Chapter 5 and 6 Notes Electromagnetic Radiation and Light Models of the Atom  There were many different models over time — Dalton-billiard ball model (1803) — Thompson – plum-pudding model (1897) — Rutherford – __________ model of the atom (1911) — Bohr – uses quantized ________ of the atom (1913) — Quantum Mechanical Model of the Atom (1  Each new model contributed to the model we use today. Even our current Quantum Mechanical model, does not give us an exact model of how _____________behave. The Bohr Model of the Atom  Bohr used the simplest element, _____________, for his model  He proposed an electron is found in specific circular paths, or orbits around the nucleus  Each electron orbit was thought to have a fixed _____________level  Lowest level-ground state  Any Higher level-_____________state The Bohr Model of the Atom cont.  One electron is capable of many _____________excited states (whenever an electron jumps to higher level)  Quantum: specific amount of ______________ an electron can ________ or lose when moving energy levels  You can excite an electron with energy like electricity, ________________, or magnets Problems with the Bohr Model  OOPS!-Model only works with _______________.  Model did not account for the _________________ behavior of atoms  WRONG: _________________ do not move around the nucleus in circular orbits  STILL VERY HELPFUL!!! How do Neon Signs Work? They have __________________gases in them. Explanation  Step 1: an electron ______________ energy and moves to a __________________ energy level  Step 2: electron drops back down to a ________ energy level  During drop it gives off _______________ called a “photon”  Sometimes this energy is ______________ light (ROYGBIV)  When a photon is emitted, energy is released. We can calculate the energy released using the equation: ____________________ Application: Atomic Emission Spectrum  Used to determine which elements are present in a sample  Used to determine which elements are present in a star (because stars are gases)  Each element has a _________________ spectrum  Only certain _________________ are emitted because the energy released relates to a specific frequency Spectroscope  A spectroscope is needed to see the atomic emission spectra, which acts similar to a prism, separating different _________________ of light

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Electromagnetic Spectrum  Electromagnetic spectrum is the range of all energies emitted from photons acting like _____________. Electromagnetic Spectrum with Visible Light Spectrum

Light  

Behaves like a ______________. Behaves like a ______________.

Characteristics of a Wave  Wavelength  (lambda) – shortest ____________________ between equivalent points on a continuous wave [Unit = meters]  Frequency  (nu) – the ____________________ of waves that pass a given point per second [Unit = 1/second = s-1 = Hertz (Hz)]  Crest – _____________________ point of a wave  Trough – ____________________point of a wave  Amplitude (a)– height from its origin to its crest (highest point) or trough (lowest point) [Unit = meters] Wavelength and Frequency  Wavelength () and frequency () are related  As wavelength goes up, frequency goes down  As wavelength goes down, frequency goes up  This relationship is ______________proportional Wavelength and Frequency cont.  c =   Speed of light (c) = 3 x 108 m/s Question Time  Calculate the wavelength () of yellow light if its frequency () is 5.10 x 1014 Hz. 

What is the frequency of radiation with a wavelength () of 5.00 x 10-8 m? What region of the electromagnetic spectrum is this radiation?

How Much Energy Does a Wave Have?  Energy of a wave can be calculated  Use the formula E= h  E= Energy   = frequency  h = Planck’s constant = 6.626 x 10-34 Joule . Sec  Joule is a unit for energy (J)  Energy and frequency are directly proportional, as frequency increases, energy _________________ Question Time  Remember that the energy of a photon is E =h  How much energy does a wave have with a frequency () of 2.0 x 108 Hz? ( h = 6.626 x 10-34 J.s)

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Visible Light, Frequency, and Energy  Red _________________ wavelength (), smallest frequency ()  Red frequency smallest (), least amount of energy (E)  Violet smallest wavelength (), largest _________________ ()  Violet frequency largest (), greatest amount of energy (E) Flame Test  The flame test is a way to determine the _________________ present in a sample  When placed in a flame, each element gives off a ________________ color  Operates the same as neon signs; electrons are excited by _________ and fall back down and give off different colors Current Model of the Atom Quantum Mechanical Model of the Atom • Quantum Mechanical Model is the current description of electrons in atoms. – It does not describe the electron’s ______________________ around the nucleus • Quantum Mechanical Model is based on several ideas including: – Schrodinger wave equation (1926) treats electrons as _______________. – Heisenberg uncertainty principle (1927) states that it is impossible to know both the ____________________ and ______________________________of a particle at the same time. Where do electrons “live”? Principal Energy Levels 1. Principal energy levels n =1 to _______. (Row # on the periodic table) • The electron’s principal energy level is based on its location around the nucleus. • Electrons closer to the nucleus are at a __________________energy level and have lower energy than those farther away from the nucleus Atomic Orbital • An __________ ____________is a region of space in which there is a _________ ___________ of finding an electron – Orbitals ____ _____ necessarily spherical Energy Sublevels (also called orbitals) and Orbitals 1. Energy sublevels – assigned letters ______, _______, ________, or f (smart people do fine) – Energy sublevels correspond to a _______________ where the electron is likely to be found. 2. Orbitals – describes the electron’s ________________________ (maximum of ______ electrons per orbital) – s sublevel has 1 orbital (2 electrons total) - spherical – p sublevel has 3 orbitals (6 electrons total) – dumb-bell shaped – d sublevel has 5 orbitals (10 electrons total )-double dumb-bells – f sublevel has 7 orbitals (14 electrons total) Electron Configurations Energy Levels, Sublevels, and Orbitals 1. Principal energy levels – n, assigned values _______________ (Like floors in a hotel) 2. Energy sublevels- s, p, d, f (Type of suite in a hotel) (Orbitals are like the number of rooms in a suite) 1. s sublevel – 1 orbital 2. p sublevel – 3 orbitals 3. d sublevel – 5 orbitals 4. f sublevel – 7 orbitals 3. Orbitals – ___________ electrons per orbital (Two people per room) Electron Configurations • Electron configuration – the ______________________ of electrons in an atom. • Example Sodium (Na) – 1s22s22p63s1 • Three rules determine electron configurations – the Aufbau Principle, – the Pauli Exclusion Principle – Hund’s rule The Aufbau Principle • Each electron occupies the ______________________ energy orbital available

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• Like filling the hotel from the bottom up Pauli Exclusion Principle • A maximum of ______________electrons may occupy a single orbital • Like only two people sharing one hotel room Hund’s Rule • If two or more orbitals of _________ energy are available, electrons will occupy them ______________ with the same spin, before filling them in pairs with opposite spins • A spin is denoted with an up or down  arrow to fill orbitals • This is like trying to find your own room in the same suite before having to share a room with someone else Writing Electron Configurations • Aufbau diagram for sodium (Na) which has 11 electrons • Na electron configuration1s22s22p63s1 Exceptions to Electron Configurations • Copper and chromium are exceptions to the ___________________ principle. Element Copper Chromium

Should be 1s22s22p63s23p63d44s2 1s22s22p63s23p63d94s2

Actually is 1s22s22p63s23p63d54s1 1s22s22p63s23p63d104s1

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Some configurations violate the Aufbau Principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations Valence electrons • Valence electrons are electrons in the ______________________ orbitals. • For A group elements the ________________________ number corresponds to number of valence electrons. • Electron-dot structures – Element’s symbol surrounded by ___________representing the valence electrons

Noble Gas Configuration What are Noble Gases? • Noble gases are found in group _____________ • The elements are called noble because they are non-reactive and very __________________. • The do not tend to form compounds Complete Electron Configuration • What is the electron configuration for Ne? • Ne: ________________________ • What is the electron configuration for Mg? • Mg: _________________________ • What do both electron configurations have in common? • 1s22s22p6 = [Ne] Noble Gas Configuration (Abbreviated Configuration) • Using neon’s configuration and then adding magnesium’s extra electrons we can get the noble gas configuration. • Ne: 1s22s22p6 = [Ne] • Mg: 1s22s22p63s2 • Noble gas configuration Mg: __________________________ • Only use noble gases in the brackets. Which Noble Gas is Used? • To figure out which noble gas to use find the noble gas that is closest to the element without going over in atomic number • Which noble gas is closest without going over? • Rb : ____

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• Cl : ____ • Ra : ____ What About the Other Electrons? • To know what to write for the other electrons that are not included in the noble gas, understanding the periodic table is important. • The periodic table is organized by blocks according to the energy ____________________ • Blocks of the Periodic Table • There are s, p, f, and d blocks of the periodic table which correspond to the energy sublevels.

s Block Elements • Write the closest noble gas without going over in brackets. • Use the row number to get the energy level. • Count the number of electrons until you get to the element in the s block. • Mg _______________________________ Question Time • Try other s-block elements. Write the noble gas configuration of the following elements • Cs _______________________________ • Ca _______________________________ • Ba _______________________________ p block elements (Between 5-18) • Write the closest noble gas without going over in brackets. • Use the row number to get the energy level. • Write s2 after the row number because you have to go through the s-block to get to the p-block. • Write the row number again • The write “p” and then count the number of p electrons you must get through to get to your element as a superscript • Si: ________________________________ Question Time • Try other p-block elements. Write the noble gas configuration of the following elements • N : ________________________________ • S : ________________________________ • Cl : ________________________________ d block elements (Between 21-48) • Write the closest noble gas without going over in brackets. • Use the row number to get the energy level. • Write s2 after the row number because you have to go through the s-block to get to the d-block. • Write one less than the row number (d-block elements are always one less than the row number)**d for down one row number • Then write “d” and count the number of d electrons you must get through to get to your element as a superscript • Co: ____________________________ Question Time • Try other d-block elements. Write the noble gas configuration of the following elements • Ti : ________________________________ • Zn : ________________________________ • Mn : ________________________________ p block elements (Between 31-53) • Write the closest noble gas without going over in brackets. • Use the row number to get the energy level. • Write s2 after the row number because you have to go through the s-block to get to the p-block.

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Write one less than the row number (d-block elements are always one less than the row number)**d for down one row number • Then write “d” and count the number of d electrons you must get through to get to your element as a superscript • Write the row number again and “p” and count over the number of p electrons until you get to your element • Br: ______________________________ Question Time • Try other p-block elements. Write the noble gas configuration of the following elements • Sn : ________________________________ • Se: ________________________________ The Modern Periodic Table Early Periodic Table – Atomic Number • In 1913 Henry Mosley discovered that each element contained a unique number of protons in the nuclei • Arranged elements in order of atomic ___________________________. • Resulted in a clear periodic pattern of properties. Periodic Law • There is a periodic repetition of chemical and physical _______________________ of elements when arranged in increasing atomic number (increasing number of protons) is called the periodic ___________ Modern Periodic Table • Organized in columns called _________________ or families • Rows are called ________________________ • Group A – representative elements (1A-____________) • Group B - ___________________ elements (1B-8B) Classification of Elements • Three classifications for elements metals, nonmetals, and metalloids (semimetals) Metals • Properties of Metals – shiny, smooth, clean solids (except mercury) – High melting and boiling points – __________________conductors of heat and – ______________________ – bended or pounded into electricity sheets – High ______________________ – Ductile – drawn into _________________ Groups of Metals • ______________________ metals – group 1A except H • Alkaline earth metals – group ____________ – Alkali metals and alkaline earth metals are chemically reactive • Transition metals – group __________ elements • Inner transition metals – Lanthanide – Actinide Organizing by Electron Configuration • Group number for group A elements represents the number of ___________________ electrons • Atoms in the same group have similar chemical properties because they have the same number of valence electrons Alkali Metals • Electron configurations for alkali metals • Lithium ________________ [He]2s1 2 2 6 1 • Sodium 1s 2s 2p 3s [Ne]3s1 2 2 6 2 6 1 • Potassium 1s 2s 2p 3s 3p 4s [Ar]4s1 2 2 6 2 6 2 10 6 1 • Rubidium 1s 2s 2p 3s 3p 4s 3d 4p 5s [Kr]5s1 • What do the four configurations have in common? • They have a _____________________ electron in their outermost energy level • They all have one valence electron, thus similar chemical properties Alkaline Earth Metals • Electron configuration for alkaline earth metals • Beryllium [He]2s2 • Magnesium [Ne]3s2 • Calcium [Ar]4s2 • Strontium [Kr]5s2

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• All alkaline earth metals have _____________ valence electrons, thus similar chemical properties. Nonmetals • Gases or brittle, dull looking solids • ______________________ conductors of heat and electricity • Usually have lower densities, melting point, and boiling point than metals. • Groups of nonmetals – Halogens ____________ – Noble gases ____________ Noble Gases • Noble gases – Group _______________ • Called inert gases because they rarely take part in a reaction – He – 1s2 – Ne – 1s22s22p6 – Ar – 1s22s22p63s23p6 – Kr – 1s22s22p63s23p63d104s24p6 • Because noble gases have completely filled s and p sublevels, they do not react with other elements Metalloids (Semimetals) • Physical and chemical properties similar to both metals and nonmetals • They are metallic-looking _________________ solids • Relatively good electrical conductivity. • Used in glasses, alloys, and semiconductors • The six elements commonly recognized as metalloids are boron, silicon, germanium, arsenic, antimony, and tellurium. Polonium and astatine are sometimes classified as metalloids

Do the Trends w/s first! Periodic Trends Atomic Radius  Defined as ___________ of the distance between two bonding atom’s nuclei Atomic Radius Across a Period • Atomic radius generally ___________________________ in size as you move left to right across the period – ___________________ positive charge in the nucleus pulls the electrons of the same energy level in.

Atomic Radius Down a Group • Atomic radius _______________________________ as you move down a group – Orbital size increases as you move down a group with increasing energy level – Larger orbitals means that outer electrons are _______________________ from the nucleus. This increased distance offsets the greater pull of the increased nuclear charge. – As additional orbitals between the nucleus and the outer electrons are occupied, the inner electrons shield the outer electrons from the pull of the nucleus this is called __________________________.

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Cation and Anion • An ion is a positively or negatively charged atom that gains or loses an ___________________________. • A cation loses electrons and produces a _________________________ charge • An anion gains electrons and produces a _________________________ charge Ionic Radius - Cations • Groups 1A, 2A, 3A, and other metals _____________________ electrons and form cations. • When atoms lose electrons they become __________________________ – The electron lost will be a valence electron leaving a completely empty outer orbital – Protons in nucleus can pull fewer electrons tighter Ionic Radius - Anions • Group 5A, 6A, and 7A tend to ________________________ electrons and form anions • When atoms gain electrons and form negatively charged ions, they become ________________________. • Protons in nucleus have more electrons to pull and cannot pull in as tight Do Ionization and Electronegativity w/s First! Ionization Energy • The energy required to _________________________ an electron from a gaseous atom • Indication of how strongly an atom’s nucleus holds onto its __________________________electron • Groups 1A, 2A, and 3A tend to have low ionization energies because they want to lose electrons. Ionization Energy Trends – Across a Period • Ionization energy generally ________________________as you move left to right – Across a period electrons are added to the same energy level (same distance away from the nucleus), yet the nuclear charge is increasing across a period increasing the attraction to the electrons. Ionization Energy Trends – Down a Group • Ionization energy __________________________ as you move down a group – Down a group electrons are added to a higher energy level (farther distance away from the nucleus), making it easier to remove an electron Octet Rule • Sodium atom 1s22s22p63s1 • Sodium ion 1s22s22p6 (Sodium atom lost 1 electron) • Neon 1s22s22p6 • Sodium ion has the same electron configuration as neon • Octet rule states that atoms gain, lose, or share electrons to acquire a full set of ___________________ valence electrons (to be like a noble gas) Electronegativity • Indicates an element’s ability to _________________________ electrons in a shared chemical bond • fluorine (F) is the most electronegative element • Cesium (Cs) and francium (Fr)are the least electronegative • Noble gases do not tend to have an electronegativity number since they tend not to form __________________ Trends with Electronegativity • Electronegativity___________________________ as you move left-to-right across a period • Electronegativity _____________________________ as you move down a group

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