Chapters+4+and+5

Atomic Structure and Electrons In Atoms: A Chapter 4 & 5 Wiki Page By: Period C Editor: Marion Burdick





Introduction to WikiPage: This page has been contructed on the fourth and fifth chapters of the Prentice Hall's //Chemistry// book. Chapter four focuses on defining and exploring early models of the atom in addition to the structure of a nuclear atom. The distinctions between atoms based on atomic number, mass number, isotopes, and atomic mass are also mentioned in chapter 4. Chapter five explores the development of the model of an atom and discusses the quantum mechanical model as well as electron configuration. The quantum mechanic theory introduces concepts of physics including that of light, wavelength, frequency, and the atomic spectra. This wiki-page outlines these two chapters in their entirety, and includes information crucial to the understanding of these chapters.

Group #1  pages 100- 109
 * Mike McShane //(co-editor)//
 * Kendyl Barron
 * Adam Shanahan
 * Christos Anastos
 * Heather Bowditch
 * Maggie Bie

=**__Early Models of the Atom (pg 101) – Kendyl Barron__**= -->The atom is the smallest particle of an element that can still be identified as that element in a chemical reaction. Subatomic particles include protons, a postively charged particle, electrons, small negatively charged particle, and neutrons, particles of a neutral charge. These particles make up an atom and give an element certain properties.

--> For years before the discovery of the indiviual atom scientists and philosophers contemplated the structure of small invisible particles that supposedly made up matter.The Greek p hilosopher Democritus first purposed the possible existence of these particles believed to be so small they were invisible to  the naked eye. He believed atoms (atomos meaning invisible) were also indestructible. Democritus' approach was not scientific, however, though his ideas were complatible with much later scientific discoveries.

=__**Daltons Atomic Theory (pg 102) – Adam Shanahan **__=

Compounds- Adam Shanahan


 * John Dalton (1766- 1844) began the modern process of discovery regarding atoms.
 * **Dalton transformed Democritus's ideas on atoms into a scientific theory by using experimental methods**
 * Dalton created hypotheses and theories based on his experiments resulting in **Dalton's atomic theory**
 * 1) All alements are made up of atoms.
 * 2) Atoms of the same element are identical. No two atoms of different elements have all of the same properties.
 * 3) Atoms of different elements can be physically mixed together or chemically combined in simple whole-number ratios to form compounds.
 * 4) Chemical reactions occur when atoms are joined, seperated, or rearranged. Atoms of one element, however, cannot be changed into atoms of another element because of a chemical reaction.

=__**Sizing up the Atom (pg 103) – Mike McShane **__= -->An atom has all of the same qualities as a piece of that thing at any size. For example: Although a copper penny is many, many times larger than one atom of copper, the copper atom has exactly the same properties as a copper penny. -->An atom is VERY small. To put this in perspective, a copper penny contains 2.4x10^22 copper atoms and there are 6x10^9 people on earth. This means that there are 4x10^12 more atoms in one tiny coin than there are people on the earth, atoms are VERY small. - Penny (Mike McShane) -->Although atoms are this small, scientists are able to observe atoms with instruments called scanning tunneling microscopes. These microscopes give scientists the ability to manipulate and arrange atoms into patterns. This is a big advancement because this moves us on our way to creating atomic sized technology. -Scanning Tunneling Microscope (Mike McShane) =__**Subatomic Particles (pg 104-105) – Christos Anastos **__= Cathode Ray (Christos Anastos)
 * =**__ ﻿ __ Atoms can be broken down into smaller particles caleed subatomic particles. There are 3 types of subatomic particles. They are... ** =
 * Electrons
 * Protons
 * <span style="font-family: Arial,Helvetica,sans-serif;">Neutrons
 * //__<span style="font-family: Arial,Helvetica,sans-serif;">Electron- __//<span style="font-family: Arial,Helvetica,sans-serif;"> A negativly charged subatomic particle
 * <span style="font-family: Arial,Helvetica,sans-serif;">Electrons were discovered by J.J. Thomson in 1897 through the invention of the cathode ray
 * <span style="font-family: Arial,Helvetica,sans-serif;">The cathode ray sent the beam of particles towards the positive side. Since opposites atract, Tomson understood that the particles must be negative.
 * <span style="font-family: Arial,Helvetica,sans-serif;">Thomson originally named them corpuscles
 * <span style="font-family: Arial,Helvetica,sans-serif;">The charge and mass of an electron were discovered by Rovert A. Millikan. The electron has a negative charge of one unit, and a mass 1/1840 of a hydrogen atom

=__<span style="font-family: Arial,Helvetica,sans-serif;">Protons and Neutrons (pg 106) – Heather Bowditch __= = =  What is left of an atom after it loses its electrons? Use these four simple ideas about electrical charges to help you answer that question: 1. Atoms have no electrical charge to begin with; they are electrically neutral. 2. Electric charges are carried by subatomic particles. 3. Charges always exist in whole numbers; there are no fractions of electric charges. 4. When a given number of negatively charged particles combine with the same number of positively charged particels, it becomes neutral. In 1886, Eugen Golstein observed a cathode ray and discovered rays travelling in opposite direction. He concluded that this particle must have a positive charge. These subatomic positively-charged particles are called protons. Protons have a mass of about 1,840 times that of an electron. In 1932, James Chadwick concluded that another subatomic particle existed, having no charge, but the same mass of a proton. This is called a neutron. PROPERTIES OF SUBATOMIC PARTICLES Particle Symbol Relative charge Relative mass (when mass of proton =1) Weight Electron e - 1- 1/1840 0 Proton p + 1 1 1 Neutron n 0 0 1 1 Pictures Eugen Goldstein Relative Size and Mass == =__**<span style="font-family: Arial,Helvetica,sans-serif;">The Atomic Nucleus (pg 106) – Maggie Bie **__= <span style="font-family: 'Times New Roman','serif'; font-size: 12pt; margin: 0in 0in 0pt;">As subatomic particles were discovered, no scientist could understand how they were put together in an atom. <span style="font-family: 'Times New Roman','serif'; font-size: 12pt; margin: 0in 0in 0pt;">J.J. Thompson (discoverer of the electro) believed electrons were evenly distributed throughout an atom filled uniformly with positively charged material.



<span style="font-family: 'Times New Roman','serif'; font-size: 12pt; margin: 0in 0in 0pt;">A student of Thomson – Ernest Rutherford – was the first to form a new, accurate ato

<span style="font-family: 'Times New Roman','serif'; font-size: 12pt; margin: 0in 0in 0pt;">mic model.

**<span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">Rutherford’s Gold-Foil Experiment (pg 107) – Maggie Bie **

<span style="font-family: 'Times New Roman','serif'; font-size: 12pt; margin: 0in 0in 0pt;">Rutherford and his coworkers decided to test the current theory of atomic structure. <span style="font-family: 'Times New Roman','serif'; font-size: 12pt; margin: 0in 0in 0pt;">They used relatively massive alpha particles – helium atoms that have lost their 2 electrons and have a double positive charge

<span style="font-family: 'Times New Roman','serif'; font-size: 12pt; margin: 0in 0in 0pt;">In the experiment: a narrow beam of alpha particles was directed at a very thin sheet of gold foil __<span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">According to the previous theory __<span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">: the particles should have passed easily through with only a slight deflection – this would have been due to the positive charge throughout the gold atoms

<span style="font-family: 'Times New Roman','serif'; font-size: 12pt; margin: 0in 0in 0pt;">Instead, the majority of particles passed through with no deflection at all. Also, a small fraction of the alpha particles bounced off the foil at large angles. Some also bounced straight back to the source

<span style="font-family: 'Times New Roman','serif'; font-size: 12pt; margin: 0in 0in 0pt;"> <span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">media type="youtube" key="5pZj0u_XMbc?fs=1" height="385" width="480"Maggie Bie

**<span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">The Rutherford Atomic Model (pg 107)– Maggie Bie **

<span style="font-family: 'Times New Roman','serif'; font-size: 12pt; margin: 0in 0in 0pt;">Rutherford proposed that the atom was mostly empty space, thus the lack of deflection of most particles. <span style="font-family: 'Times New Roman','serif'; font-size: 12pt; margin: 0in 0in 0pt;">The positive charge and almost all mass are concentrated in a small region with enough positive charge to account for the deflection at large angles for a few of the particles. <span style="font-family: 'Times New Roman','serif'; font-size: 12pt; margin: 0in 0in 0pt;">This region was called the nucleus __<span style="font-family: 'Times New Roman','serif'; font-size: 12pt;">nucleus- __<span style="font-family: 'Times New Roman','serif'; font-size: 12pt;"> the tiny central core of an atom and is composed of protons and neutrons



His discoveries are known to us as the nucleur atom. During his experiement, he observed the following things: __﻿__
 * the protons and neutrons are located in the nucleus
 * electrons are distributed around the nucleus and occupy almost all the volume of the atom
 * nucleus is tiny compared to the atom. the atom is large in density

<span style="font-family: 'Arial Black',Gadget,sans-serif;">Group #2 pages 110- 132
 * Tom DeMarco //(co-editor)//
 * Katherine Perry
 * Elizabeth Sieber
 * Nick Brault
 * Eileen Corkery

Elements are different based on the protons they have in their nucleus. The number of protons an element has is represented by its atomic number. Most of the mass of an atom is concentrated in the nucleus of an atom The total number of nuetrons and protons found in an atom is called the mass number. ex. mass number is 9, atomic number is 4. The composition of an atom can be shown in short hand. The figure below shows how an atom of gold is represented using this notation. The chemical symbol. Au appears with two numbers written to its left. The atomic number is the subscript. The mass number is the superscript. Notation for Gold (Elizabeth Sieber)
 * __Atomic Number- Elizabeth Sieber__**
 * number of electrons can change
 * number of neutrons can change
 * number of protons NEVER change for a specific element
 * atomic number represents ALL isotopes of an element
 * ex. if an atom has an atomic number of eight, it must be oxygen
 * __Mass Number- Elizabeth Sieber__**
 * nucleus is made of the protons and the nuetrons
 * Mass number can be found by adding the total number of protons and nuetrons
 * if given the atomic number and mass number, you can tell how many protons and then can also tell how many neutrons
 * number of nuetrons = mass number - atomic number
 * you can tell that ther are 4 protons
 * subtract 9-4=5
 * there are 5 nuetrons
 * Isotopes**


 * __Atomic Mass-Katherine Perry__**

The mass of atoms are very small. An atomic mass unit (amu) is defined as 1/12 of the mass of a carbo-12 atom. The mass of a single proton or neutron is 1/12 of 12 amu or 1 amu The atomic mass of an elemenet is a weighted average mass of athe atoms in a naturally occurring sample of the element. To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products. The sum is the weighted average mass of the atoms. __**EXAMPLE:**__ __YOU KNOW:__ -isotope 10X: mass = 10.012 amu relative abundance = 19.91% = 0.1991 -isotope 11X: mass = 11.009 amu relative abundance = 80.09% = 0.8009 __UNKNOWN:__ atomic mass of element X

__CALCULATE:__ for 10X : 10.012 amu x 0.1991 = 1.993 amu for 11X: 11.009 amu x 0.8009 = __8.817 amu__ for element X: atomic mass = 10.810 amu -Katherine Perry


 * __The Development of Atomic Models__**

Niels Bohr was a Danish physicist and a student of Ernest Ru
 * __The Bohr Model pg. 128-129__** Thomas DeMarco

therford. He believed that Rutherford’s atomic model needed improvements.He changed the model to include the fact that the energy of an atom changes when it absorbs or emits light. Bohr believed that the electrons of an atom could only be found in specific circular paths, also known as orbits, around the atom.

In Bohr’s model, each electron level has a fixed energy and the fixed energy these levels can have are called energy levels. The model says the more energy levels there are; the farther away the electron is from the nucleus. Also, the electrons must be found on these energy levels and they can’t be in between energy levels. A quantum of energy is the amount of energy needed to move an electron from one energy level to another. The amount of energy used varies. If the energy levels are closer they require less energy but if they are farther apart they require more energy. The Bohr model did well with experiments for the hydrogen atom but did not explain all energy transfers for atoms with more than one electron.

Erwin Schrodinger was an Austrian physicist that used new theoretical calculations and experimental results to come up with a mathematical equation describing the behavior of the electron in hydrogen. The quantum mechanical model comes from the mathematical solutions to Schrodinger’s equation.
 * __The Quantum Mechanical Model pg. 130__** Thomas DeMarco

The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus. Finding the electron in a certain location is described by probability. This model description is similar to a rotating propeller. It is almost impossible to say where the blade is at a specific time. This model represents the electrons as a cloud to show how hard it is to know where the electron is. Where the cloud is less dense it is less probable to find the electron there than where it is more dense.

**__Atomic Orbitals p131-132- Nick Brault__**

//Below visual found by: Nick Brault// //http://www.youtube.com/watch?v=F-xLQ1WBIlQ//
 * Atomic orbitals are regions of space where there is a high probability of finding an electron.
 * The Schrodinger equation can be used to find energy levels, which are then used to find atomic orbitals.
 * There are seven possible orbitals with different shapes and energy levels at each principal energy level.
 * There are different shapes of orbitals, denoted by the letters p, s, d, and f.
 * S orbitals are spherical; p orbitals are dumbbell shaped; d orbitals are usually clover shaped, and f orbitals are large and complicated
 * The first principal energy level has a 1s sublevel, and the second principal energy level has 2s and2p energy levels. THis pattern continues as you move up energy levels

<span style="font-family: 'Arial Black',Gadget,sans-serif;">Group #3 pages 133- 145
 * Courtney Gareau //(co-editor)//
 * Nina DeMeo
 * Evan Grandfield
 * Steven Denison
 * Andrew Scotti

__**Electron Configurations - Pages 133-135 - By: Andrew Sciotti**__


 * Key Concept - Aufbau principle, the Pauli exclusion principle and Hund's Rule all tell tou how to find the electron configurations of atoms.**

In atoms, electrons and the nucleus interact to make the most stable arrangement possible.


 * Electron configurations**: The way electrons are arranged in orbitals (many – 4 types: S,P,F,D) around the nuclei of an atom.

__Aufbau Principle__ Electrons occupy the lowest energy orbitals first. S energy level is always the lowest energy sublevel.

The orbitials of a sublevel of an energy level are always equal energy.



__Pauli Exclusion Principle__ An atomic orbital may only describe two electrons, 1 electron or 0 electrons at any time. Example: either 1 or 2 electrons can occupy an s or p orbital at one time.

The principle also states that if two electrons occupy 1 sublevel then they must spin different directions (opposite). Spin – quantum mechanical property of electrons that states they spin up or down.



__Hund’s Rule__ This rule states that electrons occupy orbitals of the same energy so that there is a large amount of electrons with the same spin as possible.



Shorthand for electron configuration is writing the energy level and symbol for every sublevel occupied by electrons. Write number of electrons in a sublevel with superscript. *Sum of superscript equals number of electrons in an atom*

Practice – write out the way each orbital is filled H- 1s1 Li- 1s22s1 O- 1s22s22p4 Na- 1s22s22p63s1

Remember, there can be 2 electrons in 1s, 2 in 2s, there are 3 different p orbitals where each has 2 electrons and 3s can have 2 electrons.

Answers: **H** - up arrow in 1s **Li** – 1 up 1 down in 1s, 1 up in 2s. ** O ** – 1 up 1 down in 1s, 1 up 1 down in 2s, 1 up 1 down in 2p (first level) 1 up in 2p (second level) and 1 up in 2p (third level) ** Na ** – 1 up 1 down in 1s, 1 up 1 down in 2s, 1 up 1 down in 2p (first level) 1 up 1 down in 2p (second level) and 1 up 1 down in 2p (third level), 1 up in 3s

//﻿// __ Key Concept __ – Some actual electron configurations differ from those assigned using aufbau principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations.
 * __ Exceptional Electron Configurations __** – Nina DeMeo

- some elements have an electron configuration that is an exception to the aufbau principle - correct electron configurations can be obtained for elements up to vanadium ( atomic number 23) by following the aufbau diagram · however, if you continued in that fashion, some elements (i.e. copper and chromium) would be assigned incorrect configurations - exceptions to the aufbau principle are due to subtle electron-electron interactions in orbitals with similar energies

media type="youtube" key="JNPFR-22MPA?fs=1" height="385" width="480" align="center"

Nina DeMeo

- although it is worth knowing that exceptions to the aufbau princip le occur, it is more important to understand the general rules for determining electron configurations in the many cases the aufbau principle does apply

__ Practice problems: __ -Write the electron configuration for each atom: a.) carbon 1s^2 2s^2 2p^2 b.) argon 1s^2 2s^2 2p^6 3s^2 3p^6 c.) nickel 1s^2 2s^2 2p^6 3s^2 3p^6 3d^8 4s^2 -Write the electron configuration for each atom. How many unpaired electrons does each atom have?  a.) boron 1s^2 2s^2 2p^ 1 ; one unpaired electron b.) silicon 1s^2 2s^2 2p^6 3s^2 3p^ 2 ; two unpaired electrons

__Physics and Quantum Mechanical Methods - Pg. 138-141 - Steven Denison__
- A wave cycle starts at zero, increases to its highest value, and then decreases pass zero to its lowest value. - **Amplitude** is the height of a wave's crest. - **Wavelength** is the distance between crests. - **Wave Frequency** is the number of wave cycles to pass a given point per unit of time.
 * I. Light**
 * -** All light consists of waves found in different frequencies.

Wavelength and Frequency - Courtney Gareau

- This is represented by the SI unit **Hertz (Hz)**. - A **Hertz** can be represented as s -1 - The product of **Frequency** and **Wavelength** always reaches a constant (c) which is the speed of light. - **Wavelength** and **Frequency** are inversely proportional to each other meaning that as one increases the other decreases. - **Light** consists of **Electromagnetic Radiation.** - The color of **Visible Light** depends on the frequencies. - **Visible Light** can be viewed as a **Spectrum** when it passes through a prism. - **Visible Light** starts with red light having the longest wavelength and lowest frequency to violet light which has the shortest wavelength and highest frequency.
 * -** This includes **Radio Waves**, **Microwaves**, **Infrared waves**, **Visible Light**, **Ultraviolet Light**, **X-rays**, and **Gamma Rays**.

__Example Problem:__ //Calculate the Wavelength of Light// //**1. List the Knowns and Unknowns**// Knowns: -- frequency (v) = 5.10 x 10 14 /s -- c = 2.998 x 10^8 m/s

Unknowns: -- wavelength (λ) = ? m

Solve the equation c = λv for λ λ = c/v Substitute the known values and solve. λ = c/v = 2.998 x 10^8 / 5.10 x 10^14 = 5.88 x 10^-7m
 * //2. Solve for the Unknown//**

The magnitude of the frequency is much larger than the value of the speed of light so the answer should 1 and have three significant figures.
 * //3. Does this answer make sense?//**

__Practice Problems__

 * //1.//** What is the wavelength of radiation with a frequency of 1.50 x 10^13 Hz? Does this radiation have a longer or shorter wavelength than red light?


 * //2.//** What is the frequency of radiation with a wavelength of 5.00 x 10^-8m? In what region of the electromagnetic spectrum is this radiation?

II. Atomic Spectra
- Passing electric current through a gas in a neon tube energizes the electrons. - This causes the atoms in the gas to emit light. - **When atoms absorb energy, electrons move into higher energy levels, and these electrons lose energy by emitting light when they return to lower energy levels.** - Ordinary light is made up of a mixture of wavelengths. - Light emitted by atoms consists only of specific frequencies. - Each specific frequency of light corresponds to a visible color. - Each discreet line in an emission spectrum corresponds to one exact frequency of light emitted by an atom. - Each element has a different emission spectrum. - Atomic emission spectra are useful in identifying an element. - This is used to learn about the composition of stars and other things in the universe.

**__﻿﻿Physics and Quantum Mechanical Methods Continued - <span style="font-family: 'Times New Roman',serif; font-size: 16px; line-height: 18px;">Pg. 142-146-Evan Grandfield __**
__ I. An Explanation of the Atomic Spectra __ A. Atomic line spectra were known before Bohr proposed his model of the hydrogen atom. Bohr’s model explained why the emission spectrum of hydrogen consists of specific frequencies of light and also predicted specific values of these frequencies.

Atomic Spectra - Courtney Gareau

1. According to the Bohr model, the lone electron in the hydrogen atom can have only certain specific energies, the lowest energy being the ground state. a. The **ground state** is the lowest possible energy of the electron. b. When the electron is in the ground state, the electron’s principal quantum number (n) is 1. 2. When the electron is excited by absorbing energy, the electron is lifted from the ground state to the excited state with n=2, 3, and so on to larger numbers. 3. When the electron goes back to a lower energy level, a quantum of light energy is given off. a. Electronic transition is the single step in which emission occurs. 4. The quantum of energy //E// is related to the frequency v of the emitted light by this equation: E = h x v. In this equation, h is 6.626 x 10-34J x s.

B. **The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electron** (in other words the light frequency follows a logical number pattern with the energy change—it is direct as opposed to inverse).

1. Because the light emitted has a frequency directly proportional to the energy of the electron, each transition produces a line of specific frequency within the spectrum. 2. The 3 groups of lines in the hydrogen spectrum correspond to the transition electrons from higher levels to lower levels. a. The ultraviolet lines are the Lyman series; they match expected values (n=1) because of electronic transition back to ground state from excited states. b. The visible lines are the Balmer series; they result from transitions from higher energy levels to n=2. c. The infrared lines are the Paschen series; they result from transitions from higher energy levels to n=3. d. Spectral lines result from transitions from higher energy levels to n=4 and n=5. Spectral lines occur in each group and become more closely spaced at increased values of n because the levels become closer. 1. There is an upper limit to the frequency of emitted light for each set of lines because an electron that absorbs enough energy will escape the atom entirely. 3. Because it did not explain the emission spectra of atoms with more than one electron, and because it did not explain how molecules are formed, Bohr’s atom theory was not perfect. 4. The quantum mechanical model replaced the Bohr model and is based on the description of the motion of material objects as waves.

II. Quantum Mechanics A. Albert Einstein returned to Newton’s concept of particles of light and found that light could be described as quanta of energy.

1. Quanta behave like particles and light quanta are called **photons**. B. Louis de Broglie asked the question, “Given that light behaves as waves and particles, can particles of matter behave as waves?” 1. De Broglie came up with a mathematical expression for the wavelength of a moving particle. 2. Three years later, Davisson and Germer studied the bombardment of metals with beams of electrons. They noticed that the electrons reflected from the metal surfaces in patterns similar to those obtained when electromagnetic waves reflect from metal surfaces. a. The electrons, which were believed to be particles, were reflected like waves. b. The wavelike properties of beams of electrons are useful in magnifying objects. 1. In electron microscopes, the electrons have shorter wavelengths than visible light, so they can magnify even smaller things. 3. De Broglie’s equation predicted that all moving objects have wavelike behavior. 4. The object’s mass must be very small in order for the wavelength to be observable.

C. **Classical mechanics adequately describes the motions of bodies much larger than atoms, while quantum mechanics describes the motions of subatomic particles and atoms as waves.** 1. Physicist Werner Heisenberg came up with the **Heisenberg uncertainty principle** which says that it is impossible to know both the velocity and the position of a particle at the same time. 2. This principle only applies to small particles and does not pertain to things the size of cars. -Evan Grandfield

a. This principle is demonstrated when one tries to find the location of an electron. To locate an electron, you strike it with a photon of light, but the electron has so tiny a mass that the photon alters its velocity. By finding the position of an electron, its velocity is inevitably and unpredictably changed.

3. The concept of matter waves led to Schrödinger’s quantum mechanical description of electrons in atoms. His theory led to the concept of electron orbitals and configurations; it was based on the wave-motion of matter and the uncertainty principle.