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Unit 1 Section 4 Atomic Structure Chemical Foundations: Elements, Atoms, and Ions Textbook - Chapter 4 & 11 Table of Contents 4.1 4.2 4.4 4.5 4.6 4.7 11.1 11.2 11.3 11.4 11.5 11.6 11.7 11.8 11.9 11.10 The Elements Symbols for the Elements Formulas of Compounds The Structure of the Atom Introduction to the Modern Concept of Atomic Structure Isotopes (Calculating Average Atomic Mass) Rutherford’s Atom Electromagnetic Radiation and Energy Emission The Energy Levels of Hydrogen The Bohr Model of the Atom The Wave Mechanical Model The Orbitals The Wave Mechanical Model: Further Development Electron Arrangements Electron Configurations 2 Chapter 4 Table of Contents Objectives • To learn how a formula describes a compound’s composition. • To study the atom’s structure 3 Section 4.1 The Elements • • Substances that cannot be broken down by simple chemical means 118 known: 88 found in nature, others are man made. Return to TOC 4 Section 4.1 The Elements Element Abundance Most abundant elements in the universe: hydrogen H 60% helium He 37% Most abundant elements in the entire earth: iron Fe 35% silicon oxygen O 30% Si Most abundant element in earth’s crust: oxygen O 49.2% aluminum silicon Si 25.7% Al 15% 7.50% Return to TOC 5 Section 4.1 The Elements Element Abundance Naturally radioactive elements: uranium U radium Ra radon Rn Polonium Po Elements most abundant in human body: oxygen O nitrogen N carbon C phosphorus P hydrogen H calcium Ca Return to TOC 6 Section 4.1 The Elements Return to TOC 7 Section 4.1 & Section 4.9 The Elements Diatomic Elements Diatomic elements – elements that occur naturally paired as two atoms per molecule - 7 diatomic elements H2 Cl2 F2 N2 Br2 O2 I2 GEN-U-INE Return to TOC 8 Section 4.2 Symbols for the Elements • • • Each element has a unique one- or two-letter symbol. First letter is always capitalized and the second is not. The symbol usually consists of the first one or two letters of the element’s name. • Examples: Oxygen Krypton O Kr Sometimes the symbol is taken from the element’s original Latin or Greek name. Examples: Gold Au Lead Pb aurum plumbum Return to TOC 9 Section 4.2 Symbols for the Elements Names and Symbols of the Most Common Elements Return to TOC 10 Section 4.4 Formulas of Compounds Chemical Formulas Describe Compounds • Compound – distinct substance that is composed of the atoms of two or more elements and always contains exactly the same relative masses of those elements. • Chemical Formulas – expresses the types of atoms and the number of each type in each unit (molecule) of a given compound. Return to TOC 11 Section 4.4 Formulas of Compounds Rules for Writing Formulas 1. Symbol Tells which atoms are present in compounds 2. Subscript The number of each type of atom 3. Coefficient (number in front) tells the number of molecules Coefficient 4 Return to TOC 12 Section 4.3 Atomic Dalton’sTheory Atomic Theory Scientists of the 18th century learned that: 1. Most natural materials are mixtures of pure substances. 2. Pure substances are either elements or combinations of elements called compounds. 3. A given compound always contains the same proportions (by mass) of the elements. Law of Constant Composition Return to TOC 13 Section 4.3 Atomic Dalton’sTheory Atomic Theory Law of Constant Composition • A given compound always has the same composition, regardless of where it comes from. Water always contains 8 g of oxygen for every 1 g of hydrogen. Return to TOC 14 Section 4.5 4.3 The Structure of Theory the Atom Dalton’s Atomic Objectives • Recognize that science is a progressive endeavor that reevaluates and extends what is already known. • Compare and contrast historical models of the atom. Return to TOC Section 4.5 4.6 The Structure of Introduction tothe theAtom Modern Concept of Atomic Structure The atom contains: • Electrons – found outside the nucleus; negatively charged (-1) • Protons – found in the nucleus; positive charge (+1) • Neutrons – found in the nucleus; no charge Return to TOC 16 Section 4.5 Section 4.1 The of the Atom TheStructure Elements William Thomson (Plum Pudding Model) • Plum pudding model A physicist, Thompson, believed an atom was filled with positively charged material and the electrons were evenly distributed throughout. - This model of the atom turned out to be short-lived, due to the work of Ernest Rutherford (1871–1937). Return to TOC 17 4.2 Section 4.5 4.1 The Elements Ernest Rutherford’s Portrait Return to TOC Section 4.5 4.1 Ernest Rutherford The Elements • Discovered the nucleus (1911) • Stated protons were inside of the nucleus • Gold-Foil Experiment Radioactive alpha particles “shot” through gold foil - A majority of particles passed straight through foil - A small fraction of particles bounced off gold foil at large angles or bounced straight back Return to TOC Section 4.5 Section 4.5 4.1 The Elements Rutherford’s alpha particle scattering experiment. Return to TOC 5.5 20 Section 4.5 4.1 Rutherford’s Atomic Model The Elements Rutherford discovered that: • The atom contains a tiny dense center called the nucleus – the volume is about 1/10 trillionth the volume of the atom • The nucleus is positively charged – the amount of positive charge of the nucleus balances the negative charge of the electrons • The electrons move around in the empty space of the atom surrounding the nucleus 21 Return to TOC Copyright©2004 by Houghton Section 4.5 4.1 Rutherford’s possible model The Elements Return to TOC 22 Section 4.6 Introduction to the Modern Concept of Atomic Structure Why do different atoms have different chemical properties? • • • The chemistry of an atom arises from its electrons. Electrons are the parts of atoms that “intermingle” when atoms combine to form molecules. It is the number of electrons that really determines chemical behavior. Return to TOC 23 Section 4.6 Introduction to the Modern Concept of Atomic Structure Rutherford Atomic Model Rutherford’s model turned out to be incomplete. • The Rutherford atomic model had to be revised in order to explain the chemical properties of elements. Return to TOC Section 4.6 Introduction to the Modern Concept of Atomic Structure The Modern Atom • We know atoms are composed of three main pieces - protons, neutrons and electrons • The nucleus contains protons and neutrons • The nucleus is only about 10-13 cm in diameter • The electrons move outside the nucleus with an average distance of about 10-8 cm Return to TOC 25 Section 4.6 Introduction to the Modern Concept of Atomic Structure Return to TOC 26 Section 4.3 Creating Dalton’sModels Atomic Theory 1) Create a drawing that illustrates Thomson’s view of an atom. 2) Based on what you learned about atoms in other science classes, create a diagram that represents the structure of an atom. 3) Explain how your two drawings are different. Return to TOC Section 4.7 Isotopes Objective • To learn about atomic number, mass number, and isotopes. Return to TOC 28 Section 4.7 Isotopes •Why are atoms with different numbers of neutrons still considered to be the same element? •Despite differences in the number of neutrons, isotopes are chemically alike. They have identical numbers of electrons, which determine chemical behavior. Return to TOC Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Section 4.7 Isotopes Atomic number and Mass Atomic number – the number of protons in the nucleus Mass number = protons + neutrons •The mass listed in the periodic table is the average atomic mass •It is a weighted average of the atomic masses of naturally occurring isotopes Mass Number Atomic Number A ZX Element Symbol Return to TOC 30 Section 4.7 Isotopes Atomic number and Mass • Remember that atoms are electrically neutral. In an atom, protons = electrons • Protons, neutrons, and electrons can be calculated from atomic number and mass number. How many protons, electrons, and neutrons are in each atom? boron and sodium Return to TOC 31 Section 4.7 Isotopes Isotopes • • • 1 1H Isotopes - atoms with the same number of protons but different numbers of neutrons. Show almost identical chemical properties; chemistry of atom is due to its electrons. In nature most elements contain mixtures of isotopes. 2 1H (D) 3 1H (T) 235 92 U 238 92 U Return to TOC 32 Section 4.7 Isotopes A Z • • • X X = the symbol of the element Z = the atomic number (# of protons) A = the mass number or atomic mass unit (amu) (# of protons and neutrons) A – Z = n (number of neutrons) Return to TOC 33 Section 4.7 Isotopes Two Isotopes of Sodium A – Z = n (number of neutrons) Return to TOC 34 Section 4.7 Isotopes Isotopes – An Example 14 6 • • • C C = the symbol for carbon 6 = the atomic number (6 protons) 14 = the mass number (6 protons and 8 neutrons) 12 6 C • C = the symbol for carbon • 6 = the atomic number (6 protons) • 12 = the mass number (6 protons and 6 neutrons) A – Z = n (number of neutrons) Return to TOC 35 Section 4.7 Isotopes Exercise A certain isotope X contains 23 protons and 28 neutrons. • What is the mass number of this isotope? • Identify the element. Mass Number = 51 Vanadium Return to TOC 36 Section 4.7 Isotopes Atomic Mass Unit – How is it calculated? • The value shown in the periodic table is the average atomic mass It is a weighted average of the atomic masses of naturally occurring isotopes Return to TOC 37 Section 4.7 Isotopes Atomic Mass Unit – How is it calculated? • For example: Chlorine has two isotopes Chlorine-35 and Chlorine-37 The abundance is: Cl-35 has an amu of 34.9689 with an abundance of 75.771% Cl-37 has an amu of 36.9659 with an abundance of 24.229% Return to TOC 38 Section 4.7 Isotopes Calculating Atomic Mass Unit • Cl-35 has an amu of 34.9689 with an abundance of 75.771% • Cl-37 has an amu of 36.9659 with an abundance of 24.229% Return to TOC 39 Section 4.7 Isotopes Mass Spectroscopy • an analytical tool used for measuring the molecular mass of a sample Return to TOC 40 Section 4.7 Isotopes Mass Spectrum for the element Boron • The number of isotopes: The two peaks in the mass spectrum shows that there are 2 isotopes of boron – with relative isotopic masses of 10 and 11 • The abundance of the isotopes: Can be determined by the height of the peak. Return to TOC 41 Section 4.7 Isotopes Chapter 11 Modern Atomic Theory Copyright© by Houghton Mifflin Company. All rights reserved. Return to TOC Section 4.7 Electrons in Atoms Isotopes Rutherford’s model has some limitations It did not explain the chemical properties of the elements It did not address the electrons • What are electrons doing? – How are they arranged & how do they move? Return to TOC 43 Section 4.7 Electrons in Atoms Isotopes Figure 11.1: The Rutherford atom. Return to TOC 44 Section 4.7 Electrons in Atoms Isotopes • To understand the next development, we must understand some properties of light.. Return to TOC 45 Section 4.7 11.2 Electromagnetic Radiation and Energy Isotopes By the year 1900, there was enough experimental evidence to convince scientists that light consisted of waves. • The wavelength, represented by (the Greek letter lambda), is the distance between two wave peaks. Return to TOC 46 Section 4.7 11.2 Electromagnetic Radiation and Energy Isotopes • The frequency, represented by (the Greek letter nu), is the number of wave peaks that pass a certain point per unit of time. • The SI unit of waves per second is called the hertz (Hz). Return to TOC 47 Section 4.7 Isotopes • The frequency ( ) and wavelength ( ) of light are inversely proportional to each other. • As the wavelength increases, the frequency decreases. Return to TOC 48 Section 4.7 Isotopes Return to TOC Section 4.7 Isotopes Rutherfo https://www.youtube.com/watch?v=cfXzwh3KadE Return to TOC 50 Section 4.7 11.2 Electromagnetic Radiation and Energy Isotopes According to the wave model, light consists of electromagnetic waves. Electromagnetic radiation - a form of energy that exhibits wavelike behavior as it travels through space. – All electromagnetic radiation travels at the speed of light: c = 3.0 X108 m/s Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays. Return to TOC 51 Section 4.7 Isotopes Visible light of different wavelengths can be separated into a spectrum of colors. In the visible spectrum, red light has the longest wavelength and the lowest frequency. Violet light has the shortest wavelength and the highest frequency. Return to TOC 52 Section 4.7 Wave Description of Light Isotopes • Equation relating frequency and wavelength: c = c = speed of light = wavelength = frequency • c is constant, so is , so as frequency increases, wavelength decreases (inversely proportional). Return to TOC 53 Light as a Wave: Problems Isotopes Section 4.7 1) What is the frequency of light if its wavelength is 4.34 X 10-7 m? 2) What is the wavelength of a wave with a frequency of 1019 Hz (s-1)? Return to TOC 54 Section 4.7 Energy is Quantized Isotopes • In 1900 a physcist named Max Planck proposed that matter does not emit electromagnetic energy continuously. • Max Planck suggested that the object emits energy in small specific amounts called quanta. • Quantum - the minimum quantity of energy that can be lost or gained by an atom. Return to TOC 55 56 Section 4.7 Light as a Particle: Isotopes Planck’s Equation Energy (Joules) of a quantum of radiation Frequency (Hz) E = h Plank’s Constant = 6.626 x 10-34 J•s Video - http://study.com/academy/lesson/what-is-a-photon-definition-energy-wavelength.html Return to TOC Section 4.7 Einstein Isotopes • Next Albert Einstein suggested that the electromagnetic (em) spectrum is itself quantized. • Einstein proposed that em radiation can be viewed as photons • Photon - a particle of electromagnetic radiation carrying a quantum of energy – A photon is like an energy packet Return to TOC 58 Section 4.7 Electrons in Atoms Isotopes Ephoton = h 1. What is the energy of a photon with a frequency of 9x1014 Hz? 2. What is the frequency of a photon with an energy of 5x10-22 J? Return to TOC 59 Section 4.7 Wave-Particle Isotopes Duality •In 1905 Albert Einstein introduced the idea of that electromagnetic radiation has “wave-particle duality” Light has wavelike particles and Light can be thought of as a stream of particles where each particle carries a quantum of energy. Video - http://study.com/academy/lesson/what-is-a-photon-definition-energy-wavelength.html Return to TOC Section 4.7 11.4 Energy Levels Isotopes • • • Ground State: the lowest energy state of an atom. Excited State: when an atom contains excess energy (has higher potential energy). When an excited atom returns to ground state it gives off the energy it has gained as electromagnetic radiation. Example: Neon signs Return to TOC Section 4.7 Isotopes Absorption • An electron absorbs energy (photon) and moves from the ground state to an excited state. E 4 E 3 E 2 E 1 Return to TOC Section 4.7 What goes up…must come down! Isotopes Emission • When an electron in the excited state returns to the ground state it emits a photon. E 4 E 3 E 2 E = h =E -E photon 3 1 E 1 Return to TOC Section 4.7 Absorption Isotopes and Emission • Emitting photons creates light or electromagnetic radiation • Electromagnetic radiation in the visible light spectrum has color! • These photons have wavelengths that correspond to their color. VIDEO - http://study.com/academy/lesson/the-bohr-model-and-atomic-spectra.html Return to TOC Section 4.7 Isotopes In an actual atom… • Emission and absorption happen simultaneously and on different energy levels. • Emission produces a line-emission spectrum E 4 E 3 E 2 E 1 Return to TOC Section 4.7 Line Isotopes Emission Spectrum • Line-Emission Spectrum- a beam of light separated into a series of specific frequencies (and therefore specific wavelengths) of visible light. – produced when electrons fall back to ground state – the photons emitted in the fall give off specific patterns (colors) of light Return to TOC Section 4.7 Isotopes The Hydrogen-Atom Line Emission Spectrum Return to TOC Section 4.7 Isotopes The Hydrogen-Atom Line Emission Spectrum Return to TOC Section 4.7 11.5 The Bohr Model Isotopes In 1913, Niels Bohr develops a new atomic model Bohr stated that the electrons orbit the nucleus like the planets orbit the sun. Return to TOC 69 Section 4.7 Isotopes Each possible electron orbit in Bohr’s model has a fixed energy. The fixed energies an electron can have are called energy levels. Each energy level further from the nucleus is of greater energy Energy levels are regions where an electron is likely to be moving Return to TOC 70 Section 4.7 Isotopes The rungs on this ladder are like the energy levels in Bohr’s model. A person on a ladder cannot stand between the rungs. Similarly, the electrons in an atom cannot exist between energy levels. Return to TOC 71 Section 4.7 Isotopes The Rutherford model could not explain why elements that have been heated to higher temperatures give off different colors of light. The Bohr model explains how the energy levels of electrons in an atom change when the atom emits light. Return to TOC 72 Section 4.7 Bohr’s model Isotopes Return to TOC 73 Section 4.7 Bohr’s Model of the Atom Isotopes • Bohr’s explained emission spectra using the idea of fixed orbits. • This idea is close but not true… Return to TOC Section 4.7 11.6 The Wave (Quantum) Mechanical Isotopes Model Unfortunately, Bohr’s model did not apply to other atoms That led scientists to question his model They wondered why the electron had to be located in a precise orbit Return to TOC 75 Section 4.7 The Isotopes Wave (Quantum) Mechanical Model • That led to the Heisenburg uncertainty principle states that it’s impossible to determine the position and velocity of an electron at the same time Video clip - https://www.youtube.com/watch?v=H-AlfuvjPYM Return to TOC 76 The Wave (Quantum) Mechanical Isotopes Model Section 4.7 • Further developments led to the wave (quantum) mechanical model The wave (quantum) mechanical model describes mathematically the position of electrons in an atom Return to TOC 77 The Wave (Quantum) Mechanical Isotopes Model Section 4.7 Like the Bohr model, the quantum mechanical model of the atom restricts the energy of electrons to certain values. However, the quantum mechanical model does not specify an exact path the electron takes around the nucleus. Return to TOC 78 The Wave (Quantum) Mechanical Isotopes Model Section 4.7 • Erwin Schrodinger (1887-1961) developed the quantum mechanical model further It determined the allowed energies an electron can have Also, Schrodinger developed an equation to determine how likely it is to find an electron in a particular location around the nucleus of an atom. Return to TOC 79 The Wave (Quantum) Mechanical Isotopes Model Section 4.7 Return to TOC 80 The Wave (Quantum) Mechanical Isotopes Model Section 4.7 Rutherfo Return to TOC 81 Section 4.7 Isotopes Concept Check TRUE or FALSE In the modern atomic model, electrons are moving around the nucleus in a circular path. a) True b) False VIDEOCLIP - https://www.youtube.com/watch?v=H-AlfuvjPYM Return to TOC Copyright © Cengage Learning. All rights reserved 82 Section 4.7 11.7 The Orbitals Isotopes • The model explained that an electron exists in certain regions called orbitals. A 3D region around the nucleus that indicates the probable location of an electron Return to TOC 83 Section 4.7 Quantum Numbers Isotopes The s orbitals are spherical. The p orbitals are dumbbell shaped. Return to TOC 84 Section 4.7 11.7 The Orbitals - Quantum Numbers Isotopes • In order to specify the properties of atomic orbitals and electrons in orbitals, chemists use quantum numbers Principle quantum number (n) Angular momentum quantum number (l) Magnetic quantum number (m) Spin quantum number Return to TOC 85 Section 4.7 Hydrogen Isotopes Energy Levels • Hydrogen has discrete energy levels. Called principal energy levels Principal energy levels Labeled as n = 1, n=2, 3, 4, and so forth The energy level corresponds to the periods of the periodic table 1st energy level = period 1 Return to TOC Copyright © Cengage Learning. All rights reserved 86 Section 4.7 Sublevels Isotopes • Energy levels can be divided into sublevels s, p, d, f are sublevels • Orbitals exist in sublevels An orbital can be empty or it can contain one or two electrons Return to TOC 87 Section 4.7 Isotopes • Each principal energy level is divided into sublevels. Labeled with numbers and letters Indicate the shape of the orbital Return to TOC Copyright © Cengage Learning. All rights reserved 88 Section 4.7 Figure 11.18: Principal levels can be divided into sublevels Isotopes Return to TOC 89 Section 4.7 Quantum Numbers Isotopes The spin quantum number indicates the spin of the electron • It may be thought of as clockwise or counterclockwise. • A vertical arrow indicates an electron and its direction of spin ( or ) Return to TOC 90 Section 4.7 Quantum Numbers Isotopes Rutherfo Return to TOC 91 Section 4.7 11.9 and 11.10 Electron Arrangement Isotopes To describe the arrangement of the electrons in an atom we use electron configuration To describe spin we use orbital notation •Remember – an atom tends to assume the lowest energy configuration possible Return to TOC 92 Section 4.7 Electron Arrangement Isotopes To determine electron configuration, follow three simple rules 1.The Aufbau principle states that an electron occupies the lowest energy orbital that can receive it oWe fill lowest to highest energy Return to TOC 93 Section 4.7 Orbital Diagram Isotopes Rutherfo Return to TOC 94 Section 4.7 Electron Arrangment Isotopes 2. The Pauli Exclusion principle states that an atomic orbital can hold a maximum of two electrons, and those two electrons must have opposite spins. 3. Hunds Rule states that orbitals of the same energy must be occupied by one electron before it can be occupied by a second electron. Return to TOC 95 Section 4.7 Isotopes Hund’s Rule Orbital Diagrams Three electrons would occupy three orbitals of equal energy as follows. Electrons then occupy each orbital so that their spins are paired with the first electron in the orbital. Return to TOC Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Orbital Diagram and Electron Configuration Isotopes Look at the orbital filling diagram of the oxygen atom. Section 4.7 • An oxygen atom contains eight electrons. Electron Configurations of Selected Elements Element 1s 2s 2px 2py 2pz 3s Electron configuration H 1s1 He 1s2 Li 1s22s1 C 1s22s22p2 N 1s22s22p3 O 1s22s22p4 F 1s22s22p5 Ne 1s22s22p6 Na 1s22s22p63s1 Return to TOC Copyright © Pearson Education, Inc., or its affiliates. All Rights Section 4.7 Orbital Diagram and Electron Configuration Isotopes Return to TOC 98 Section 4.7 DO NOW Isotopes Draw the orbital diagrams for the following: • Nitrogen • Cobalt 1s 2s 2p 3s 3p 4s 3d Return to TOC 99 Section 4.7 11.10 Electron Configuration Isotopes Orbitals, Sublevels & Electrons • for a many electron atom, build-up the energy levels, filling each orbital in succession by energy 2 2 2 6 2 6 10 6 2 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d 6 2 14 10 6 2 14 10 < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p 10 6 - each s sublevel has 1 s orbital (can hold 2 e-) - each p sublevel has 3 p orbitals (can hold 6 e-) - each d sublevel has 5 d orbitals (can hold 10 e-) - each f sublevel has 7 f orbitals (can hold 14 e-) Return to TOC 100 Section 4.7 H Atom Isotopes • Electron configuration – electron arrangement 1s2 • Orbital diagram – orbital is a box grouped by sublevel containing arrow(s) to represent electrons Return to TOC Copyright © Cengage Learning. All rights reserved 101 Section 4.7 Isotopes Li Atom Orbital diagram Electron configuration Return to TOC Copyright © Cengage Learning. All rights reserved 102 Section 4.7 O Atom Isotopes • The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons in a particular set of degenerate (same energy) orbitals. Electron Configuration Orbital Diagram Return to TOC Copyright © Cengage Learning. All rights reserved 103 Section 4.7 Orbitals Isotopes • Parts of the periodic table corresponds to each orbital shape Groups 1 & 2 – s block Groups 13-18 – p block Groups 3-12 – d block Bottom two rows – f block Return to TOC 104 Section 4.7 Orbitals Isotopes Return to TOC 105 Section 4.7 Isotopes • The electron configurations in the sublevel last occupied for the first eighteen elements. Return to TOC Copyright © Cengage Learning. All rights reserved 106 Section 4.7 11. 10 Abbreviating Electron Configuration Isotopes To avoid writing the inner-level electrons, we often abbreviate the configurations. For example: for sodium 1s22s22p63s1 [Ne]3s1 for titanium 1s22s22p63s23p64s23d2 [Ar]4s23d2 Return to TOC 107 Section 4.7 Abbreviating Electron Configuration Isotopes Try abbreviating for: a) Silicon 1s22s22p63s23p2 b) Vanadium 1s22s22p63s23p64s23d3 c) Strontium 1s22s22p63s23p64s23d104p65s2 Return to TOC 108 Section 4.7 Isotopes Do Now What is the electron configuration for Germanium? How many electrons can the p sublevel hold? How many electrons can principal energy level 3 hold? Return to TOC 109 Section 4.7 Isotopes Do Now Write the noble gas configuration for the following elements. 1. Calcium [Ar] 3s2 2. Iodine [Kr] 5s24d105p5 Return to TOC 110 Section 4.7 Valence electrons Isotopes • Highest energy level for any atom is called the valence shell – electrons in the valence shell are called valence electrons Return to TOC 111 Section 4.7 Valence electrons Isotopes Example: • Carbon’s (with an electron configuration of 1s22s22p2) highest energy level is principal energy level 2. • There are 2 electrons from the 2s sublevel and 2 electrons from the 2p sublevel. • 2 + 2 = 4 4 valence electrons for carbon Return to TOC 112 Section 4.7 Valence electrons Isotopes Example: • Lithium’s (with an electron configuration of 1s22s1) highest energy level is principal energy level 2. • There is 1 electron from the 2s sublevel. • 1 valence electron for lithium How many valence electrons are there in Ca? I? Return to TOC 113