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Unit 3: Electron Arrangement & The Periodic Table This chapter will introduce the chemistry needed to understand Section 1: History of Atomic Theory & The Bohr Model Section 2: Electron Configuration Section 3: Orbital Notation Section 4 Electron Configuration for Ions Section 5: Light Section 6: The Periodic Table Section 7: Trends REMEMBER….. The atom is defined as the smallest particle of an element that retains the properties of that element. Section 4.0—The History of Atomic Theory The Beginning of the “ATOM” In 400 B.C., Democritus, a Greek philosopher expressed his idea that matter was made of very small, indivisible particles that he named “ATOMOS” Next in line: John Dalton John Dalton (17661844), an English school teacher and chemist, studied the results of experiments by other scientists. Dalton’s Atomic Theory : in 1803 1. All matter is made of tiny indivisible particles “atoms” 2. Atoms of the same element are identical; those of different atoms are different. 3. Atoms of different elements combine in wholenumber ratios to form compounds 4. Chemical reactions involve the rearrangement of atoms. Atoms cannot be created, divided, destroyed or changed into other types of atoms Dalton’s Atomic Model: tiny solid ball that could not be broken up into parts. What was wrong with Dalton’s Theory? Theories change The discovery of the subatomic particles led to the idea that the atom is not “Indivisible” Atoms of the same element can have different masses Next : J.J. Thomson & Cathode Ray Tubes A cathode ray is a ray of light traveling in a vacuum (no other particles inside) The ray travels from one metal plate to another as the plates are connected to electricity Cathode ray Metal plate (cathode) releases stream Metal plate (anode) to which stream travels Cathode Ray Tubes & Charge In the late 1800’s, JJ Thomson put charged plates outside the tube Negatively charged plate - + Positively charged plate Ray is deflected away from negative plate and towards positive plate It made no difference what type of metal he used in the tube—all material produced this stream that curved towards the positive charge Discovery of Subatomic Particles 1897 Thompson discovers the electron through the use of the cathode ray tube. He knows it is negatively charged and has an extremely small mass. Cathode Ray Movie Clip http://www.youtube.com/watch?v=O9Goys cbazk New Model Proposed Goldstein determines there are positive particles called protons and hey have a mass 1837 times heavier than the electron! Thomson’s evidence showed Dalton’s idea of solid, uniform atoms was incorrect. Thomson developed the “plum pudding” model. Since most of us aren’t familiar with plum pudding, you can think of it as a chocolate cookie dough theory Plum Pudding Model A.K.A. Cookie Dough Model The “chips” are the negative electrons. The “dough” is the positive portion The “chips” are stationary and don’t move within the “dough” Remember, officially this theory is called “plum pudding” but it’s the same idea! Next: Rutherford & Gold Foil Experiment: Early 1900’s He led a team of scientists They bombarded very thin gold foil with radioactive particles (alpha particles “”) They expected these relatively heavy particles to go through the atoms with a small deflection Gold Foil Experiment Basic video http://www.youtube.com/watch?v=5pZj0u_X Mbc Relates to the models http://www.youtube.com/watch?v=XBqHkraf 8iE What happened in the Experiment? What did he see? Most of the alpha particles passed straight through with no deflection These particles did not run into anything Some did deflect slightly These particles ran into something much smaller than themselves A few were reflected back the direction they came from These particles ran into something very dense Rutherford discovers the nucleus through the gold foil experiment & that atoms are mostly empty space. Electrons (small particles) caused the small deflections A small area of the atom with most of its mass (the protons) that caused the reflections. Rutherford’s Model Sphere with dense, middle center called the nucleus with electrons dispersed around it 1932: Chadwick: A third particle The protons and electrons could explain the charges of the various parts of the atom They could not explain the total mass of the atoms Chadwick confirms the neutron which has a mass similar to a protonand no charge. They were located in the nucleus Section 1: Next: Niels Bohr:1913 Bohr performed experiments with hydrogen atoms & light He determined that electrons are in levels according to how much energy they have and that only certain energy amounts were allowed. Think of Energy levels as rungs of a ladder. The Bohr Model: 1913 The farther away an energy level is from the nucleus, the more energy it attains. The circle closest to the nucleus contains the lowest energy electrons Energy Levels Electrons can move from one energy level to the next by gaining or losing energy (quanta). Ground State: An electron is as close to the nucleus as it can get. “excite d state” Excited State: An electron in a higher energy level than it should be. “ground state” Drawing Simple Bohr Models When assigning electrons, a max of 2 electrons are placed in the first shell, up to 8 in the 2nd shell, up to 18 in the 3rd shell, etc. Only 8 electrons can be placed in the 3rd shell at first, then 2 electrons will move into the 4th shell and the remaining of the 18 will be placed back in the 3rd shell for a total of 18. **(Just know this : it will be explained in the next section)** Valence electrons are the outermost electrons is found in the highest, outermost energy level of the atom Bohr Atomic Model http://www.youtube.com/watch?v=ZkNPBy 2i0ss http://www.youtube.com/watch?v=Cn6v5y gyZHQ Lewis Dot Diagrams A way to show the # of valence electrons in an atom. The symbol represents the nucleus and inner core electrons. The dots represent the valence electrons. Only 2 dots per side. Only 4 sides. Max. of 8 valence electrons. Add dots one at a time on each side until each side is full. Then add a second dot to make a pair when needed. Exception is Helium. It can look like this: He: or He Bohr Model In and then Out! We no longer believe electrons are in concentric circles, but this is still a convenient way to show energy levels on 2-dimensional paper 1920: At Last: Modern Atomic Theory: Quantum Mechanical Model Modern atomic theory uses calculus equations to show how the subatomic electrons act as both particles and waves These equations show the most probable location of electrons in the atom (known as atomic orbitals) Quantum Mechanical Model Section 2—Electron Structure The Electron Hotel Analogy Shopping Center Parking Garage Restaurant A man built an hotel for electrons with a restaurant next door. But he was making so much money that he decided to add on with some more rooms and a parking garage. He still had high demand and decided to add on some more rooms and a shopping center. He used the last space he could to put some rooms above the shopping center. How the Electron Hotel Fills Shopping Center Parking Garage Restaurant This man had some very strange ideas about how to run his hotel. He insisted four things: • The lowest possible must be used first (actually it was the fire inspector that insisted on this one) • There can only be one person in a room until all rooms at that level have someone • No more than 2 people to a room • When two people are in a room, they must be of opposite sex If 8 people come to the hotel, where would he put them? Another Example Shopping Center Parking Garage Restaurant This man had some very strange ideas about how to run his hotel. He insisted four things: • The lowest possible must be used first (actually it was the fire inspector that insisted on this one) • There can only be one person in a room until all rooms at that level have someone • No more than 2 people to a room • When two people are in a room, they must be of opposite sex If 21 people come to the hotel, where would he put them? You Try Shopping Center Parking Garage Restaurant This man had some very strange ideas about how to run his hotel. He insisted four things: • The lowest possible must be used first (actually it was the fire inspector that insisted on this one) • There can only be one person in a room until all rooms at that level have someone • No more than 2 people to a room • When two people are in a room, they must be of opposite sex If 42 people come to the hotel, where would he put them? Where do electrons really live? They don’t live in a hotel…They are in the area outside of the nucleus, the ELECTRON CLOUD Electron Clouds can be broken into: LEVEL = SHELL Electron Hotel Electron cloud Which section of the hotel Principal energy levels The electron cloud is made of energy levels. Which floor Sublevels Energy levels are composed of sublevels Which room Orbitals Sublevels have orbitals. Energy Levels There are 7 energy levels. Energy levels are also called shells The period number on the periodic table corresponds to the energy level Example: The energy level of the furthest electrons is: Ca is in energy level 4 Cl is in energy level 3 Sublevels (subshells): a set of orbitals with equal energy 4 sublevels exist “S” subshell Spherical shaped Only 1 orientation (position)= 1 orbital Maximum of 2 electrons Represented on the periodic table as groups 1A and 2A + helium First seen in the 1st energy level “P” Subshell Dumbbell shaped There are 3 orientations (positions)= 3 orbitals Maximum of 6 electrons Represented on the periodic table as groups 3A -8A First seen in the 2nd energy level “D” Subshell Four lobed shaped There are 5 orientations (positions)= 5 orbitals Maximum of 10 electrons Represented on the periodic table as the transition metals (group B) First seen in the 3rd energy level “F” Subshell Too complex to name shaped There are 7 orientations (positions)= 7 orbitals Maximum of 14 electrons Represented on the periodic table as the inner transition metals (lower block) First seen in the 4th energy level Within a Sublevel are Orbitals Orbital– Area of high probability of the electron being located. Each orbital can hold 2 electrons To calculate the total number of orbitals, use 2(n)2 Summary Electron Configuration Is an address of an electron Electrons must be placed in the lowest possible energy levels first (ground state) 4p 1 Energy Level # of electrons subshell 3 rules that govern electron configurations Aufbau Principle: Electrons must fill the lowest 1 available subshells and orbitals before moving on to the next higher energy subshell/orbital. Filling order is: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p Can use the periodic table as a guide to fill from lowest to highest sublevels or memorize filling order An orbital’c cloud space can overlap within different energy levels The 4s is lower in energy than the 3d Energy and Subshells-Filling Order 6p 6s 5p 5d 4f 4d 5s 4p 3d 4s 3p 3s 2p Energy 2s Subshells are filled from the lowest energy level to increasing energy levels. Does this look familiar? Electron Hotel! 1s How Electron Configurations Relate to the Organization of the Periodic Table s p d f Figure 11.31: Orbitals being filled for elements in various parts of the periodic table. Hund’s Rule 2 Hund’s Rule: Place electrons in unoccupied orbitals of the same energy level before doubling up. How does this work? If you need to add 3 electrons to a p subshell, add 1 to each before beginning to double up. Pauli Exclusion Principle 3 Pauli Exclusion Principle: Two electrons that occupy the same orbital must have opposite spins. “Spin” is designated with an up or down arrow. How does this work? If you need to add 4 electrons to a p subshell, you’ll need to double up. When you double up, make them opposite spins. Electron Configurations 1 Determine the number of electrons to place 2 Follow Aufbau Principle for filling order 3 Fill in subshells until they reach their max (s = 2, p = 6, d = 10, f = 14) or use periodic table as a guide. 4 The total of all the superscripts is equal to the number of electrons. Example: Write electron configuration for S No charge written Charge is 0 Atomic number for S = 16 = # of protons 0 = 16 - electrons 1s 2 2s 2 2p 6 3s 2 3p 4 Electrons = 16 2 + 2 + 6 + 2 + 4 = 16 Electron Configurations Example: Write electron configuration for K No charge written Charge is 0 Atomic number for K = 19 = # of protons 0 = 19 - electrons Electrons = 19 1s 2 2s 2 2p 6 3s 2 3p6 4s1 2 + 2 + 6 + 2 + 6 + 1 = 19 Electron Configurations Example: Write electron configuration for Ti No charge written Charge is 0 Atomic number for Ti = 22 = # of protons 0 = 22 - electrons Electrons = 22 1s 2 2s 2 2p 6 3s 2 3p6 4s2 3d2 2 + 2 + 6 + 2 + 6 + 2 + 2 = 22 Unit 3: Other Notations What is Orbital Notation? It shows the grouping and position of electrons in an atom. Orbital Notation use boxes or lines for orbitals and arrows for electrons. Drawing Orbital Notation (boxes & arrows) 1 Aufbau Principle: Electrons fill subshells (and orbitals) so that the total energy of atom is the minimum 2 Hund’s Rule: Place electrons in unoccupied orbitals of the same energy level before doubling up. 3 Pauli Exclusion Principle: Two electrons that occupy the same orbital must have different spins. Example: Write the orbital notation for Cl Drawing Orbital Notations 1 Aufbau Principle: Electrons fill subshells (and orbitals) so that the total energy of atom is the minimum 2 Hund’s Rule: Place electrons in unoccupied orbitals of the same energy level before doubling up. 3 Pauli Exclusion Principle: Two electrons that occupy the same orbital must have different 3p spins. No charge written Charge is 0 Example: Write the orbital notation for Cl 1s Atomic number for Cl = 17 = # of protons 0 = 17 - electrons Electrons = 17 2s 2p 4 9 70 6 5 3 2 1 8 17 16 15 14 13 12 11 3s 3p Drawing Orbital Notations No charge written Charge is 0 Example: Write the orbital notation for Fe Atomic number for Fe = 26 = # of protons 0 = 26 - electrons Electrons = 26 1s 2s 2p 3s 3p 4s 3d Shorthand Notation (aka: Noble Gas Notation) Noble Gas– Group 8 of the Periodic Table. They contain full valence shells. The Noble gas is used to represent the core (inner) electrons and only the valence shell is shown. Br Complete electron configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 Noble gas [Ar] 4s 2 3d 10 4p 5 The “[Ar]” represents the core electrons and only the valence electrons are shown Which Noble Gas Do You Choose? How do you know which noble gas to use to symbolize the core electrons? Think: Price is Right. How do you win on the Price is Right? By getting as close as possible without going over. Choose the noble gas that’s closest without going over! Noble Gas # of electrons He 2 Ne 10 Ar 18 Kr 36 Xe 54 Noble Gas Notation Example 1 Determine the number of electrons to place 2 Determine which noble gas to use 3 Start where the noble gas left off and write electron configuration for the valence electrons Example: Write noble gas notation for Fe Noble Gas Notation Example 1 Determine the number of electrons to place 2 Determine which noble gas to use 3 Start where the noble gas left off and write spectroscopic notation for the valence electrons No charge written Charge is 0 Atomic number for Fe= 26= # of protons Example: Write noble gas notation for Fe 0 = 26- electrons [Ar] 4s 2 3d 6 Electrons = 26 Closest noble gas: Ar (18) Ar is full up through 3p 18 + 2 + 6 = 26 Noble Gas Notation Example Example: Write noble gas notation for Ba No charge written Charge is 0 Atomic number for Ba= 56 = # of protons 0 = 56 - electrons Electrons = 56 Closest noble gas: Xe (54) Xe is full up through 4p [Xe] 6s 2 54 + 2 = 56 Section 4: Electron Configuration for Ions 1 Determine the number of electrons to place. Positive ions lose electrons; negative ions gain electrons. 2 Follow Aufbau Principle for filling order 3 Fill in subshells until they reach their max (s = 2, p = 6, d = 10, f = 14) or use periodic table as a guide. 4 The total of all the superscripts is equal to the number of electrons. Example: Write electron configuration for S-2 Atomic number for S = 16 = # of protons -2 = 16 - electrons 1s 2 2s 2 2p 6 3s 2 3p 6 Electrons = 18 2 + 2 + 6 + 2 + 6 = 18 Example 1 Determine the number of electrons to place. Positive ions lose electrons; negative ions gain electrons. 2 Follow Aufbau Principle for filling order 3 Fill in subshells until they reach their max (s = 2, p = 6, d = 10, f = 14) or use periodic table as a guide. 4 The total of all the superscripts is equal to the number of electrons. Example: Write electron configuration for K+1 Atomic number for K = 19 = # of protons +1 = 19 - electrons 1s 2 2s 2 2p 6 3s 2 3p 6 Electrons = 18 2 + 2 + 6 + 2 + 6 = 18 Another Example 1 Determine the number of electrons to place. Positive ions lose electrons; negative ions gain electrons. 2 Follow Aufbau Principle for filling order 3 Fill in subshells until they reach their max (s = 2, p = 6, d = 10, f = 14) or use periodic table as a guide. 4 The total of all the superscripts is equal to the number of electrons. Example: Write electron configuration for Fe+2 Atomic number for S = 26 = # of protons +2 = 26 - electrons 6 1s 2 2s 2 2p 6 3s 2 3p 6 3d Electrons = 24 2 + 2 + 6 + 2 + 6 + 6 = 18 Exceptional Configuration Half-filled or completely filled d & f sublevels have LOWER energies and are more stable than partially filled d’s and f’s. This means that an atom can “borrow” one of its “s” electrons from the previous orbital to become more stable. ___ ↑↓ __ ↑↓ __ __ __ __↑ ↑↓ ↑↓ ↑↓ 5s 4d becomes ↑ ___ 5s ↑↓ __ ↑↓ __ ↑↓ __ ↓↑ __ ↑↓ __ 4d Because the 4d sublevel is now full, the atom is at a lower energy state and therefore more stable. YOU TRY: Write the exceptional configuration for copper, Cu This means that an atom can “borrow” one of its “s” electrons from the previous orbital to become more stable. ↑↓ ___ ↑__ ↓ 4s ↑ ↓ __ ↑ ↓ __ ↑ ↑ ↓ __ __ 3d becomes ↑ ___ 4s ↑↓ __ ↑ ↓ __ ↑↓ __ ↓↑ __ ↑↓ __ 3d Because the 4d sublevel is now full, the atom is at a lower energy state and therefore more stable. Section 5—Light Light is Electromagnetic Radiation Electromagnetic radiation is all energy that travels through space at the speed of light which equals 3.0 x 108 m/s Examples of Electromagnetic Radiation…visible, microwaves, infrared, ultraviolet, radio waves, gamma, x-rays ElectroMagnetic Radiation Spectrum Wave Properties—Wavelength Wavelength () is the distance from trough to trough of a wave (measured in meters “m”) 1nm = 1 x 10-9 m wavelength Wave Properties—Frequency Frequency () is the number of times a wave completes a cycle in one second (cycles per second is “Hertz” or “Hz”) Lower frequency Higher frequency Reference Sheet Lower frequency Lower Energy Higher frequency Higher Energy Relationship between wave properties Notice the shorter the wavelength, the higher the frequency. This is an INVERSE relationship. c = Notice as frequency increases, the amount of energy increases. This is a DIRECT relationship. Visible Light The visible light region is a small region within the spectrum that has wavelengths/frequencies that are eyes can detect. ROY G BIV is a mnemonic to help you remember the colors of the spectrum. Red is near infrared and violet is near ultra violet. Reference Sheet Examples Using Reference Sheet 1. Which of the following forms of electromagnetic radiation has the shortest wavelength? a) gamma b) visible c) infrared d) radio As the frequency of electromagnetic radiation increases, its wavelength ______________. a) increases b) decreases c) remains constant d) is impossible to determine. 3. Which of the following forms of radiation has photons with greatest amount of energy? a) red light b) yellow light c) green light d) violet light A 2. B D Interesting superhero facts: • Superman has x-ray vision. • The Incredible Hulk was “created” by an accidental overdose of gamma radiation. • The Fantastic Four were “created” •by cosmic rays. Bohr Model 4. What type of electromagnetic radiation is represented by a wavelength of 1870 nm? a)Infrared light c) Ultraviolet A b) visible d) x-ray Examples Using Reference Sheet C 5. What type of electromagnetic radiation is represented by a wavelength of 4.7x10-1m? a) gamma rays b)infrared c) microwaves d) visible light LIGHT & MATTER: Visible Range Wavelength increases Frequency decreases Energy decreases 400 nm 700 nm Visible light White light is made of all the colors…a prism can separate white light into a rainbow! Continuous Spectrum: Sun light (or white light) will produce a range of color because there are no specific wavelengths Line Spectrum Is when individual atoms emit light of only certain wavelengths. Each element has its own line spectrum, or fingerprint. How can a line spectrum be explained? Electrons Absorbing Energy Energy packets called photons or quanta come into contact with an atom & collide with an electron. + The electron is “excited” to a higher energy level with is newly increased energy from absorbing the photon. Electrons Absorbing Energy Photon coming into atom collides with electron. Photons are energy. + Excitation The process of an electron absorbing a photon of light (energy) and being promoted to a higher energy level from its “ground state” And later… The electron cannot remain in that excited state indefinitely + And later… The electron cannot remain in that excited state indefinitely + Energy is released during relaxation Relaxation The process of an electron releasing a photon of light (energy) and falling back down to a lower energy level. Energy of photon and levels jumped The higher the energy of the photon, the greater the electron jump! A photon of UV light has more energy than a photon of Infrared light The UV photon would cause a higher energy jump (jump up more levels) than the IR photon. Total energy in = Total energy out However much energy was absorbed must be released again, but it can be released in smaller packets A high energy photon might be absorbed, but two lower energy photons might be released as the electron falls in a “stepwise” manner. Photons must match energy changes The energy of the photon must exactly match the energy change of the electron. If the photon is not an exact match, the photon will pass through unabsorbed. + Hydrogen Line Spectrum The colored lines are the wavelengths of light that are emitted when an electron moves from a higher E level to a lower E level, This was proof that atoms had fixed energy levels! Emission Spectrum Hydrogen Spectrum Neon Spectrum How hydrogen produces the four visible photons Reference Sheet Examples 1. On the energy level diagram below, draw an arrow representing the electron in hydrogen’s ground state being excited to the fourth energy level. Examples 2. An electron in the hydrogen atom makes the transition n = 5 n = 3. a. Determine the wavelength of light associated with this transition. Include units. a) 434 nm b) 434 m A c) 1282 nm b. Classify the type of electromagnetic radiation this wavelength represents: a) infrared b) visible light D c) ultraviolet d) x-ray c. Is this energy emitted by the atom or absorbed by the atom ? emitted d) 1282 m Flame Tests Metals can be identified by the wavelength of light they emit. When metals absorb energy from a flame, the electrons absorb energy and are raised to higher energy level. When they return to their ground state, they release the energy they absorbed in the form of radiation. The wavelength of light for some metals fall in the visible light portion of the spectrum. This allows us to see their color. Ways of producing light Fluorescence: visible light is absorbed and visible light is emitted at the same time— the relaxation happens very quickly after excitation Phosphorescence: Visible light is absorbed and then a while later is emitted—relaxation occurs after a period of time Section 6—The Periodic Table History of the Periodic Table Different scientists organized the elements differently—this lead to confusion In 1869, Dimitri Mendeleev designed a periodic table based on atomic mass. This way showed patterns in properties that repeated across rows and similarities down columns He couldn’t find elements to fit all the property trends, so he left holes Mendeleev’s Periodic Table History of the Periodic Table The holes Mendeleev left were later filled in as more elements were discovered The modern periodic table is arranged by atomic number rather than atomic mass. This caused a few “switches” in placement, but overall is very similar to Mendeleev’s Henry Mosley is given credit for the modern periodic table. Modern Periodic Law: Elements in columns have similar properties! Organization: Groups and Periods Periods Rows are called “periods” Groups Columns are called “groups” or “families” Information for Each Element Most periodic tables give the following information, but it can be in a different location Atomic Number Element Symbol If there’s a second letter, it’s lower-case 6 C Carbon 12.01 Average Atomic Mass Number with decimals Gives the mass for 1 mole of atoms, in grams THIS IS NOT THE MASS NUMBER! Whole number— elements are ordered by this on the periodic table. Element Name Modern Periodic Table Divisions The rows at the bottom Most periodic tables are written with 2 rows at the bottom. This is done to allow the font to be bigger on a piece of paper. The rows at the bottom Most periodic tables are written with 2 rows at the bottom. This is done to allow the font to be bigger on a piece of paper. But they really belong here! Follow the atomic numbers on your periodic table to see it! Properties of Metals All solid except for mercury Formability Malleable: can be flattened into thin sheets Ductile: can be drawn into fine wire Great conductors of heat & electricity High luster High strength Properties of Nonmetals Mostly gases, few solids and 1 liquid, Br Poor formability Poor conductors of heat & electricity dull brittle Sulfur Properties of Metalloids(semi-metals) All solids Share properties of both metals & nonmetals Semiconductors Electron Configuration Patterns Look at the electron configurations for the Halogens F 1s2 2s22p5 Cl 1s2 2s2 2p6 3s23p5 Br 1s2 2s2 2p6 3s2 3p6 4s2 3d104p5 I 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d105p5 All of the elements in Group 7 end with 5 electrons in a p subshell. Every Group ends with the same number of electrons in the highest energy subshell. (Similar Valence electrons of a group explains why elements in the same column have similar properties.) Configurations and the Periodic Table s-block p-block d-block s1 s2 p1 p2 p3 p4 p5 p6 d1 d2 d3 d4 d5 d6 d7 d8 d9 d10 f-block f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 2 Types of Elements Representative elements(main group): Group A See special families info. In the next set of slides Transition Elements: Group B Properties Have more than 1 oxidation number or ion charge Typically the least reactive of the metals Hardest, densest, and highest melting points Greatest conductors of heat and electricity Parts of the Periodic Table Parts of the Periodic Table Family: Alkali Metals Group 1A Most reactive metals; always found in compound form Reacts violently with water to produce hydrogen gas and a base Soft and silver in color Low density & melting points Has 1 valence electron and forms +1 ions Cool Clip Showing Reactivity of Alkali Metals http://www.youtube.com/watch?v=m55kgy ApYrY Family: Alkaline Earth Metals Group 2A less reactive than 1 A but still extremely reactive; always found in compound form Reacts with water to produce hydrogen gas and a base – Not Explosive Extracted from ores-mineral rocks Harder,higher density & melting points than 1A Has 2 valence electrons and forms +2 ions Family: Halogens Group 7A all are non metals most reactive nonmetals; always found in compound form Most form the diatomic molecules Has 7 valence electrons and forms -1 ions Family: Noble Gases Group 8A All nonmetals exist as gases nonreactive elements- most stable Called inert gases All have 8 valence electrons except Helium which only has 2 and does not form ions Other Important Facts Hydrogen is a diatomic non metal & in a group by itself; can have a +1 or -1 charge 58-71 Lanthanides: 1st row of lower block; called Rare-Earth metals 89-103 Actinides: 2nd row of lower block; all are radioactive and most are manmade Section 7—Trends &Periodicity What is periodicity on the periodic table? The predictable pattern by which properties of elements change across or down the periodic table. There are always exceptions to these periodicity trends…each of the trends is a “general” trend as you move across a period or down a group. Three Important Terms to help explain TRENDS FOR A PERIOD TREND ONLY: “NUCLEAR CHARGE”: the number of protons in the nucleus; determines the strength of attraction to the electrons FOR A GROUP TREND ONLY: “ENERGY LEVEL”: location of the electrons; the higher the energy level, the farther distance the electrons are from the nucleus FOR A GROUP TREND ONLY:”SHIELDING” : energy levels that are between the nucleus and valence electrons cause this Trend 1 :Atomic Radius Defined as: Half of the distance between the nuclei of two bonded atoms. H H Distance between nuclei Atomic radius of hydrogen atom Atomic Radii Trends Period trend :Decreases Group trend: Increase s Atomic Radii Period Trend: DECRESES WHY? As the number of protons increases, the nuclear charge increases. e e e n p n p n p e Lithium atom Move across the periodic table Radius decreases e e pn p p p n nn e Beryllium atom As the nuclear charge increases, the attraction between the positive nucleus and negative electron cloud increases. This attraction “pulls” in on the electrons. Atomic Radii: Group Trend: INCREASES WHY? Nuclear Charge also increases as you go down a group. HOWEVER: The electrons are added into new energy levels. e e e + Move down the periodic table e e Radius increases e e e e e e e Lithium atom e + e Sodium atom The inner electrons “shield” the new outer electrons from the pull of the nucleus, therefore it doesn’t pull in like the last slide. Examples List the following in INCREASING order of atomic radius . Li, Cs, K List the following in DECREASING order of atomic radius . Ca, Be, Ba, Sr Who has the largest atomic radius? Mg, Cl, Na, P Li < K < Cs Ba > Sr > Ca > Be Na Trend 2: Ionization Energy Defined as: The energy needed to remove the outermost electron. Ionization Energy Trends Period Trend: Increases Group Trend: Decreases Ionization Energy Period Trend: INCREASES WHY? Moving left to right, the size of the atom decreases as more protons pull harder on more electrons. Move across the periodic table e e n p n p n p e e Radius decreases IE increases Lithium atom e e pn p p p n nn e Beryllium atom When an atom is smaller, the electrons are closer to the nucleus, and therefore feel the pull more strongly. Thus, it requires more energy to remove electrons away from these smaller atoms. Ionization Energy Group Trend: DECREASES WHY? As you move down a group, the radius increases as more electrons shells are added. e e e + Move down the periodic table Radius increases IE decreases e e e e e e e e e Lithium atom e + e Sodium atom As the outer electrons (those involved in bonding) are farther from the nucleus, they will feel the “pull” of the nucleus less. It is easier to remove an electron from a larger atom thus requires less energy. Examples List the following in INCREASING order of ionization energy . Li, Cs, K List the following in DECREASING order of ionization energy . Ca, Ba, Be, Sr Who has the highest ionization energy? Cl, Na, I, In Cs < K < Li Be > Ca >Sr > Ba Cl Trends in the Periodic Table “Successive Ionization Energies” “Successive Ionization Energies” means the energy required to remove a 2ndor a 3rdelectron from an atom. Removing more and more e-’s requires more& more energy. Why? nucleus. The remaining e-’s are more tightly bound tothe Trend 3: Electronegativity Defined as: Tendency of an atom to steal an electron when combining with another element F is the highest Fr is the lowest Electronegativity Trends Period trend: Increases Group Trend: Decreases Electronegativity Period Trend: INCREASES WHY? Moving left to right, the radius of the atom decreases as more protons pull on more electrons. Move across the periodic table e e n p n p n p e e Radius decreases EN increases Lithium atom e e pn p p p n nn e Beryllium atom When an atom is smaller, the nucleus pulls more strongly. This can attract & draw an electron away from a different atom. Electronegativity Group Trend: DECREASES WHY? As you move down a group, the radius increases as more electrons shells are added. e e e + Move down the periodic table Radius increases e e e e e e EN decreases e + e e e Lithium atom Sodium atom The larger atom has weaker attractions so it cannot draw electrons away very easily from different atoms. e Examples Who has the highest electronegativity? Ba, Br, Ca Br List the following in DECREASING order of electronegativity. I, Cl, Br Cl > Br > I Trend 4: Ionic Radius Review Some Definitions An Ion is defined as– atom with a charge. Cation– are positively charged ions. Results from a metal losingelectrons. Anion– are negatively charged ions. Results from a nonmetals gaining electrons. Cations Are Smaller than a Neutral Atom WHY? Atoms lose electrons to create positive ions e e e + Lithium atom Creating a cation, losing electrons Radius decreases e e + Li+ ion When electrons are lost, there are now more protons than electrons Therefore, the protons have a greater “pull” on each of the electrons – making a smaller ion. Trends in the Periodic Table Atomic Size vs. Ion Size Anions are larger than a neutral atom WHY? Atoms gain electrons to create negative ions e e e e Creating an anion, gaining electrons e + e e e e e e e e + Radius increases e e e Oxygen atom e e O2- ion When electrons are gained, there are now more electrons than protons Therefore, the protons have a weaker “pull” on each of the electrons. They loosen up & spread out making a larger ion Ionic Radii Period Trend for METALS & NONMETALS: decreases Group Trend: Increases Examples Arrange in order of decreasing ionic radius? P -3 Mg +2 Cl-1 Arrange in order of INCREASING ionic size K+1 Cs+1 Li +1 P -3 > Cl-1 > Mg+2 Li +1 < K+1 < Cs+1 Trend 5: Reactivity Defined as: How chemically active an atom is to another Metal Reactivity Trends Period Trend Decreases Group Trend Increases METAL Reactivity Trend: DECREASES WHY? Move across the periodic table e e n p n p n p e e Radius decreases IE increases Lithium atom e e pn p p p n nn e Beryllium atom The most reactive metals have a low ionization energy. Nonmetal Reactivity Trends Period Trend Increases Group Trend decreases NON Metal Reactivity Trend: INCREASES WHY? Move across the periodic table e e n p n p n p e e EN increases IE increases Lithium atom e e pn p p p n nn e Beryllium atom The most reactive nonmetals have a high Ionization energy & electronegativity!