Ch.+4+and+5,+The+Atom+(Period+A)

__Chapters 4 and 5: THE ATOM__ Editor: Ali Fortier, Period A __ Introduction: __ Chapters 4 and 5 are all about the atom. Chapter 4 explains what an atom is, along with what’s inside of an atom and how that works. Several experiments, ideas and theories are also explained. In Chapter 5, we are presented with different atomic models. Chapter 5 also explains atomic orbitals and electrons in atoms. **Chapter 4: Atomic Structure**

__ **4.1 Defining the Atom** __
**Co-Editor: Catherine Murray** **Member: Kyle Gallagher**

Scientists use many devices to enhance their ability to see, but they can't always see everything. Scientists then must obtain experimental data to fill in the missing parts
 * Connecting to your world:**




 * Early Models of the Atom**
 * __Atom:__ the smallest particle of an element that retains its identity in a chemical reaction.
 * We cannot observe individual atoms, but we can propose ideas about the atoms


 * Democritus's Atomic Philosophy**
 * Greek philiosopher Democritus (460 BC-370 BC) was the first to propose the existance of atoms. He believed they were invisible and indestructible
 * His theory lacked experimental support and did not explain chemical behavior, but agreed with later ideas.

** Dalton’s Atomic Theory **
 * The modern process of discovery regarding atoms began with John Dalton, an English chemist and school teacher, more than 2000 years after Democritus
 * By using experimental methods, Dalton transformed Democritus’s ideas on atoms into a scientific theory
 * He studied the ratios in which elements combine in chemical reactions
 * Based on results of experiments, made hypotheses and theories to explain observations, resulting in Dalton’s Atomic Theory
 * 1) All elements are composed of tiny indivisible particles called atoms.
 * 2) Atoms of the same element are identical. The atoms of any one element are different from those of any other element. (shown by parts "a" and "b" in the picture below)
 * 3) Atoms of different elements can physically mix together or can chemically combine in simple whole-number ratios to form compounds. (shown by parts "c" and "d" in the picture below)
 * 4) Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction.



** Sizing up the Atom **
 * A coin the size of a penny and made of pure copper (Cu) shows Dalton’s concept of the atom
 * Grind the coin into a fine dust and each speck of the small pile of dust would still have the properties of copper
 * If you continued to make the particles smaller, you would get to a particle that could not be divided while still having the chemical properties of copper
 * This is an atom
 * There are about 2.4 x 10^22 atoms in a pure copper coin the size of a penny
 * Earth’s population is about 6 x 10^6 people
 * There are about 4 x 10^12 times as many atoms in the coin as there are people on Earth
 * If you line up 100,000,000 copper atoms side by side, they would produce a line only 1 cm long
 * The radii of most atoms are usually 5 x 10^-11 m to 2 x 10^-10 m
 * Despite their small size, individual atoms are observable with instruments such as scanning tunneling microscopes.
 * Individual atoms can even be moved around and arranged in patterns, holding future promise for creation of atomic-sized electronic devices (atomic-scale or “nanoscale”)
 * Could become essential to future applications in medicine, communications, solar energy, and space exploration



__ **4.2 Structure of the Nuclear Atom** __
**Co-Editor: Maddie Myers** **Member: Brynna Harum** - scientists questioned how subatomic particles were put together in an atom
 * The Atomic Nucleus**
 * most thought that electrons were evenly distributed throughout an atom
 * including __JJ Thomson__(discoverer of the electron)
 * "plum-pudding model"
 * electrons were stuck into a lump of positive charge, like raisins stuck in dough
 * short lived due to groundbreaking discoveries by his student, __Ernest Rutherford__


 * Rutherford's Gold-Foil Experiment**
 * decided to test the current theory of atomic structure
 * used massive alpha particles
 * helium atoms that have lost two electrons and have a double positive charge
 * a narrow beam of them was directed at a thin sheet of gold foil
 * should have passed easily through the gold, with a slight deflection
 * to everyones surprise, the majority of the alpha particles passed straight through the gold atoms, w/o deflection
 * a small fraction of them bounced off the gold foil at large angles


 * The Rutherford Atomic Model**
 * suggested a new theory of the atom
 * the atom is mostly empty space
 * which explains the lack of deflection of most of the alpha particles
 * all positive charge, and almost all of the mass are concentrated in a small region
 * called the __nucleus__
 * the tiny central core of an atom and is composed of protons and neutrons
 * known as the nuclear atom
 * in it, the protons and neutrons are located in the nucleus. The electrons are distributed around the nucleus and occupy almost all the volume of the atom.
 * nucleus is tiny in comparison to whole atom (football stadium vs. marble)

__** 4.3 Distinguishing Among Atoms **__

 * Co-Editor: Nikki Steiner **
 * Member: Ally Luongo **
 * Member: Lauren O'Reilly **

- atoms are composed of protons, electrons and neutrons - protons and neutrons make up the nucleus - electrons surrond the nucleus - elements are different because they have different numbers of protons - atomic number is the number of protons in the nucleus of an atom of that element - atomic number identifies an element example: hydrogen atoms have 1 proton so the atomic number of hydrogen is 1
 * Atomic Number**
 * oxygen atoms have 8 protons so the atomic number of oxygen is 8



- most of the mass of an atom is concentrated in its nucleus and depends on the number of protons and electrons - mass number: the total number of protons and neutrons in an atom - if you know the atomic number and mass number of an atom of any element, you can determine its compostition example: oxygen has an atomic number of 8 and a mass number of 16 atomic number = number of protons (and also the number of electrons in a neutral atom) so you know oxygen has 8 protons and 8 electrons mass number is equal to the number of protons + number of neutrons the number of neutrons in an atom= mass number - atomic number - the compostitoin of anu atom can be represented in shorthand notation by using atomic number and mass number (pg. 112-114)
 * Mass Number**


 * Isotopes**
 * there are three different kinds of neon atoms
 * all have 10 protons and 10 electrons but each have a different numberof neutrons
 * Isotopes are atoms that have the samenumberof protons but different numbers of neutrons
 * because isotopes have different numbers of neutrons, they also have different mass numbers.
 * isotopes are chemically alike because they have the same number of protons and electrons
 * subatomic particles responsible for chemical behavior
 * Three known isotopes of hydrogen
 * each isotope of hydrogen has one proton in its nucleus
 * most common hydrogen isotope has no neutrons
 * has a mass number of 1 and is called hydrogen -1, or simply just hydrogen
 * the second isotope has one neutron and a mass number of 2, called either hydrogen -2 or deuterium
 * third isotope has two neutrons ans a mass number of 3
 * called hydrogen -3 or tritum



(pg. 115-116) - most elements occur as a mixture of two or more isotopes in nature - each has a fixed mass and natural percent abundance - example: chlorine-35 and chlorine-37 - atomic mass - a wighed average mass of the atoms in a naturally occurring sample of the element - knowing that the atomic mass is a wighted average od the masses of its isotopes, you can determine the atomic mass based on relative abundance
 * Atomic Mass**
 * ======actual mass of a proton or a neutron is 1.67 x 10 to the negative 24 grams======
 * mass of an electron is 9.11 x 10 to the negative 28 g
 * largest atom is incredibly small
 * able to determine mass of smaller things by using a mass spectrometer
 * isotope of carbon was assigned a mass of exactly12 atomic massunits
 * atomic mass unit(amu): one twelfth of the mass of a carbon-12.
 * a carbon- 12 atom has six protons and six neutrons in its nucleus and its mass is set at 12 amu
 * the six protons and six neutrons make up nearly all of the mass
 * the mean of these two masses would be 35.968 amu but this is higher than the actual value because you need to take into consideration that chlorine-35 accounts for 75% of natural chlorine atoms and chlorine-37 is 25%
 * a weighted average effects both mass and the relative abundance of the isotopes as they occur in nature
 * 3 values must be known in order to do this
 * the number of stable isotopes of the element
 * the mass of each isotope
 * the natural percent abundance of each isotope
 * 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
 * resulting sum is the weighted average mass of the atoms of the element as they occur in nature

This website is helpful for calculating the atomic mass based on relative abundance: []

Here are two examples for calculating the atomic mass: Bromine has two isotopes, Br-79 and Br-81. Both exist in equal amounts. Calculate the relative atomic mass of bromine.
 * //1) Example://**
 * //Solution://**

The neon element has three isotopes. They are 90.92% of, 0.26% of and 8.82% of
 * //2) Example://**
 * //Solution://**

**Chapter 5: Electrons In Atoms**

__** 5.1 Models of the Atom **__

 * Co-Editor: Jordyn Renaghan **
 * Member: Jordan Majka **
 * Member: Andrea Vale **


 * The Development of Atomic Models**
 * After discovering the atomic nucleus,Rutherfordused existing ideas about the atom and proposed an atomic model in which the electrons move around the nucleus.
 * This model explained only a few simple properties of atoms. It could not explain the chemical properties of elements; this requires a model that better describes the behavior of electrons within atoms
 * The Bohr Model**
 * Niels Bohr; Danish physicist
 * BelievedRutherford’s model needed improvement; in 1913, Bohr changedRutherford's model to include newer discoveries about how the energy of an atom changes when it absorbs or emits light.
 * Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus
 * Energy levels – The fixed energies an electron can have
 * Energy levels are like the rungs of a ladder: The lowest rung of the ladder is like the lowest energy level. Someone can climb up or down a ladder by going from rung to rung, and an electron can jump from one energy level to another in the same way. A person cannot stand on a ladder between the rungs, and an electron cannot be between energy levels. To move from one rung to another, a person must move just the right distance – to move from one energy level to another, an electron must gain or lose just the right amount of energy. The higher the electron is on the energy ladder, the farther it is from the nucleus.
 * Quantum – The amount of energy required to move an electron from one energy level to another energy level
 * The amount of energy an electron loses or gains is not always the same
 * The energy levels in an atom are not equally spaced; higher energy levels are closer together, so it takes less energy to climb from one rung to another near the top of the ladder.
 * The higher the energy level occupied by an electron, the less energy it takes to move from that energy level to the next higher energy level.



In 1926, an Austrian physicist named Erwin Schrödinger used results from the Bohr and Rutherford planetary models to explain the movement of electons in an atom. The quantum mechanical model comes from mathematical solutions to Schrödinger's equations.
 * The Quantum Mechanical Model**
 * This model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus.
 * How likely it is to find an electron is determined by probability.
 * This model describes the movement of electrons to be the same of a rotating propellor blade. The model shows that electrons are more likely to be found closer to the nucleus. Picture the model as being able to mold a sack around the cloud of electrons outside the nucleus so that the electrons were in the sac 90% of the time.



Solving the Schrödinger equation gives the energies an electron can have.
 * Atomic Orbitals**
 * Called energy levels
 * the Schrödinger equation gives each energy level a mathematical expression
 * known as an atomic orbital
 * Atomic Orbital: the region of space in which the is a high probability of finding an electron
 * for each principle energy level, there may be several orbitals with different shapes and at different energy levels
 * constitute energy sublevels
 * EACH ENERGY SUBLEVEL CORRESPONDS TO AN ORBITAL OF A DIFFERENT SHAPE, WHICH DESCRIBES WHERE THE ELECTRONS ARE LIKELY TO BE FOUND
 * different atomic orbitals are denoted by letters: s are spherical, p are dumbell-shaped, d are leaf shaped, f are more complicated
 * S orbitals: because of sphere shape, the probability of finding an electron at any given distance from the nucleus does not depend on the direction
 * P orbitals: 3 kinds, that have different orientations in spaced




 * The numbers and kinds of atomic orbitals depend on the enerdy sublevel
 * lowest principle energy level has one sublevel (n=1) 1s.
 * second principle energy level has two sublevels (n=2) 2s and 2p
 * 2p sublevel is of higher nenery and consists of 3p orbitals of equal energy
 * third principle energy level has three sublevels (n=3) 3s, 3p, 3d
 * consists of 5d orbitals of equal energy
 * fourth principle energy level has four sublevels (n=4) 4s, 4p, 4d, 4f
 * consists of 7f orbitals of 18 energy

__** 5.2 Electron Arrangement in Atoms **__

 * Co-Editor: Abby White **
 * Member: Ryan McSweeney **
 * Member: Billy Arruda **


 * Electron Configurations**
 * In the atom, the nucleus and electrons interact to make the most stable arrangement possible
 * __Electron Configuration:__The ways in which electrons are arranged in various orbitals around the nuclei of atoms
 * This can be determined through:
 * Aufba Principle
 * Pauli Exclusion Principle
 * Hunds Rule
 * Aufbau Principle:
 * Electrons occupy the orbitals of lowest energy first
 * The orbitals for any sub level of a principal energy are always of equal energy
 * The "s" sub level is always the lowest energy sub level
 * The filling of atomic orbitals does not follow a certain pattern
 * Pauli Exclusion Principle:
 * An atomic orbital may describe at most 2 electrons
 * Either one of the two electrons can occupy an "s" or "p" orbital
 * In Order to occupy the same orbital, two electrons must have opposite spins
 * The elections spins must be paired
 * Spin is a quantum mechanical property
 * A vertical arrow indicates an electron and its direction of spin
 * Hund's Rule:
 * Electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible
 * Second electrons occupy each orbital so that their spins are paired with the first electron in the orbital
 * As a result, each orbital can eventually have two electrons with paired spins

As shown in the diagram below, an oxygen atom contains 8 electrons. -The orbital of the lowest energy, 1s, has one electron, then a second electron of the opposite spin. -The next orbital to fill is 2s. It also has one electron then a second electron of opposite spins. -One electron then occupies each of the three 2p orbitals of equal energy. -The remaining electron now pairs with an electron occupying one of the 2p orbitals. The other 2p orbitals remain only half filled, with one electron each. -A shorthand for writing an electron configuration involves writing the energy level and the symbol for every sublevel occupied by an electron. -The number of electrons is written in a superscript. For example: hydrogen with one electron is 1s^1. -Example: for oxygen with two electrons in a 1s orbital, two elctrons in a 2sx orbital, and four elctrons in a 2p orbital it is written 1s^2 2s^2 2p^4. -When the configurations are written the sublevels within the same principal energy are generally written together. This is not always the same as in the aufbau diagram. For example the 3d sublevel is written before the 4s sublevel. __**This is a standard electron configuration pattern below.**__

- ex: chromium has a half-filled //d// sublevel. - Filled energy sublevels are more stable than partially filled sublevels.
 * Exceptional Electron Configurations**
 * You may obtain correct orbital configurations up to vanadium.
 * Some actual electron configurations differ from those assigned using the aufbau principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations.
 * The exceptions are due to subtle electron-electron interactions in orbitals with very similar energies.

**__ 5.3 Physics and the Quantum Mechanical Model __**

 * Co-Editor: Grant Casey **
 * Member: Ian Kelly **
 * Member: Brittany Morgan **
 * Light**
 * the quantum mechanical model grew out of the study of light
 * light consists of waves
 * the **amplitude**of a wave is the wave's height from zero to the crest
 * the **wavelength**, represented by the Greek letter lambda ([[image:http://www.clker.com/cliparts/8/4/e/3/1206565371256380727Anonymous_lambda.svg.med.png width="15" height="29"]]), is the distance between the crests
 * the **frequency,** represented by the Greek letter nu ([[image:http://upload.wikimedia.org/wikipedia/commons/thumb/4/41/Greek_lc_nu.svg/400px-Greek_lc_nu.svg.png width="147" height="39"]]), is the number of wave cycles to pass a given point per unit of time


 * the units of frequency are usually cycles per second, in SI called a __hertz__.
 * According to the wave model, light consists of electromagnetic waves
 * __Electromagnetic radiation__includes radio waves, visible light, UV waves, X-rays, and gamma rays
 * sunlight consists of light with a continuous range of wavelengths and frequencies
 * the color of light for each frequency found in sunlight depends on its frequency
 * when sunlight passes through a prism, the different frequencies separate into a __spectrum__of colors
 * a rainbow is an example of this phenomenon
 * RED light has the lowest frequency and longest wavelength, whereas VIOLET has the highest frequency and shortest wavelength in visible light.


 * Atomic Spectra**
 * 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.
 * When an electric current is passed through a gas in a neon tube, it energizes the electrons of the atoms of the gas and causes them to emit light.
 * Ordinary light is made up of a mixture of all forms of light.
 * When light passes through a prism, the frequencies of light emitted by an element separate into discrete lines to give an elements __Atomic Emission Spectrum__.
 * Every line of the emission spectrum correlates to one frequency of light emitted by the atom
 * Each element has its own unique emission spectrum.
 * This helps us to understand the makeup of stars and other bodies of the universe.
 * An Explanation of Atomic Spectra**
 * The Bohr model helped to explain the emission spectrum of Hydrogen and can be applied to further prediction.[[image:http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/graphics/hydrogen.gif caption="Atomic emission spectrum of Hydrogen"]]
 * The lowest possible energy state of an electron is its **Ground State.**
 * In this state the electron's principal quantum number is 1.
 * When it absorbs energy it raises from a ground state to an excited state.
 * Light is formed when the electron drops a quanta in a single abrupt step called an electronic transition.
 * E= h x v
 * The light emitted by an electron moving from a higher energy level to a lower energy level has a frequency directly proportional to the energy change of the electron.
 * However, Bohr's theory of the atom was unhelpful regarding atoms with more than one electron. media type="youtube" key="NdMkEmRWqs8" width="425" height="350"


 * Quantum Mechanics**


 * In 1905 Albert Einstein re-examined Newton’s idea of particles of light. He successfully explained experimental data by proposing that light could be described as quanta of energy. Quantum Mechanics
 * Light quanta are called __photons__.
 * In 1924 Louis de Broglie referred to the wavelike behavior of particles as matter waves.
 * He then reached a mathematical expression for the wavelength of a moving particle.
 * Experiments confirmed the validity of this theory.
 * The actual experiment was done by Clinton Davisson and Lester Germer.
 * They experimented on how electrons when reflected from metal flew in wave-like patterns.
 * De Broglie and Davisson received Nobel Prizes for their work.
 * Today wavelike properties are used in magnifying images.
 * The smaller an object the greater it’s wavelength; and the easier it is to observe the wavelength.
 * __Quantum Mechanics__ describes the motions of subatomic particles and atoms as waves.
 * Werner Heisenberg created the __Heisenberg uncertainty principle__and it stated that it is impossible to know exactly the velocity and position of a particle at the same time.
 * This principle applies to small particles not normal sized objects.
 * Predicting a location and velocity is impossible because trying to find the position of an electron alters its velocity making its velocity uncertain.
 * This discovery of waves led the way for Schrodinger’s quantum mechanical description of electrons in atoms.
 * His theory led to the concept of electron orbitals and configurations and it includes the wavelike motions of matter and the uncertainty principle.