CHEM-16 Midterm Coverage PDF

Principles of Chemistry Chem 16 Chapter 1  Lecture Outline 1.1. Atoms 1.2 Dalton’s Atomic Theory 1.3 Subatomic Particles and Atomic Structure 1.4 The Periodic Table 1.5 Atomic Number, Mass Number and Atomic Symbols 1.6 Isotopes, Cations and Anions 1.7 Average atomic mass ...

Principles of Chemistry Chem 16 Chapter 1  Lecture Outline 1.1. Atoms 1.2 Dalton’s Atomic Theory 1.3 Subatomic Particles and Atomic Structure 1.4 The Periodic Table 1.5 Atomic Number, Mass Number and Atomic Symbols 1.6 Isotopes, Cations and Anions 1.7 Average atomic mass  Chemistry The study of: o the composition of matter o the changes that matter undergoes Matter is anything that occupies space and has mass. “Therefore, MATTER is EVERYTHING around us!” 1.1. Atoms the basic unit of matter the defining structure of elements from the greek word “atomos”, meaning indivisible Electrons occupy a region around the nucleus, AND NOT AN ORBIT 1.1. Atoms Bohr Model Quantum Mechanical Model  1.2 Dalton’s Atomic Theory (1804) All matter is composed of extremely small particles called atoms. Atoms cannot be subdivided, created, or destroyed. (Atoms can be created or destroyed or subdivided.) Atoms of a given element are identical in size, mass, and other properties. Atoms of different elements differ in size, mass, and other properties. (Isotops have different mass but same other properties) Atoms of different elements can combine in simple whole number ratios to form chemical compounds. In chemical reactions, atoms are combined, separated, or rearranged. An atom of one element cannot change into an atom of another element. (Nuclear reaction can change an atom into another atom) 1.3 Subatomic Particles and Atomic Structure  The Atomic Mass Masses of atoms and subatomic particles are very small, a conversion unit must be used to measure them. Atomic mass unit (amu or u): defined as 1/12 of the mass of 1 carbon atom. 1 amu = 1.661 × 10-24 g Atomic weight (amu): average of the atomic masses of an element’s isotopes found on earth. Average Atomic Mass = Atomic Weight 1.4 The Periodic Table 1.5 Atomic Number, Mass Number and Atomic Symbols Atomic Number = Number of Protons = Number of Electrons Mass Number Mass Number - Number of Protons = Number of Neutrons 1.5 Atomic Number, Mass Number and Atomic Symbols O 8 16 8 8 8 K 19 38 19 19 19 Br 35 80 35 45 35 Zn 30 65 30 35 30 1.6 Isotopes, Cations and Anions Atoms: Isotopes same atomic number (same number of protons) different atomic mass (due to different number of neutrons) Abundance: 1.6 Isotopes, Cations and Anions Atoms: Ions happens when atoms lose or gain electrons 1.6 Isotopes, Cations and Anions Atoms: Ions happens when atoms lose or gain electrons 1.7 Average atomic mass 𝐶𝑎𝑟𝑏𝑜𝑛 12 99% 𝐶𝑎𝑟𝑏𝑜𝑛 13 1% Atomic weight = (12 𝑥 0.99) + (13 𝑥 0.01) = 11.88 + 0.13 = 12.01 1.7 Average atomic mass Sample Problem 2 Calculate the average atomic mass of Li. Atomic weight = (6.02 𝑥 0.075) + (7.02 𝑥 0.925) = 0.4515 + 6.4935 = 6.945 Quiz Time 1. Given the table below, calculate the average atomic mass of neon, Ne. (3 pts) 2. Find the number of subatomic particles Name of Atomic Atomic # of # of # of Symbol Element Number Mass protons neutrons electrons Oxygen 16O2- 1.) 2.) 3.) 4.) 5.) Oxygen 15O2- 6.) 7.) 8.) 9.) 10.) Calcium 11.) 20 40 12.) 13.) 18 14.) 15.) 16) 17) 15 18.) 18 Iron 19.) 20.) 21) 22) 30 24 Chapter 2 The Mole Concept and the Anatomy of Common Chemical Reactions Lecture OUTLINE 2.1 The mole concept and molar mass 2.2 Formula mass and molecular mass 2.3 Interconversion between mass, moles and elementary entities 2.4 Percent composition of compounds 2.5 Common Chemical Reactions 2.6 Balancing Chemical Reactions 2.1 The mole concept and molar mass The mole (mol) is the amount of a substance that contains as many elementary entities as there are atoms in exactly 12.00 grams of Carbon 12 1 mol = NA = 6.0221367 x 1023 Avogadro’s number (NA) The mole concept and molar mass 1 mol = NA = 6.0221367 x 1023 atoms in 12 grams of Carbon 12 (12C). The mole concept and molar mass 1 mol Element/Molecule = 6.0221367 x 1023 atoms Element = 6.0221367 x 1023 Molecule = amu of Element (ex. 12g 12C, 6.9g Li) The mole concept and molar mass The molar mass (M) of any atom, molecule or compound is the mass (in grams) of one mole of that substance. – The molar mass in grams is numerically equal to the atomic mass or molecular mass expressed in u (or amu). For any element, the atomic mass and molar mass are numerically equal The mole concept and molar mass Molar mass is the mass of 1 mole of amu of an atom/molecule in grams. Molar mass of 12C: = 12g per mol of Carbon or 1mol per 12g of Carbon 12𝑔 𝐶 𝑚𝑜𝑙 𝐶 = or 𝑚𝑜𝑙 𝐶 12𝑔 𝐶 The mole concept and molar mass Molar mass is the mass of 1 mole of amu of an atom/molecule in grams. Molar mass of NH4: 14𝑔 𝑵 1𝑔 𝑯 18𝑔 𝑵𝑯4 𝑚𝑜𝑙 𝑵 + ( )4 = 𝑚𝑜𝑙𝑵𝑯4 𝑚𝑜𝑙 𝑯 2.3 Interconversion between mass, moles and elementary entities 1 𝑚𝑜𝑙 6.022 𝑥 1023 # 𝑜𝑓 𝑎𝑡𝑜𝑚𝑠 𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠, 𝑔 1 𝑚𝑜𝑙 1 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒 Atoms in a Grams Moles Particles particle 𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠, 𝑔 1 𝑚𝑜𝑙 1 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒 1 𝑚𝑜𝑙 6.022 𝑥 1023 # 𝑜𝑓 𝑎𝑡𝑜𝑚𝑠 Multiply by the value indicated by the arrow to convert 2.3 Interconversion between mass, moles and elementary entities Sample Problem 1 What is the molar mass CH4? 12𝑔 𝑪 1𝑔 𝑯 16𝑔 𝑪𝑯4 𝑚𝑜𝑙 𝑪 + ( )4 = 𝑚𝑜𝑙𝑪𝑯4 𝑚𝑜𝑙 𝑯 What is the mass of 0.25 mol CH4? 2.3 Interconversion between mass, moles and elementary entities 1 𝑚𝑜𝑙 6.022 𝑥 1023 # 𝑜𝑓 𝑎𝑡𝑜𝑚𝑠 𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠, 𝑔 1 𝑚𝑜𝑙 1 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒 Atoms in a Grams Moles Particles particle 𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠, 𝑔 1 𝑚𝑜𝑙 1 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒 1 𝑚𝑜𝑙 6.022 𝑥 1023 # 𝑜𝑓 𝑎𝑡𝑜𝑚𝑠 Multiply by the value indicated by the arrow to convert 2.3 Interconversion between mass, moles and elementary entities Sample Problem 1 What is the molar mass CH4? 12𝑔 𝑪 1𝑔 𝑯 16𝑔 𝑪𝑯4 𝑚𝑜𝑙 𝑪 + ( )4 = 𝑚𝑜𝑙𝑪𝑯4 𝑚𝑜𝑙 𝑯 What is the mass of 0.25 mol CH4? 0.25 mol 𝑪𝑯4 16𝑔 𝑪𝑯4 = 4.0g 𝑪𝑯4 𝑚𝑜𝑙𝑪𝑯4 2.3 Interconversion between mass, moles and elementary entities Problem 2 (Seat Work) How many molecules are there in 0.5 mol CH4? 2.3 Interconversion between mass, moles and elementary entities 1 𝑚𝑜𝑙 6.022 𝑥 1023 # 𝑜𝑓 𝑎𝑡𝑜𝑚𝑠 𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠, 𝑔 1 𝑚𝑜𝑙 1 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒 Atoms in a Grams Moles Particles particle 𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠, 𝑔 1 𝑚𝑜𝑙 1 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒 1 𝑚𝑜𝑙 6.022 𝑥 1023 # 𝑜𝑓 𝑎𝑡𝑜𝑚𝑠 Multiply by the value indicated by the arrow to convert 2.3 Interconversion between mass, moles and elementary entities Problem 2 (Seat Work) How many molecules are there in 0.5 mol CH4? What is the mass of 7.5275×1023 molecules of CH4? 2.3 Interconversion between mass, moles and elementary entities 1 𝑚𝑜𝑙 6.022 𝑥 1023 # 𝑜𝑓 𝑎𝑡𝑜𝑚𝑠 𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠, 𝑔 1 𝑚𝑜𝑙 1 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒 Atoms in a Grams Moles Particles particle 𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠, 𝑔 1 𝑚𝑜𝑙 1 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒 1 𝑚𝑜𝑙 6.022 𝑥 1023 # 𝑜𝑓 𝑎𝑡𝑜𝑚𝑠 Multiply by the value indicated by the arrow to convert 2.3 Interconversion between mass, moles and elementary entities Problem 2 (Seat Work) How many molecules are there in 0.5 mol CH4? 0.5 mol 𝑪𝑯4 6.022 x 1023 𝒎𝒐𝒍𝒆𝒄𝒖𝒍𝒆𝒔 𝑪𝑯4 3.0 x 1023 molecules 𝑪𝑯4 = 𝑚𝑜𝑙 𝑪𝑯4 What is the mass of 7.5275×1023 molecules of CH4? 2.3 Interconversion between mass, moles and elementary entities Problem 2 (Seat Work) How many molecules are there in 0.5 mol CH4? 0.5 mol 𝑪𝑯4 6.022 x 1023 𝒎𝒐𝒍𝒆𝒄𝒖𝒍𝒆𝒔 𝑪𝑯4 3.0 x 1023 molecules 𝑪𝑯4 = 𝑚𝑜𝑙 𝑪𝑯4 What is the mass of 7.5275×1023 molecules of CH4? 𝑚𝑜𝑙 𝑪𝑯4 16𝑔 𝑪𝑯4 7.5275×1023 molecules 𝑪𝑯4 = 20.0g 𝑪𝑯4 6.022 x 1023 𝒎𝒐𝒍𝒆𝒄𝒖𝒍𝒆𝒔 𝑪𝑯4 𝑚𝑜𝑙𝑪𝑯4 2.3 Interconversion between mass, moles and elementary entities Problem 2 (Continuation) How many hydrogen atoms are there in 15.0g of CH4? 2.3 Interconversion between mass, moles and elementary entities 1 𝑚𝑜𝑙 6.022 𝑥 1023 # 𝑜𝑓 𝑎𝑡𝑜𝑚𝑠 𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠, 𝑔 1 𝑚𝑜𝑙 1 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒 Atoms in a Grams Moles Particles particle 𝑀𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠, 𝑔 1 𝑚𝑜𝑙 1 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒 1 𝑚𝑜𝑙 6.022 𝑥 1023 # 𝑜𝑓 𝑎𝑡𝑜𝑚𝑠 Multiply by the value indicated by the arrow to convert 2.3 Interconversion between mass, moles and elementary entities Problem 2 (Continuation) How many hydrogen atoms are there in 15.0g of CH4? 𝑚𝑜𝑙𝑪𝑯4 6.022 x 1023 𝒎𝒐𝒍𝒆𝒄𝒖𝒍𝒆𝑪𝑯4 4 𝑎𝑡𝑜𝑚𝑠𝑯 15.0 g 𝑪𝑯4 16𝑔𝑪𝑯4 𝑚𝑜𝑙 𝑪𝑯4 1𝒎𝒐𝒍𝒆𝒄𝒖𝒍𝒆𝑪𝑯4 = 2.26 x 1024 atoms 𝑯 2.3 Interconversion between mass, moles and elementary entities Problem 3 (Assignment) Chloral hydrate (C2H3Cl3O2) is a drug formerly used as a sedative and hypnotic. a) Calculate the molar mass of chloral hydrate. b) What amount (moles) of C2H3Cl3O2 are in 500.0 g chloral hydrate? c) What is the mass in grams of 2.0 x 1022 molecules of chloral hydrate? 2.4 Percent composition of compounds #𝑎𝑡𝑜𝑚𝑠 𝑥 𝑎𝑚𝑢 𝑎𝑡𝑜𝑚 % atom = 𝑎𝑚𝑢 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒 x 100 2.4 Percent composition of compounds #𝑎𝑡𝑜𝑚𝑠 𝑥 𝑎𝑚𝑢 𝑎𝑡𝑜𝑚 % atom = x 100 𝑎𝑚𝑢 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒 Example: the percentage by mass of C in ethane, C2H6, is: 2 𝑥 12.01 𝑎𝑚𝑢 %C= x 100 30.068 𝑎𝑚𝑢 24.02 𝑎𝑚𝑢 = x 100 30.068 𝑎𝑚𝑢 = 79.89% Chemical Equation  a shorthand notation to describe a chemical reaction  Just like a chemical formula, a chemical equation expresses quantitative relations Subscripts tell the number of atoms of each element in a molecule Coefficients tell the number of molecules 2 H2O Chemical Equation 2.5 Common Chemical Reactions 1. Synthesis: Two or more reactants combine to produce ONE compound 𝑨 + 𝑩 𝑨𝑩 2. Decomposition: ONE compound is broken down to produce two or more compounds 𝑨𝑩 𝑨 + 𝑩 2.5 Common Chemical Reactions 3. Single Displacement: Element + compound produces element + compound 𝑨 + 𝑩𝑪 𝑨𝑪 + 𝑩 4. Double Displacement : Two compounds produce two different compounds 𝑨𝑩 + 𝑪𝑫 𝑨𝑫 + 𝑪𝑩 2.5 Common Chemical Reactions 5. Combustion: Burning – carbon compound combines with oxygen to release heat/energy 𝐶𝑜𝑚𝑝𝑜𝑢𝑛𝑑 + 𝑶𝟐 𝑯𝟐𝑶 + 𝑪𝑶𝟐 Always the product of complete combustion 2.5 Common Chemical Reactions Sample Problem 1 Identify the type of chemical reaction for the ff equations: a) ZnCl2 + Mg → MgCl2 + Zn b) BaCl2 + Na2SO4 → BaSO4 + 2NaCl c) CO2 + H2O → H2CO3 d) CaCO3 → CaO + CO2 2.6 Balancing Chemical Reactions 1. Write the correct formula for each substance 𝐻2 + 𝐶𝑙2 𝐻𝐶𝑙 2.6 Balancing Chemical Reactions 2. Add coefficients so the number of atoms of each element are the same on both sides of the equation 𝐻2 + 𝐶𝑙2 2𝐻𝐶𝑙 2 𝑎𝑡𝑜𝑚𝑠 𝑜𝑓 𝐶𝑙 2 𝑎𝑡𝑜𝑚𝑠 𝑜𝑓 𝐶𝑙 2 𝑎𝑡𝑜𝑚𝑠 𝑜𝑓 𝐻 2 𝑎𝑡𝑜𝑚𝑠 𝑜𝑓 𝐻 𝐵𝑜𝑡ℎ 𝑠𝑖𝑑𝑒 𝑎𝑟𝑒 𝑛𝑜𝑤 𝑏𝑎𝑙𝑎𝑛𝑐𝑒𝑑 2.6 Balancing Chemical Reactions 2. Add the states of reactants and products on both sides of the equation 𝐻2(𝑔) + 𝐶𝑙2(𝑔) 2𝐻𝐶𝑙 (𝑎𝑞) Combustion Analysis  The process of burning, the combination of a substance with oxygen to produce a flame  When an organic compound burns in oxygen, the carbon reacts with oxygen to form CO2, and the hydrogen forms water, H2O 2.6 Balancing Chemical Reactions 𝐵𝑎𝑙𝑎𝑛𝑐𝑖𝑛𝑔 𝐶𝑜𝑚𝑏𝑢𝑠𝑡𝑖𝑜𝑛 𝑅𝑒𝑎𝑐𝑡𝑖𝑜𝑛𝑠 1 Assume one molecule of the most complicated substance 𝐶5𝐻12 + 𝑂2 CO2 + 𝐻2𝑂 2.6 Balancing Chemical Reactions 𝐵𝑎𝑙𝑎𝑛𝑐𝑖𝑛𝑔 𝐶𝑜𝑚𝑏𝑢𝑠𝑡𝑖𝑜𝑛 𝑅𝑒𝑎𝑐𝑡𝑖𝑜𝑛𝑠 2 Adjust the coefficient of CO2 to balance C 𝐶5𝐻12 + 𝑂2 5CO2 + 𝐻2𝑂 2.6 Balancing Chemical Reactions 𝐵𝑎𝑙𝑎𝑛𝑐𝑖𝑛𝑔 𝐶𝑜𝑚𝑏𝑢𝑠𝑡𝑖𝑜𝑛 𝑅𝑒𝑎𝑐𝑡𝑖𝑜𝑛𝑠 3 Adjust the coefficient of H2O to balance H 𝐶5𝐻12 + 𝑂2 5CO2 + 6 𝐻2𝑂 2.6 Balancing Chemical Reactions 𝐵𝑎𝑙𝑎𝑛𝑐𝑖𝑛𝑔 𝐶𝑜𝑚𝑏𝑢𝑠𝑡𝑖𝑜𝑛 𝑅𝑒𝑎𝑐𝑡𝑖𝑜𝑛𝑠 5 Check the balance by counting the number of atoms of each element. 𝐶5𝐻12 + 8 𝑂2 5CO2 + 6 𝐻2𝑂 C=5 C=5 H = 12 H = 12 O = 16 O = 16 2.6 Balancing Chemical Reactions Seat Work Quantum Theory and the Electronic Structure of Atoms Chapter 3  Lecture Outline 3.1. 3.2 3.4 The nature of light Quantum Mechanics Atomic Orbitals 3.1.1 The Wave Nature of Light 3.1.2 The Particle Nature of Light 3.3 Quantum Numbers 3.5 3.6 3.7 Electron Configuration Condensed Electron Valence Electrons 3.5.1 Aufbau Principle Configuration 3.5.2 Pauli’s Exclusion Principle 3.5.3 Hund’s Rule  Energy vs. Matter “Matter and Energy were distinct.” Energy:  in the form of light is described as a wave  massless and delocalized Matter:  composed of particles  has mass and can be located anywhere in space  Energy vs. Matter But it turns out, LIGHT behaves both as a particle and a wave.  Wave Nature of Light Visible Light is a type of electromagnetic radiation. It has wavelike properties which means that it is characterized by wavelength, frequency and speed.  Energy vs. Matter Frequency – the time it takes for a wave to repeat a cycle per second. Units can be 1 Hz, s-1, or. Symbol is 𝑠 v(nu). Wavelength – the distance over which a wave’s shape repeats. Given by the symbol, λ Electromagnetic radiation also consists of light that is invisible such us UV, infrared, gamma rays etc  Energy vs. Matter To understand current atomic theory, you need to know about electromagnetic radiation (also called electromagnetic energy, or radiant energy). Particle Nature of Light Particle Nature of Light  Particle Nature of Light The Photoelectric Effect - electrons are emitted from a surface of a metal when light strikes it - minimum amount of energy is needed to remove an e- - Einstein assumed that the light is like a “particle of energy” called a photon. Particle Nature of Light The type of electromagnetic wave the photon is will depend on its wavelength, λ.  Quantized Energy State An electron can radiate or absorb energy as radiation only in limited amounts or bundles called quanta This means that electrons can move only on specific energy states (n=1,2, only integers, Z)  Particle Nature of Light Atomic Spectra Bohr’s Theory of the Hydrogen Atoms - the energy of electron in an H atom is quantized (line spectrum)  Particle Nature of Light Atomic Spectra Bohr’s Theory of the Hydrogen Atoms the energy of electron in an H atom is quantized (line spectrum) n is called the principal quantum number.  Quantum Mechanics  Currently accepted atomic model. Location of electrons can only be estimated mathematically. “The location of an e- is not a point but a region of space called orbitals.”  Quantum Numbers An electron’s location in an atom are specified by a set of numbers called quantum numbers. Can be thought of as an electron’s address. The 4 Quantum Numbers: Principal QN, n Azimuthal (Secondary) QN, l Magnetic QN, ml Spin QN, ms  Quantum Numbers Principal QN, n related to the size and energy of the orbital has integral values: 1, 2, 3,… increasing n means larger orbital and greater energy corresponds to the shell Quantum Numbers  Quantum Numbers Azimuthal (Secondary) QN, l related to the shape of the orbital has values from 0 to n – 1 with assigned letters a.k.a Angular Momentum QN corresponds to the subshell Quantum Numbers l=0, s subshell (sphere) 1 orbital l=1, p subshell (dumbbell) 3 orbitals l=2, d subshell l=3, f subshell 5 orbitals 7 orbitals  Quantum Numbers Magnetic QN, ml determines the orientation of orbital in space numbers are limited by l ml = -l to l Ex: for l = 1, ml = -1,0,1 for l = 2, ml = -2,-1,0,1,2 Each ml value corresponds to one orbital in the subshell  Quantum Numbers Spin QN, ms determines the rotation of electron around itself value is either +½ (↑) or –½ (↓) Quantum Numbers Quantum Numbers  Quantum Numbers Energy Levels Energy states and atomic orbitals of an atom is described with specific terms and are associated with one or more QN. 1. Level - atom’s energy levels or shells = n 2. Sublevel - levels are divided into sublevels or subshells: ▪ l = 0 = s subshell ▪ l = 2 = d subshell ▪ l = 1 = p subshell ▪ l = 3 = f subshell  Quantum Numbers Energy Levels 3. Orbital - each combination of QN tells the size (energy), shape, and spatial orientation of one of an atom’s orbital Example ▪ 2s subshell: n = 2, l = 0, ml = 0; - 1 value of ml means 1 orbital for this subshell ▪ 3p subshell: n = 3, l = 1, ml = -1,0,1; - 3 values of ml means 3 orbitals for this subshell Quantum Numbers Energy Levels Quantum Numbers Energy Levels  Electron Configurations Each orbital can have up to two electrons The s subshell has one orbital, up to 2 electrons The p subshell has three orbitals, up to 6 electrons The d subshell has five orbitals, up to 10 electrons The f subshell has seven orbitals, up to 14 electrons Quantum Numbers Quantum Numbers  Electron Configurations s p d f 1 1s Construction of the 2 2s 2p Electron Configurations of 3 3s 3p 3d an atom has 3 rules 4 4s 4p 4d 4f 5 5s 5p 5d 5f 6 6s 6p 6d 6f 7 7s 7p 7d 7f  Electron Configurations Aufbau Principle s p d f electrons fill orbitals starting with the 1 1s lowest energy levels 2 2s 2p Energy Ranking: 3 3s 3p 3d 1s < 2s < 2p < 3s < 3p < 4s 4 4s 4p 4d 4f 3d < 4p < 5s < 4d < 5p < 6s 5 5s 5p 5d 5f 4f < 5d < 6p < 7s < 5f < 6d 6 6s 6p 6d 6f 7p < 6f < 7d < 7f 7 7s 7p 7d 7f  Electron Configurations 1. Aufbau Principle s p d f the energy of subshells in the same shell 1 1s is in the order of: 2 2s 2p ns < np < nd < nf 3 3s 3p 3d Orbitals in the same subshell are degenerate 4 4s 4p 4d 4f (same energy) 5 5s 5p 5d 5f 6 6s 6p 6d 6f 7 7s 7p 7d 7f  Electron Configurations s p d f 2. Pauli’s Exclusion Principle only 2 electrons of opposite spins per 1 1s orbital no two electrons has the same set of 2 2s 2p quantum numbers 3 3s 3p 3d Example For 1s orbital: 4 4s 4p 4d 4f One electron: n = 1, l = 0, ml = 0, ms = +½ 5 5s 5p 5d 5f Another electron: n = 1, l = 0, ml = 0, ms = -½ 6 6s 6p 6d 6f 7 7s 7p 7d 7f 1s  Electron Configurations s p d f 3. Hund’s Rule for degenerate orbitals, electrons fill each 1 1s orbital singly with parallel spins before any orbital gets a second electron 2 2s 2p 3 3s 3p 3d Example: Write the orbital diagram of C= 4 4s 4p 4d 4f 5 5s 5p 5d 5f 6 6s 6p 6d 6f 7 7s 7p 7d 7f Electron configuration of Carbon atom  Electron Configurations s p d f 3. Hund’s Rule for degenerate orbitals, electrons fill each 1 1s orbital singly with parallel spins before any orbital gets a second electron 2 2s 2p 3 3s 3p 3d Example: Write the orbital diagram of C= 1s2 2s2 2p2 4 4s 4p 4d 4f 5 5s 5p 5d 5f 6 6s 6p 6d 6f 1s 2s 2p 7 7s 7p 7d 7f 6 e- (electrons) of Carbon atom  Electron Configurations s p d f 3. Hund’s Rule for degenerate orbitals, electrons fill each 1 1s orbital singly with parallel spins before any orbital gets a second electron 2 2s 2p 3 3s 3p 3d Example: Write the orbital diagram of C = 1s2 2s2 2p2 : 4 4s 4p 4d 4f 5 5s 5p 5d 5f 6 6s 6p 6d 6f 1s 2s 2p 7 7s 7p 7d 7f Write the electron configuration of the following: 1. Kr 2. Cl- 3. Mg+2 4. Cu 5. Cr  Condensed Electron Configurations Core electrons – electrons located in completely filled subshells Noble gas core notation – noble gas electron configurations are used to abbreviate the core electrons of all elements. Noble Gases is in G8A or G18  Condensed Electron Configurations Core electrons – electrons located in completely filled subshells Noble gas core notation – noble gas electron configurations are used to abbreviate the core electrons of all elements. Noble Gases is in G8A or G18  Condensed Electron Configurations The underlined subshells contain electrons in the outermost shell (highest n), also called the valence electrons Condensed Electron Configurations Condensed Electron Configurations  Unusual Electron Configurations Most stable arrangement of electrons in subshells are fully-filled or half-filled Electronic Configuration and the Periodic Table Map for outermost subshells of each element  Unit 4 Chem 16 | Principles of Chemistry Periodic Trends of the Elements  Lecture Outline 4.1 4.2 4.3 Development of the Organization of the Factors that Periodic Table Periodic Table Determine the Periodic Trends 4.3.1 4.3.2 4.4 Energy Shell/Level Effective Nuclear Periodic Trends in Atomic Radius, Ionic Radius, Charge Ionization Energy, Electron Affinity Electronegativity and Metallic Character 4.1 Development of the Periodic Table Development of the Periodic Table Dmitri Mendeleev and Lothar Meyer arranged the elements in order of increasing atomic weight Henry Moseley rearranged it based on increasing atomic number so arrangement is consistent based on the elements’ properties 4.2 Organization of the Periodic Table  Organization of the Periodic Table Period/series – elements in a row Family/group – elements in a column Classification of Elements 1. Based on properties metals, nonmetals, metalloids 2. Based on electronic configuration representative/main group elements, transition elements, inner transition elements, noble/inert gases Organization of the Periodic Table  Metals and nonmetals Metals Nonmetals Shiny Non-shiny Good conductor of heat Poor conductor of heat Good conductor of electricity VS Poor conductor of Malleable electricity (easy to mould) Not malleable Ductile (hard to mould) (can bend without breaking) Ductile All solid at room temp (brittle) (except Mercury, Hg)  The periodic table of elements 1 2 H He 3 4 5 6 7 8 9 10 Li Be B C N O F Ne Recently 11 12 Metal Metalloid Nonmetal 13 14 15 16 17 18 discovered Na Mg Al Si P S Cl Ar 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 55 56 71 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 Cs Ba Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 87 88 103 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 Fr Ra Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og Lanthanides 57 58 59 60 61 62 63 64 65 66 67 68 69 70 La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb 89 90 91 92 93 94 95 96 97 98 99 100 101 102 Actinides Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No 4.3 Factors that Determine the Periodic Trends  Organization of the Periodic Table Based on electronic configuration 1. Representative / main group elements elements in which the last electron added enters an s or p orbital in the outermost shell but in which this shell is incomplete (Group 1A – 7A) 2. Transition elements elements that have filled or partially-filled inner d subshell (Group 1B – 8B)  Organization of the Periodic Table Based on electronic configuration 3. Inner transition elements (f- block) elements that have filled or partially-filled inner f subshell (lanthanides and actinides) 4. Noble / inert gas have filled valence shell (Group 8A) and are thus very stable since they have closed shells (ns2 np6)  Effective Nuclear Charge 1. Valence electrons electrons on the outermost shell of an atom (highest n). Often corresponds with group number. Mg: 1s2 2s2 2p6 3s2 : 2 valence electrons F: 1s2 2s2 2p5 : 7 valence electrons 2. Effective nuclear charge charge or attraction experienced by an electron on a many-electron atom due to the protons inside the nucleus  Shell/Level 1. The higher the principal QN, n, the further the valence (outer shell) electrons are from the nucleus. Thus, the effective nuclear charge of the valence electrons decreases the higher the n. 2. Rank the effective nuclear charge of the val e- of Na, Li and K in decreasing order: Li>Na>K 3. Down a group, the effective nuclear charge on valence electrons decreases.  Effective Nuclear Charge Across a Period 1. Across a period, the principal QN, n, does not increase but the number of protons increase, causing stronger attraction to the valence electrons and thus stronger nuclear charge 2. Rank the effective nuclear charge of the val e- of C, N, O and F in increasing order: C Mg2+ > Al3+  Periodic Trends: Ionization Energy 1. First ionization energy, I1 – amount of energy required to remove an electron from a gaseous atom Na(g) Na+(g) + e- 1. The larger ionization energy, the more difficult it is to remove the electron 2. The higher the effective nuclear charge, the higher the first Ionization energy Periodic Trends: Ionization Energy  Periodic Trends: Ionization Energy Second ionization energy, I2 energy required to remove an electron from a gaseous ion: Na+(g) Na2+(g) + e-  Periodic Trends: Ionization Energy It is harder to remove inner shell electrons than valence (outer shell) electrons as shown by the high jump on ionization energies on the inner shell electrons  Periodic Trends: Electron Affinity Electron affinity energy change when a gaseous atom gains an electron to form a gaseous ion: The stronger the effective nuclear charge, the stronger the electron affinity. Cl(g) + e- Cl-(g) Periodic Trends: Electron Affinity  Periodic Trends: Electron Affinity Electronegativity the ability of a bonded atom to attract electrons to itself Stronger effective nuclear charge causes stronger electronegativity 1. Within a group - EN decreases from top to bottom 2. Within a period - EN increases across a period  Periodic Trends: Metallic Character Metallic character the properties of metals Metals have low ionization energies and tend to form cations  Periodic Trends: Non-metallicity Within a group - nonmetallic character decreases from top to bottom Within a period - nonmetallic character increases from left to right  Properties of Metals and Nonmetals 1.Metals – low IE, low EA o Easily ionize (tend to form cation) and hardly accepts an added electron 2.Nonmetals – high IE, high EA o Does not easily ionize but easily accepts an added electron (tend to form anion)