerium is 464.1205, what is the percentage of each isotope of terrerium? Now, go on to POGIL: Weighing in On All the AMU Issues! Part 01 Assignment Task: Complete the following questions from your text book. Show all required work. Questions: 2.11 2.21 2.51 2.53 2.59 3.2 3.5 3.6 3.32 3.93 page 2 of 2 Revised - 2018-2019

Part 02 - Notes: Photons, Bright-Line Spectra, and Photoelectric Effect Objectives: Calculate the wavelength, frequency, or energy of light and state relationship between energy, wavelength and frequency. Explain the origins of the atomic emission spectrum of an element. Differentiate between ground and excited states and explain how it is easier to remove an electron from a higher energy level. Identify, define, and explain: spectrum, atomic emission spectrum, Planck s constant, photons, photoelectric effect, ground state, de Broglie equation, and Heisenberg uncertainty principle. Text Reference: Chang and Goldsby (Chemistry: The Essential Concepts, McGraw-Hill, 2014) pp. 213-217, 218-220, and 220-224. Text Vocabulary Amplitude (p214) vertical distance from the middle of the wave to the peak or the trough Frequency (ν) (p214) number of waves that pass through a particular point per unit time Wave (p214) vibrating disturbance by which energy is transmitted Wavelength (λ) (p214) distance between two identical points on successive waves Electromagnetic radiation (p215) emission and transmission of energy in the form of electromagnetic waves Electromagnetic wave (p215) wave that has an electric field component and a mutually perpendicular magnetic field component Quantum (p217) smallest quantity of energy that can be emitted or absorbed in the form of electromagnetic radiation Photoelectric effect (p218) phenomenon in which electrons are ejected from the surface of certain metals exposed to light of at least a certain minimum frequency Photon (p218) particle of light Emission spectrum (p220) continuous or line spectrum or electromagnetic radiation emitted by a substance Excited level (or state) (p221) state that higher energy than the ground state of the system Ground level (or state) (p221) the lowest energy state of a system Line spectrum (p221) spectrum produced when radiation is absorbed or emitted by a substance only at some wavelengths The wave-particle duality of electromagnetic wave: How is electromagnetic radiation transferred to matter? A wave is a wiggle in time and space. A wave is a disturbance that repeats regularly in time and space and that is transmitted from one place to another with no net transport of matter. Waves transmit energy from one place to another without actually transmitting the matter through which it moves. page 1 of 3

Part 02 - Notes: Photons, Bright-Line Spectra, and Photoelectric Effect Speed of a wave: Speed of an electromagnetic wave: Energy of a photon of electromagnetic radiation: Relationships: What type of relationship exists between frequency and wavelength? What type of relationship exists between frequency and energy of a photon? What type of relationship exists between wavelength and energy of a photon? Practice: (a) (b) Wavelength? Given that 6.0x10 8 crests of the wave pass a given point in 1.0 microsecond, what is the frequency of the wave in hertz? (c) Speed of the wave? (d) The light wave in the figure corresponds to green light. Describe how the value of the wavelength would differ for a figure showing a wave of red light. (e) How would the frequency of the red light wave compare to that of green light? The Photoelectric Effect Photoelectrons - electrons emitted from the surface of a metal when light of sufficient energy is shined on the metal Work function - amount of energy required to be put into metal to dislodge the electron from the surface of the metal If the incident light does not have enough energy... If the incident light is equal to the threshold energy... page 2 of 3

Part 02 - Notes: Photons, Bright-Line Spectra, and Photoelectric Effect If the incident light is above threshold energy... If the intensity of the incident light is increased... The number of electrons ejected depends on... The energy of the ejected electrons depends on... E = W + KE hυ = W + KE Example: Work function of titanium metal is 6.93x10 19 J. (a) What is the minimum frequency of light required to release electrons from titanium? (b) Calculate the kinetic energy of the ejected electrons if light of frequency 2.50x10 15 Hz is used for irradiating the metal? The Bright-Line Spectrum Consider the experiment when white light is produced and it is then shined on a thin slit. The light that passes through the slit is then incident on a prism. Laws of classical physics there is no limit to how large or how small the energy gained or lost by an object may be. So according to this, the bright-line spectrum should be a continuous rainbow. But it is NOT. We cannot always apply macroscopic laws to subatomic events. Max Planck German physicist 1858-1947 Question: Why does an object change color when heated? Heating iron causes it to change from black to yellow to red to white to blue as its temperature is increased. Answer: The energy of a body changes only in small discrete units quanta small packages of energy. It appears that thermal energy may continuously be supplied to heat liquid water to any temperature between 0 o and 100 o C. Actually the water temperature increases by infinitesimally small steps, which occurs as individual water molecules absorb quanta of energy. An ordinary thermometer is unable to detect small changes in temperature. Thus your everyday experiences give you no clue to the fact that energy is quantized. You do NOT need to memorize Equations 7.5 (page 221) or Equation 7.6 (page 222). page 3 of 3

Part 03 - Notes: Electrons and Orbitals in Atoms Objectives: Explain the significance of quantized energies of electrons in relation to the quantum mechanical model. Identify, define, and explain: energy level, quantum, quantum mechanical model (wave mechanical model), atomic orbitals, quantized, sublevel, ground state, excited state, probability, and principal quantum number. Explain Aufbau Principle, Hund s Rule, and Pauli Exclusion Principle and how they are used to locate an electron in an atom. List, describe, and differentiate between s, p, d, and f atomic orbitals. Relate the principal quantum number to the number of orbital types and to the number of individual orbitals in a given energy level. Text Reference: Chang and Goldsby (Chemistry: The Essential Concepts, McGraw-Hill, 2014) pp. 224 228, & 230-234. Text Vocabulary: Node (p225) point at which the amplitude of the wave is zero Heisenberg uncertainty principle (p227) It is impossible to know simultaneously both the momentum and the position of a particle with certainty. Atomic orbital (p228) wave function of an electron in an atom Electron density (p228) probability that an electron will be found at a particular region in an atomic orbital Many-electron atom (p228) atoms that contain two or more electrons Quantum numbers (p228) numbers that describe the distribution of electrons in atoms Recall the Bohr Model of the atom: The energies of electrons are QUANTIZED meaning only certain values are allowed. But noone, not even Bohr, could explain why the energies of the hydrogen electron was quantized. Einstein originally proposed that light can behave like a stream of particles (photons) - Wave-Particle Duality Louis Victor De Broglie and Erwin Schrödinger extended Einstein s wave-particle description of light to all matter in motion. (P225) If an electron does behave like a standing wave in the hydrogen atom, the length of the wave must fit the circumference of the orbit exactly. Note the distance between the nodes on this standing wave is equal to half a wavelength. De Broglie s Equation: Energy-mass conversion: E = m c 2 Energy of a photon: E = h ν Example 1: (Text 7.40) Protons can be accelerated to speed near that of light in particle accelerators. Calculate the wavelength (in nm) of such a proton moving at 2.90x10 8 m/s. The mass of a proton is 1.675x10 27 kg. Heisenberg Uncertainty Principle It is impossible to know simultaneously both the momentum (mass x velocity) and the position of a particle with certainty. This uncertainty is more obvious and significant with small objects than with large objects. Comparing an electron and a baseball there is the uncertainty with the baseball but it is so small it is nearly immeasurable. But, with regard to an electron, the uncertainty is much more significant due to the extremely small size of an electron. page 1 of 2

Part 03 - Notes: Electrons and Orbitals in Atoms The Quantum Mechanical Model a modern description of electrons in atoms Principal energy levels Principal Quantum Number: Sublevels: Atomic Orbital: An orbital is a probability map, a region in which the electron has a 90% ( or 95%) probability of being found. There are four basic types of orbitals: s, p, d, and f. There are additional orbitals designated g, h, i, j, k, Shape Groups of Begin where? Relative energy s-orbital p-orbital d-orbital f-orbital Let s explore several relationships within the atom by visiting The Aufbau Hotel... page 2 of 2

Part 04 - Notes: Electron Configuration & Diagrams Objectives: Apply the Aufbau principle, the Pauli Exclusion Principle, & Hund s Rule in writing electron configurations and drawing orbital diagrams. Identify, define, and, explain: electron configuration, Aufbau principle, Hund s Rule, Pauli Exclusion Principle, and orbital diagram. Explain why all of the 3 rd energy level does not fill before the 4 th energy level starts to fill. Text Reference: Chang and Goldsby (Chemistry: The Essential Concepts, McGraw-Hill, 2014) pp. 234 240 and 241-244. Text Vocabulary: Many-electron atom (p228) atoms that contain two or more electrons Boundary surface diagram (p231) diagram of the region containing about 90 percent of the electron density in the atomic orbital Electron configuration (p235) distribution of electrons among the various orbitals of an atom or molecule Pauli exclusion principle (p235) no two electrons in an atom can have the same four quantum numbers Hund's rule (p238) The most stable arrangement of electrons in atomic subshells is the one with the greatest number of parallel spins. Aufbau principle (p241) As protons are added one by one to the nucleus to build up the elements, electrons similarly are added to the atomic orbitals Noble gas core (p241) the noble gas that most nearly precedes the element being considered; used in writing electron configuration Key Items to Keep in Mind... 1. Energy levels, n, increase by integer values. 2. The number of orbitals per energy level is equal to n 2. The first energy level will have one orbital. The second energy level, n = 2, will have 4 orbitals. The third energy level, n = 3, will have nine orbitals. And so on. 3. Every energy level has an s orbital. The s-orbital is sphere shaped. 4. Energy levels n = 2 and higher contain a set of three p-orbitals: px, py, and pz. The p-orbitals have two lobes. 5. Energy levels n = 3 and higher contain a set of five d-orbitals: dxy, dxz, dyz, dz2, and dx2-y2 that have four lobes. 6. Energy levels n = 4 and higher contain a set of seven f-orbitals. The subscripts are not important for our purposes; just know the total number. They also have shapes more complex than can be simply described. 7. Aufbau Principle Electrons fill orbitals from lowest to highest energy. 8. Pauli Exclusion Principle No two electrons in the same atom can have the same four quantum numbers. In other words, an orbital can hold at most two electrons and if there are two electrons, they have paired spins. 9. Hund s Rule When electrons occupy a set of degenerate orbitals, one electron enters each orbital until all orbitals contain one electron with their parallel spins. How to Read Electron Configuration: 1 s 2 Read it as... The 1 represents The s represents Box Diagrams: The superscript 2 indicates Electron Configuration and Orbital Diagrams for the Selected Elements At this time, all elements are considered to be in their ground state (the state of least energy). hydrogen helium lithium page 1 of 4

Part 04 - Notes: Electron Configuration & Diagrams beryllium boron carbon nitrogen oxygen fluorine neon sodium magnesium aluminum silicon phosphorus sulfur chlorine argon Abbreviated Configurations - based on the noble gas configuration of the noble gas the most nearly precedes the element What happens after the 18 th electron fills into the 3p sublevel? What orbital do you think will fill next? Examine the Orbital Filling Chart. Why is this so? potassium calcium page 2 of 4

Part 04 - Notes: Electron Configuration & Diagrams Now, scandium has 21 electrons. Where does the 21 st electron go? scandium So what is the filling order? Practice: nickel arsenic krypton lead Valence Electrons Valence electrons - electrons in the outermost energy level of an atom highest E level # of valence e calcium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 nitrogen 1s 2 2s 2 2p 3 nickel 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 In the above table of electron configurations for sodium through argon, list the number of valence electrons in the final column. Notice any pattern? How many valence electrons in... nickel? krypton? arsenic? lead? The valence electrons are the most significant to chemists because those are the electrons involved in bond breaking and bond formation. The inner electrons are known as core electrons. The core electrons are not involved in bonding atoms to one another. page 3 of 4

Part 04 - Notes: Electron Configuration & Diagrams Consider the electron configurations of the first three elements in the alkali metal family. Abbrev Config # val. e hydrogen 1 s 1 lithium 1 s 2 2 s 1 sodium 1 s 2 2 s 2 2 p 6 3 s 1 Patterns? Conclusions? What does this information say about the similarity of chemical properties within elements of the same chemical family? page 4 of 4

Part 05 - Notes: Anomalies & Quantum Numbers Objectives: Write abbreviated configurations for elements. Explain why the electron configurations for the chromium and copper columns do not fit the standard pattern. Correctly write the electron configurations for chromium molybdenum, tungsten, copper, silver, and gold, and explain why it is written in such a manner. Identify, define, and explain: quantum number, principal quantum number, angular momentum quantum number, magnetic quantum number, and spin quantum number. Explain why no two electrons in a given atom can have the same four quantum numbers. Text Reference: Chang and Goldsby (Chemistry: The Essential Concepts, McGraw-Hill, 2014) pp. 228 230 and 234-244. Text Vocabulary: Quantum numbers (p228) numbers that describe the distribution of electrons in atoms Diamagnetic (p236) repelled by a magnet; a diamagnetic substance contains only paired electrons Paramagnetic (p236) attracted to a magnet, a paramagnetic substance contains one or more unpaired electrons Transition metals (p241) elements that have incompletely filled d sublevels or readily give rise to cations that have incompletely filled d sublevels Actinide series (p243) elements that have incompletely filled 5f subshells or readily readily give rise to cations that have incompletely filled 5f subshells Lanthanide (rare earth) series (p243) elements that have incomplete 4f subshells or readily give rise to cations that have incompletely filled 4f subshells Rare earth series (p243) see lanthanide series Anomalies of Electron Configuration Based on your current knowledge, write the electron configuration of chromium and copper. Chromium: Copper: These electron configurations are actually INCORRECT!!! Take a moment to cross them out. Atoms in the ground state have the lowest possible energy. That is why some of them form ions and some form covalent compounds, to obtain a state of minimum energy. The Aufbau Principle indicates that electrons fill into the orbitals in a manner that minimizes the energy of the atom. The correct electron configuration for chromium: Why does this happen? The correct electron configuration for copper: Why does this happen? page 1 of 3

Part 05 - Notes: Anomalies & Quantum Numbers There are a number of other anomalies that occur throughout the periodic table. You will be responsible for the anomalies of chromium, copper, molybdenum, silver, tungsten, and gold. Mo, W, Ag, and Au follow the same valance configuration as chromium and copper, based upon the column in the periodic table. There are other anomalies for which you are not held responsible. Note the following rules for writing the electron configurations: 1. In a principal energy level that has d orbitals, the s-orbitals from the next level before the d-orbital in the current level. 2. After lanthanum, which has the electron configuration [Xe] 6s 2 5d 1, a group of fourteen elements called the lanthanide series occurs. This series of elements corresponds to the filling of the seven 4f orbitals. 3. After actinium, which has the electron configuration [Rn] 7s 2 6d 1, a group of fourteen elements known as the actinide series occurs. This series corresponds to the filling of the seven 5f orbitals. Paramagnetic Versus Diamagnetic Consider the following electron configurations and determine if the elements are paramagnetic or diamagnetic. Also determine the number of valence electrons and how many (if any) unpaired electrons exist in the atom. tellurium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 4 technetium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 5 mercury 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 osmium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 6 rutherfordium 1s2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 6d 2 5f 14 silver 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 4d 10 Quantum Numbers The exact location of an electron cannot be pinpointed; however, the electron may be described using quantum numbers. There are four quantum numbers. The numbers represent an electron s address within an atom. No two electrons in the same atom can have the same address at the same time; this is the Pauli Exclusion Principle. The Principal Quantum Number = n page 2 of 3

Part 05 - Notes: Anomalies & Quantum Numbers The Angular Momentum Quantum Number = l The Magnetic Quantum Number = m Spin Quantum Number = ms The Pauli Exclusion Principle says that no two electrons in the same atom may have the same four quantum numbers. Justify this statement. page 3 of 3