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Unit 5: -lecture Regents ’14-‘15 Mr. Murdoch Unit 5: Electrons

1.

Student Name: ______Class : ______Page 1 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

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Page 2 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Unit 5 Vocabulary:

1. Anion: a negatively charged . 2. Cation: A positively charged ion. 3. : A particle with a net charge of -1 and a mass 1/1836th of an amu located in the levels outside the nucleus. Electrons are lost, gained, or shared in the formation of a . (e-) 4. : An ’s attraction to electrons in a chemical bond. Electronegativity is used to determine bond types, polarity of a molecule, and attractive force type and strength. 5. : A condition where an atom’s electrons occupy higher energy levels than they normally would. 6. Frequency (ν) & (Hz): The number of wavelengths that pass a fixed point in one second. 7. Ground State: A condition where an atom’s electrons are occupying the lowest possible energy state. 8. Ion: A charged atom or molecule formed by the gain (or loss) of electrons. 9. Ionic Radius: The measure of the size of an ion. 10. Energy: The energy required to remove an atom’s most loosely held electron, measured when the element is in the gas phase. 11. Kernal: The atom beneath the valence electrons, including the rest of the electrons in the lower energy levels and the nucleus. 12. Orbital: A region of space around the nucleus that is the most likely location to find an electron in an atom. 13. Orbital Notation: Also known as “box diagrams”, these schematics describe the location and spin of the electrons in an atom. 14. Oxidation: The loss of electrons from an atom or ion. 15. : An infinitesimally small particle that travels in a wave-like fashion after being released when electrons fall from the excited state to the ground state. A photon is also known as quanta. 16. Planck’s Constant (h): A proportionality constant that converts Hz (frequency) to Joules (energy). Planck’s Constant = 6.6 x 10-34 J/Hz 17. Reduction: The gain of electrons by an atom or ion. 18. Quantum: A four-digit series of numbers that identifies the location of a specific electron around the nucleus based on PEL, sublevel, orbital, and spin. 19. Shell (Principal Energy Level/PEL): The most probable location an electron may be found around the nucleus.

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20. Stable Octet: An that is reached when gain, lose, or share electrons in an attempt to achieve a electron configuration of eight valence electrons. is an exception to this “Rule of Eight”. 21. Sublevel: The regions of space electrons occupy making up a principal energy level. 22. Speed of Light (c): The velocity (speed) of light in a vacuum, or 299,792,458 meters/second. 23. Valence electrons: The electrons that reside in the outermost principal energy level of an atom. These electrons are lost, gained, or shared in the formation (or decomposition) of a chemical bond. 24. Wavelength (λ): The distance from one crest to the next crest in a wave. Measured in meters.

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Unit 5 Homework Assignments:

Assignment: Date: Due:

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Topic: An Overview of the Electron Objective: How did the electron come about to be understood?

History of sub-atomic particle discovery:

Watch History of Subatomic Particles video https://www.youtube.com/watch?v=kBgIMRV895w

The of the atom and electrons:

Watch Bozeman Science The Bohr Atom video https://www.youtube.com/watch?v=GhAn8xZQ-d8

The Quantum-Mechanical Model of the atom:

 The current model of the atom has the atom containing a small, dense positively charged nucleus surrounded by electrons that travel in a wave-like motion around the nucleus. This motion is modified by mass and charge interactions between electrons and the nucleus. The interactions and the fast speed of the electron make it impossible to know both where an electron is and where it is going at any particular moment. All that can be known is a general area of probable space where the electron might be. Electrons travel in principal energy levels made up of sublevels, with each sublevel

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made up of orbitals that contain two electrons in each orbital. Electrons in the same orbital will spin in opposite directions.  The end result is that the current model of the atom shows the electron’s around a nucleus closely resemble a cloud of gnats (the electrons) buzzing around your head (the nucleus).

The Development of the Atom Model:

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Topic: Light Objective: Why does high matter emit light?

1. Electrons (charge of -1 and having a mass of 1/1836th amu) surround the nucleus of an atom in distinct energy levels. Electrons (e-) occupy the lowest possible energy levels when the atom is in the ground state.

2. When electrons have energy added (in the form of light, heat, or electricity), electrons will rise in energy level by the same amount of energy that the electrons were given. The higher the energy added, the higher the electrons rise in level. The higher energy level is called the excited state. This is in accordance with the Law of Conservation of Energy, stating energy may not be created or destroyed by physical or chemical change.

3. As electrons are negatively charged, they are attracted to the positively charged nucleus, and will eventually release the excited state energy as they fall back to the ground state.

4. The energy released as an electron returns to the ground state is in the form of photons. Photons are the smallest known particles, and are essentially massless. Photons travel at the fastest theoretical

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speed possible, 3 x 108 m/sec, or the speed of light. Photons are considered particles of light.

5. The color of the light emitted is determined by the amount of energy released by the electron as it dropped back to the ground state. Light particles travel in a wave pattern. The more energy a photon has, the shorter its wavelength. Photons with wavelengths shorter than visible light (high energy) are found in gamma rays, X- rays, and . Visible light photons make up a very small part of the electromagnetic spectrum and are in the middle of the photon energy range. For visible light, the higher energy photons are violet, and the lower energy photons are red. Photons with wavelengths longer than visible light (lower energy) are infrared, radar, microwave, and radio.

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Topic: Light Objective: Why does high temperature matter emit light?

6. There are three properties of light waves: energy (Ɛ, in joules); wavelength (λ, in meters); and frequency (ν, in wavelengths per second, or Hz). Note that frequency, (ν), is the number of wavelengths that pass a given point in one second. All photons travel at the speed of light (3 x 108 m/sec), so all photons have the same speed. Therefore, a greater number of short wavelength photons will pass a certain point each second than will long wavelength photons.

Watch Khan Academy Introduction to Light video https://www.youtube.com/watch?v=rLNM8zI4Q_M

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Topic: Electromagnetic Spectrum Objective: How do wavelength, frequency, and energy all relate?

As the energy of a photon increases, the wavelength shortens and the frequency increases as more short wavelengths may pass a point in one second.

The Continuous Electromagnetic Spectrum (from 2006 Ref Table)

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Topic: Electron Energy levels Objective: What are the different drops of energy for electrons?

Each atom has many electron levels. Electrons can drop from the excited state in different pathways. If an electron is excited from the first to the fourth energy level, the electron may fall back in one of the following pathways.

i. From the 4th to the 1st energy level; ii. From the 4th to the 3rd to the 1st energy level; iii. From the 4th to the 2nd to the 1st energy level; iv. From the 4th to the 3rd to the 2nd to the 1st energy level.

Each different drop in energy level emits light with a different amount of energy, and therefore a different wavelength (color). A sample of pure element subjected to electrical current excites the sample’s electrons. When the excited electrons fall back toward the ground state, the differing amounts of energy will give off different colors of light. The sample will look as if it glows a certain color. If the light emitted is projected through a prism onto a white backdrop, the individual of the emissions may be seen as bright lines (spectra) of different colors. Each element has unique spectra, allowing the testing of unknown materials by exciting the electrons of a sample

Page 12 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch and then comparing the results with known standard spectra. This process of viewing spectra is known as spectroscopy, and the tool for observing spectra is the spectroscope.

The spectrum above is the bright-line spectrum for hydrogen in the visible (to humans) range of approximately 780 nm to 390 nm. Note that there are no lines for energy level drops for the 4th to 3rd energy levels or the 2nd to 1st energy levels. That’s because those two examples emit photons outside of the visible range. If the emitted photons are outside our visible range, other detectors (gamma- and X- ray, ultraviolet light, infrared, microwave, and radio wave detectors).

Watch Spectral Lines video https://www.youtube.com/watch?v=fKYso97eJs4

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Notes page:

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Student name: ______Class Period: ______Please carefully remove this page from your packet to hand in. The Electron Model and Light homework Circle your answers for the multiple choice questions below.

1. The current quantum-mechanical model of the atom says that electrons a) Are mixed in evenly with positive charge b) Are found orbiting a negatively-charged nucleus c) Are found orbiting a positively-charged nucleus d) Are in regions of probability called orbitals around a nucleus

2. How is light formed, in terms of energy levels?

3. What is a particle of light called? ______

4. If the frequency of a photon is increased, what happens to the photons wavelength?

5. If the energy of a photon is decreased, what happens to the photons frequency?

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6. Why may bright-line spectra of an element be able to identify that specific element?

An unknown sample is analyzed using spectroscopy. The unknown material generated the spectrum shown at the bottom of the diagram below. The known spectra of four other elements, ‘A’ through ‘D’, are also shown.

7. Based on the reference spectra for elements ‘A’ through ‘D’, what elements make up the unknown sample? Circle the letter of the element(s) you choose.

8. Explain why you selected the element(s) you chose.

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Notes page:

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Topic: Electron Configuration Objective: How do we best show atomic structure?

The electron configuration of an atom tells you where the electrons are located. This is important because the outermost (valence) electrons in an atom are responsible for all the physical and chemical properties of the elements and any compounds they form.

The Four Types of Electron Configurations

Shell configuration:

The Shell Configuration is found at the bottom left of each element box on the Period Table of Elements. The Shell Configuration numbers tell you how many electrons are located in each principal energy level. This is along the idea of the original Bohr model energy level model.

i. For hydrogen (H), the shell configuration is simply 1, as it only has 1 electron. Page 18 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

ii. For (Cu), the shell configuration is 2 electrons in the 1st PEL, 8 in the 2nd PEL, 18 in the 3rd PEL, and 1 electron in the 4th PEL. iii. For lead (Pb), the shell configuration is not fully shown. For elements with a large number of electrons, the innermost electrons are left off the chart to make reading the chart easier. Therefore, Pb starts with -18 on the Periodic Table, meaning that the first 2-8 electrons are in the first 2 PELs are assumed to be there. Pb has electrons in more than four PELs. iv. Now look at (U). Note that there isn’t a shell configuration shown for uranium. That is because since uranium has multiple isotopes, there is no stable electron count.

Shell configuration for (Ca): How many electrons does calcium (Ca) contain?

This is the Bohr model for calcium. 1. Note the nucleus: 20 protons (atomic # of 20) with 20 neutrons (mass # of 40) 2. Note the 1st ring around the nucleus (1st energy level) contains 2 e-, spaced as far apart as possible (e- repel each other) 3. Note the 2nd ring (2nd energy level) contains 8 e- 4. Note the 3rd ring (3rd energy level) contains 8 e-

5. Note the 4th ring (4th energy level) contains 2 e-

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Topic: PEL and the Periodic Table Objective: What is the significance of PELs in the Periodic Table?

Elements in the same PERIOD (horizontal row) across the Periodic Table all have the same number of energy levels (shells). Therefore, all of the elements in Period 3 have three energy levels in their electron configuration, and all of the elements in Period 4 have four energy levels in their electron configuration.

Period 3 has three PELs

Period 4 has four PELs

Watch Configuration video https://www.youtube.com/watch?v=R-DW2hGta2I

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Notes page:

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Topic: Electron Configuration Objective: How do we best show atomic structure?

Sublevel Configuration:

Principal Energy Levels (PELs) are made up of sublevels; much like a town is made up of streets. The expanded configuration tells you how many electrons are found in each sublevel of each PEL. Most of the time (and especially in Regents Chemistry class!) one sublevel must fill up completely before the next sublevel will gain any electrons.

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Writing Sublevel Configurations: ’s shell configuration is 2-1. Sublevel: 1s2 2s1

1s2 means that the 2 e- in the 1st energy level are in the 1s sublevel 1 - nd 2s means that the 1 e in the 2 energy level is in the 2 s sublevel *Note the first sublevel filled is ALWAYS the 1s sublevel ’s shell configuration is 2-5. Sublevel: 1s2 2s2 2p3

1s 2 means that the 2 e- in the 1st energy level are in the 1s sublevel 2s2 means that 2 e- in the 2nd energy level are in the 2s sublevel and the other 3 e- go into the 2p sublevel until they run out. ’s shell configuration is 2-8-2. Sublevel: 1s2 2s2 2p6 3s2

1s2 means that the 2 e- in the 1st energy level are in the 1s sublevel 2s2 means that 2 e- in the 2nd energy level are in the 2s sublevel and the other 6 e- go into the 2p sublevel until they run out of room. The 3rd energy level has 2 e- that go into the 3s sublevel. ’s shell configuration is 2-8-7. Sublevel: 1s2 2s2 2p3 3s2 3p5

1s2 means that the 2 e- in the 1st energy level are in the 1s sublevel 2s2 means that 2 e- in the 2nd energy level are in the 2s sublevel - and the other 6 e go into the 2p sublevel until they run out of room. The 3rd energy level has 7 e- that place 2 e- into the 3s sublevel, then the other 5 e- go into the 3p sublevel.

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Topic: Orbital Box Diagrams Objective: How do show number and motion of electrons in orbitals?

Orbital Box Diagrams:

 Orbital box diagrams are graphic depictions of how many electrons are in each orbital of each sublevel, as well as the spin (rotation) of the electrons movement. As sublevels are further classified into orbitals, each orbital can be diagramed. Orbitals contain a maximum of two electrons, and those electrons spin (rotate) as they orbit around the nucleus. If there is only one electron in an orbital, it will have an UP spin. If a second electron is in that same orbital, it will have a DOWN spin.  Orbitals are shown as boxes with arrow(s). The arrow(s) represent the spin of the electron(s) present. Orbitals are filled in order, always placing the UP arrows in each empty box before placing any DOWN arrows in any box.

Orbital Box Diagram Rules:

i. The number of orbitals is equal to one-half the number of electrons EACH sublevel may contain.

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ii. When two electrons occupy the SAME orbital, they MUST have OPPOSITE spin. When electrons are placed into a box, all UP arrows are entered first. iii. A single orbital with ONE electron looks like: iv. A single orbital with TWO electrons looks like: v. The sublevel name (1s, 2p, etc.) is written ABOVE each orbital box.

vi. The lowest energy level (1s) is fully filled before any electrons are placed into the next higher (2s) energy level.

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Topic: Orbital Box Diagrams Objective: How do show number and motion of electrons in orbitals?

Orbital Box Diagram Rules (Cont’d):

Notice for nitrogen (N) above in the 2p sublevel there are three orbitals with ONLY up (spin) arrows. 2p3 states that there are three electrons in the 2p sublevel. As there are only three orbitals in the 2p sublevel, we place all the UP spin electrons FIRST in each available 2p box.

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Looking at magnesium (Mg) and chlorine (Cl), we see that orbital notation is useful for finding UNPAIRED electrons in an orbital. You will see soon enough that the unpaired electrons are what bond (join) atoms together, so understanding the number and locations of the unpaired electrons will help in understanding how the atoms bond. If you see two arrows in the same orbital box, the electrons are PAIRED. Again, we fill ALL orbital boxes with UP arrows in a sublevel BEFORE any DOWN arrows are entered.

Kernal:

 A Kernal is everything underneath the valence shell. This would include the nucleus (nucleons) and any filled non- shells.  Kernal electrons take no part in chemical bonding process for compound formation.

Watch Bozeman Science Orbitals & Electron Configuration Video https://www.youtube.com/watch?v=2AFPfg0Como

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Topic: Valence Electrons Objective: How are the outermost electrons shown in an atom?

The outermost electrons in an atom are known as valence electrons. The number of valence electrons that an atom has may be determined by the last number in the basic electron configuration as listed on the Periodic Table for that element.

Nitrogen has Lithium has TWO Magnesium has Chlorine has has THREE TWO energy energy levels. THREE energy THREE energy energy levels. levels. The 2nd The 2nd energy levels. The 3rd levels. The 3rd The 3rd energy energy level level contains 1 energy level energy level level contains 8 contains 5 e-. e-. This is the contains 2 e-. contains 7 e-. electrons. These These are the valence e-. These are the These are the are the valence valence e-. valence e-. valence e-. e-. 8 valence e- is the most (noble) e- an element may contain.

The number of valence electrons that any atom has determines both the physical and chemical properties of that atom. The number of valence electrons will influence almost all aspects of chemistry you will continue to learn.

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Topic: e- Configuration Order Objective: How are the sublevels of electrons ordered?

Electron Configuration Order:

To determine the electron configuration of a particular atom, start at the nucleus and add electrons one by one until the number of electrons equals the number of protons in the nucleus. Each added electron is assigned to the lowest-energy sublevel available. The first sublevel filled will be the 1s sublevel, then the 2s sublevel, the 2p sublevel, the 3s, 3p, 4s, 3d, and so on. This order is difficult to remember and often hard to determine from energy-level diagrams.

A more convenient way to remember the order is to use the diagram below. The principal energy levels are listed in columns, starting at the left with the 1s level. To use this figure, read along the diagonal lines in the direction of the arrow. The order is summarized under the diagram.

The arrow shows a second way of remembering the order in which sublevels fill.

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Topic: Lewis Dot Diagrams Objective: How may we diagram the electrons around an atom?

Lewis Dot Diagrams:

Lewis Dot Diagrams use dots (single or paired) around the symbol of an atom to represent the valence electrons. For NYS Regents Chemistry you should only need to know valence electrons occupying the s and p orbitals. As you have learned, the ‘s’ orbital fills with electrons first, then the ‘p’ orbitals fill with ‘up’ electrons before any ‘down’ electrons are added.

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Topic: Valence e- & the Period Table Objective: Why are valence electrons important in the periodic table?

Importance of Valence Electrons in the Periodic Table:

The Period Table of Elements (Reference Table p.9) is arranged in order of increasing . Dmitri Mendeleev designed the Period Table of Elements in 1869 based on the fact that as atomic number increases there were repeating groups of atoms with similar chemical properties. Mendeleev grouped these similar chemical property atoms in VERTICAL columns called GROUPS. All elements in the same vertical share similar chemical properties with each other atom in that vertical group. It wasn’t until later that atoms in the same vertical group were found to have the same number of valence electrons. These vertical groups with the same number of valence electrons form with the same nuclear charges, and therefore have similar chemical properties.

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Topic: Stable Objective: Why are 8 valence electrons important in elements?

Stable Octet:

A Stable Octet is an atom that contains eight valence electrons. All atoms on the Periodic Table want to gain or lose electrons until they end up with a stable octet of valence electrons. Noble gases (Group 18) already contain eight valence electrons, which is why noble gases do not readily react with other elements.

Watch Sci Show's Mendeleev Periodic Table Video https://www.youtube.com/watch?v=-wu0LixSBpk

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Student name: ______Class Period: ______Please carefully remove this page from your packet to hand in. Electron Configuration homework Using the one box with information, fill in the missing information in the table.

1s22s22p63s23p64s23d104p2

1. How many valence electron(s) are there in the noble gases (Ne, Ar, Kr, etc.)?

2. How many valence electron(s) are there in Group 1 (Li, Na, K, etc.) elements?

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Topic: Excited and Ground States Objective: How may we practically use the quantum atomic model?

Excited state and Ground State Electrons:

 Ground State: The Ground State is when electrons occupy the lowest possible PEL and sublevels. The Ground State is the configuration shown on the Periodic Table of Elements.  Excited State: The Excited State is when electrons absorb energy and rise to a higher (excited) energy level.  Light: When Excited State electrons fall back towards (or to) the Ground State, they release energy in the form of photons of light. The greater the distance traveled towards the Ground State, the higher is the photon’s energy.

1. Note that in the many possible configurations for the excited state, the number of TOTAL electrons remains the SAME.

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Topic: Excited and Ground States Objective: How may you tell if Ground or Excited State electrons?

Shell Configuration (PEL):

2. Add up all the electrons in the given configuration. Use the total number of electrons (that # is also the atomic number) to identify the element the given configuration belongs too. If the number of given electrons in the given configuration matches the periodic table configuration for the element with the same number of protons as the given number of electrons, then you have the ground state for that element. If the given electron configuration does NOT match the periodic table configuration for the element, it is in an excited state.

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Topic: Excited and Ground States Objective: How may you tell if Ground or Excited State electrons?

Sublevel Configuration:

3. Add up all the electrons in the configuration. Use the total number of electrons (that # is also the atomic number) to identify the element the given configuration belongs too. Write the given expanded configuration of the element. If the expanded configuration matches the configuration you find on the Periodic Table, then the configuration is in the ground state. If the expanded configuration given does not match that element’s Periodic Table configuration, then it is in the excited state.

Watch NYS Regents Question on Excited electrons https://www.youtube.com/watch?v=j5G-8RvnsPs&index=13&list=PLpSkLXW6NBslqGv27LrC0YGQXJ-ZqqIer

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Topic: Excited and Ground States Objective: How may you tell if electrons positions are full or not?

Occupied: An occupied PEL, sublevel, or orbital is considered OCCUPIED if there is at LEAST one electron in that PEL, Sublevel, or orbital.

Full: A full PEL, sublevel, or orbital is considered FULL if the maximum number of electrons are in that space.

Orbitals with ONE arrow are OCCUPIED; with TWO arrows are FULL.

Occupied and Full Orbital Examples:

1. In a (Na) atom

(#)

(letters) (↑↓)

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2. In a cobalt (Co) atom:

Watch Crash Course Chemistry The Electron video https://www.youtube.com/watch?v=rcKilE9CdaA

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Notes page:

Page 40 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Student name: ______Class Period: ______Please carefully remove this page from your packet to hand in. Excited and Ground Interpretation homework

1. Identify the following electron configurations as being in the ground or excited state.

Complete the following data charts. 2. Complete the data charts below for the element (O).

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3. Complete the data charts below for the element (K).

Circle your answers for the multiple choice questions below. 4. What is the total number of valence electrons in an atom of in the ground state? a) 2 b) 4 c) 8 d) 14

5. An atom in the ground state has seven valence electrons. This could be an atom of which element? a) Oxygen b) Sodium c) Calcium d)

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Topic: Electronegativity Objective: How do the different elements relate in attraction?

Electronegativity:

 Electronegativity is an atom’s attraction to electrons when the atom is involved in a chemical bond. Remember that an individual atom will be neutral (# of protons = # of electrons) when alone. The higher an elements electronegativity, the stronger the attraction that element has for another elements’ electrons. Therefore, when two atoms of different elements bond, the atom with the higher electronegativity will tend to gain electrons from the atom with the lower electronegativity.  The electronegativity scale was developed by based on fluorine (F) having the HIGHEST (strongest) electron attraction, with an electronegativity value of 4.0. All other elements have electronegativity values based on fluorine as given in Reference Table S.

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Topic: Ionization Characteristics Objective: What constitutes the characteristics of forming ions?

Ionization Energy:

is the amount of energy (kJ/ of atoms) that an atom must absorb to remove the most loosely bound valence electron from a gas state atom forming a +1 ion of that atom.

 The higher the ionization energy, the more energy required to remove the most loosely bound valence electron in that gas state element. Ionization Energy values are given in Reference Table S.

Atomic Radius:

 The Atomic Radius of an atom is the size of the entire atom, from the center of the nucleus to the outer edge of the electron cloud.  Atomic Radius INCREASES from top to bottom within a group (vertical column) as the elements are shown on the Periodic Table of Elements. This is a function of increasing number of PELs going from top to bottom within a group (column), placing the valence electrons farther from the nucleus.

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Atomic Radius examples:

Group 1 elements

 Atomic radius DECREASES as the elements on the Periodic Table from left to right across a period (horizontal row). The elements within period (row) have the same number of energy levels, but with increasing atomic number (# of protons), the increasing positive nuclear charge attracts the negatively charged electrons more, drawing the electrons closer to the nucleus.

Period 2 elements

Therefore, the general TREND is the closer to (Fr) has a larger atomic radius, and the closer to fluorine (F) has a smaller atomic radius.

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Topic: Electronegativity Trends-Ions Objective: How does electronegativity drive loss or gain of ions?

Loss of Electrons is Oxidation (LEO):

1. Atoms with LOW electronegativity () tend to LOSE electrons (oxidation) and form positive (+) ions, called CATIONS. 2. Atoms have a neutral charge (# of + protons = # of - electrons). When an atom loses an electron, the number of protons (+ charge) is greater than the number of electrons (- charge), and the ion has a positive charge (cation). 3. The charge of the positive ion is + (# of electrons lost). 4. If an atom loses 2 e-, the ion is +2 charged; 3 e- lost is +3 charged.

Gain of Electrons is Reduction (GER):

1. Atoms with HIGH electronegativity () tend to GAIN electrons (reduction) and form negative (-) ions, called ANIONS). 2. Atoms have a neutral charge (# of + protons = # of - electrons). When an atom gains an electron, the number of protons (+ charge) is less than the number of electrons (- charge), and the ion has a negative charge (anion). 3. The charge of the ion = - (# of electrons gained). 4. If an atom gains 1 e-, the ion is -1 charged; 2 e- gained is -2 charged.

Page 46 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

The charge of an ion may be found on the Periodic Table in the upper corner of each element box. If negative charges are listed first (e.g. Cl), use the first charge listed. Some elements are capable of forming more than one charge, either negative or positive. More on that later... 

When atoms form ions, the ions will have a stable octet of 8 valence electrons mimicking a noble gas configuration (Group 18).

Chlorine (Cl):  The first listed charge is -1, which is chlorine’s ionic charge. The other charges deal with later chemistry.  To form the Cl-1 ion, chlorine must GAIN 1 e- (reduce)  A chlorine atom has 17 e- in an electron configuration of 2-8-7. When Cl becomes Cl-1, the new configuration is 2- 8-8, or the same as the stable octet noble gas argon (Ar).

Magnesium (Mg):  Magnesium forms a +2 ion. To form the Mg+2 ion, Mg must LOSE 2 e- (oxidation).  A magnesium atom has 12 e- in an electron configuration of 2-8-2. When Mg becomes Mg+2, the new configuration is 2-8, or the same as the stable octet noble gas (Ne).

Page 47 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Topic: Ionic Radius Objective: How will the overall size of the atom differ with charge?

Ionic Radius:

1. Negative ions are larger than the original atom as the gained electrons repel each other. i. When a chlorine (Cl) atom (2-8-7) gains an electron to form a chlorine (Cl-1) ion (2-8-8), the extra valence electron repels the other electrons, making the negative ion larger than the original atom. 2. Positive ions are smaller than the original atom as the outermost electrons have been lost. i. When a magnesium (Mg) atom (2-8-2) loses two electrons to form a magnesium (Mg+2) ion (2-8), the third energy level is empty, making the positive ion smaller than the original atom.

Page 48 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Page 49 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Topic: Ionic Dot Diagrams Objective: How do we diagram ionic configurations of atoms?

Ionic Dot Diagrams:

 Positive ions (CATIONS) have a pair of brackets around the element symbol, and the positive (+) charge number on the upper right side of the brackets. There are NO dots shown inside the brackets, as the positive ion lost the same number of outermost valence electrons as the charge number.

 Negative ions (ANIONS) have a pair of brackets around the element symbol, and the negative (-) charge amount on the upper right side of the brackets. There are EIGHT dots shown inside the brackets, as the negative ion gained the same number of electrons as the charge number to form a stable octet of eight valence electrons.

Page 50 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Page 51 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Topic: Subatomic Particles in Ions Objective: What atomic particles change during ionization of atoms?

When an atom becomes an ion, the only accounting that changes is the number of electrons. The number of protons (the atomic number) remains the same, and the number of neutrons is still the mass number minus the atomic number.

Subatomic ionization examples:

41 +2 1. How many protons, neutrons, and electrons are found in a 20Ca ion? i. The atomic number of Ca is always 20, so there are 20 protons. ii. The mass number is 41, so there are (41-20) = 21 neutrons. iii. The ion charge is +2, meaning the Ca atom lost 2 e-. In any neutral atom, the number of protons (20 in Ca) equals the number of e- (20), and since a +2 e- charge means a more positive ion (loss of 2 negative e-), there ends up being 18 electrons.

35 -1 2. How many protons, neutrons, and electrons are found in a 17Cl ion? i. The atomic number for Cl is always 17, so there are 17 protons. ii. The mass number is 35, so there are (35-17) = 18 neutrons.

Page 52 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch iii. The ion charge is -1, meaning the Cl atom gained 1 e-. A Cl atom has 17 protons, therefore the Cl atom started with 17 e- and the -1 charge means a more negative ion (gain of 1 negative e-), with an end result of having 18 electrons.

Page 53 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Topic: Naming Positive Ions Objective: How do we name atoms when they lose electrons?

Naming Positive Ions:

 When naming positive ions (most elements), if only ONE positive charge is listed in the elements box on the periodic table, then positive ion name will be the same as the element name.

NA+1 is named a Cd+2 is named a Sr+2 is named a H+1 is named a sodium ion cadmium ion strontium ion hydrogen ion

 If two (or more) positive charges are listed, write the charge as roman numerals in parentheses after the element name.

Cu has two Mn has two charges shown: charges shown: Cu+1 is named Mn+2 is named copper(I); CU+2 is (II); named copper (II) Mn+3 is named manganese (III); +4 Mn is named manganese (IV); Mn+7 is named manganese (VII)

Page 54 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Topic: Naming Negative Ions Objective: How do we name atoms when they gain electrons?

Naming Negative Ions:

When naming negative ions (most nonmetals), the ion charge is the first charge listed. Ignore all of the other charges. The name of the negative ion becomes the first syllable of the elements name with the suffix “-ide” added.

Cl is chlorine; Cl-1 O is oxygen; O-2 becomes C is ; C-4 becomes becomes chloride carbide

Watch Naming Ions video https://www.youtube.com/watch?v=WijZBA6tXfw

Page 55 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Notes page:

Page 56 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Student name: ______Class Period: ______Please carefully remove this page from your packet to hand in. Ions homework

Circle your answers for the multiple choice questions below. 1. Which atom has the strongest attractions to other atom’s electrons? a) Na b) C c) N d) F

2. An atom of which element below is most likely to lose an electron when bonding with oxygen? a) Na b) C c) N d) F

3. An atom of which element below is most likely to gain an electron when bonding with magnesium? a) Na b) C c) N d) F

4. Which atom listed below will require the most energy to remove the most loosely held valence electron? a) Cl b) Li c) Rb d) Ne

5. Which atom listed below will be most likely to lose valence electrons and form a positively charged ion? a) Cl b) Li c) Rb d) Ne

6. How many valence electrons does a nitrogen atom have? a) 4 b) 5 c) 7 d) 8

7. According to Reference Table S, as the elements in Group 2 are considered from top to bottom, what happens to atomic radius? a) Increases b) Decreases c) Remains the same

8. According to Reference Table S, as the elements Na to Cl are considered from left to right, what happens to the atomic radius of the atoms? a) Increases b) Decreases c) Remains the same

Cont’d next page

Page 57 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

9. What is the reason why the radius of is larger than the radius of fluorine? a) More PEL’s c) More neutrons b) More electrons d) More nuclear charge

10. What is the reason why the radius of oxygen is smaller than the radius of ? a) More PEL’s c) More neutrons b) More electrons d) More nuclear charge

11. What must a nitrogen atom do to form an ion with a stable octet? a) Gain 3 e- (reduction) c) Lose 3 e- (reduction) b) Gain 3 e- (oxidation) d) Lose 3 e-(oxidation)

12. How many electrons does a Na+1 ion have? a) 0 b) 1 c) 10 d) 11

13. When nitrogen forms a -3 ion, it forms the electron configuration of which element? a) Be b) N c) Ne d) Na

14. What is the charge of an ion that has 10 electrons and 9 protons? a) +1 b) -1 c) +9 d) -10

15. What is the name of the Fe+3 ion? a) Iron b) Iron(I) c) Iron (II) d) Iron (III)

16. What is the name of the S-2 ion? a) c) Sulfide b) Sulfur (II) d) Sulfide (II)

Page 58 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Student name: ______Class Period: ______Please carefully remove this page from your packet to hand in. Ions homework (cont’d) Complete the chart below for the subatomic particles for ions in these isotopes.

For each of the elements below, write the ion charge from the periodic table, draw the dot diagram for the ion, including the charge, and then give the number of electrons lost (e.g. lost 2 e-, gained 1 e-) to form the ion. Compare your ion configuration (dot diagram) to the Periodic Table to find which Noble Gas shares the same Electron Configuration as the ion you drew.

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Complete the chart below using the clues for the missing selections.

Complete the chart below by naming the given ions and then writing the ion symbols and charges for the given ion names.

Page 60 of 61 Website upload 2014 Unit 5: Electrons-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Notes page:

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