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Chemistry 141 Final Review

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by: Alina Levy

Chemistry 141 Final Review Chemistry 141

Marketplace > Michigan State University > Chemistry 141 > Chemistry 141 Final Review
Alina Levy
GPA 3.0

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This is for the CLUE course with Cooper.. Although it may not be the same as other classes it still is based off of basic principals of Chemistry
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"If Alina isn't already a tutor, they should be. Haven't had any of this stuff explained to me as clearly as this was. I appreciate the help!"
Ronny Grady

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This 13 page Study Guide was uploaded by Alina Levy on Sunday December 6, 2015. The Study Guide belongs to Chemistry 141 at Michigan State University taught by in Fall 2014. Since its upload, it has received 7 views.

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If Alina isn't already a tutor, they should be. Haven't had any of this stuff explained to me as clearly as this was. I appreciate the help!

-Ronny Grady


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Date Created: 12/06/15
1.0-Chemistry  Chemistry is the study of matter and the changes that it undergoes  Changes in matter = energy changes Matter and Energy  (Most) matter is made up of atoms  Energy doesn’t have atoms in it Atoms  Atoms are the smallest unit of an element  There are 92 naturally occurring elements  Elements contain only one kind of atom  Atoms combine together to make molecules  The radius of an atom is ~ 100 pm (0.1nm) Molecules  Atoms combine together to make molecules  O is the symbol for oxygen, 02 is the symbol for oxygen molecule  Molecules can have more than one kind of element in them 1.1- Chemistry, Life, the Universe & Everything Scientific Question  Can be answered by doing experiments making observations, taking measurements, in a replicable way (ie. If someone else can’t repeat then it is not science)  You need evidence if you are going to make a claim  Must support your claim with data A Scientific Theory  Based on observations/experiments  Makes testable predictions  May change over time as new evidence becomes available  Is falsifiable (ie. Can be proven false by experiments and data)  Restricted to natural phenomena  Explains how and/or why Scientific Model  Can be a drawing or a graph or diagram or an equation, physical or mental  We use models to help us make sense- predict what would happen  They make testable, quantifiable predictions What’s the difference between Law and Theory?  Law is the description of the phenomenon while the theory explains the phenomenon Constructing a Scientific Explanation  3 main components o Claim: the claim is your conclusion o Data: the date is the evidence that is used to support the claim o Reasoning: The reasoning explains the relationship between the claim and the data 1.2- Intro to Atoms  Only 5% of the matter in the Universe appears to be made up of atoms  Cells, Air, and Gold all have atoms  A cell is bigger than atoms and molecules – atoms are the smallest Evidence of Atoms  The original idea of atoms came from the Greeks  “Atomos” – not to be cut o Not based on experimental evidence – based on philosophy o Elements: Earth, fire, air, and water  Atoms were thought to be constant motion- based on watching the movement of dust motes in sunlight- and that there was nothing or a “void” between them (later called Brownian motion from Einstin) Dalton’s Atomic Theory  Elements are composted of small indivisible, indestructible particles, called atoms.  All atoms of an element are identical and have the same mass and properties  Atoms of a given element are identical and have the same mass and properties  Compounds are formed by combinations of atoms of two or more elements  Chemical reactions are due to the rearrangement of atoms; atoms (matter) are neither created nor destroyed during a reaction. (still true) The Electron: JJ Thompson  The electron was the first subatomic particle to be discovered. Thompson’s experiment showed that:  “Particles” emerged from one disc (cathode) & moved to the other (anode)  These particles could also be deflected by magnetic fields  The particles carried the electrical charge – that is if the ray was bent, for example by a magnetic field, the charge went with it  The metal that the cathode was made of did not affect the behavior of the ray- so whatever the composition of the ray – it appeared to be independent of the element that is came from.  Atoms contain electrons “embedded” in the atom like raisins in plum pudding  Thompson discovered the electron because electrons are easier to remove from the atom  If atom was like a plum pudding, all the a particles should go straight through Ernest Rutherford  a particles that hit the nucleus were deflected and the rest passed through  Rutherford discovered the nucleus  Atom are made up of mostly empty space  Small dense positive nucleus in the center of the atom Alpha Particle  Has two protons and two neutrons Rutherford’s Planetary Model Neutrons (discovered 1932)  Hard to detect because they are neutral in charge  Slightly heavier than protons What an atom looks like Fuzzy electron cloud 1.3 – Atom Interactions  Atom is electrically neutral  Very small nucleus  Nucleus contains protons and neutrons  “cloud” of electrons- takes up most space Force Description Gravity Responsible for attraction between objects that have mass Electromagnetic Responsible for attraction between objects that have electric charge Strong Short range interactions that occur between objects made of quarks (like protons and neutrons, which are held together in the nucleus despite electrical repulsions between protons) Weak Short range intersection that occurs between elementary particles; weaker than electromagnetic or strong force. When a ball is thrown up in the air The Gravitational Force is the only force acting on the ball. When interacting objects increases, force increases. When distance between objects increase, force increases M 1 2 Gravitational Force can be Modeled by this equation: F∝ 2 r Fields mediate gravitational forces  Objects with mass are sources of gravitational fields and are affected by the gravitational fields of all other objects with mass.  Gravitational forces are always attractive. Electromagnetic force  Is mediated by electric and magnetic fields  All these forces act at a distance  Much stronger than gravity.  Can be both attractive and repulsive.  The electromagnetic force stops us from falling through the Earth, - and makes us hover over chairs (sort of) Coulomb’s Law q1q 2 F ∝ 2 r  Q1 and Q2 are the charges, and r is the distance between them  Forces between the charges can be modeled by this Energy  Any change in matter is accompanied by a change in energy  Changes in energy are caused by changes in force  Energy is conserved  System- Part of the universe you are looking at  Can monitor energy changes between system and surroundings  Energy is never lost- it can be transferred or transformed. o First Law of Thermodynamics  Two types of energy: Kinetic and potential Kinetic Energy • KE = 1/2mv 2 • Energy associated with motion • When the ball moves toward the ground, the kinetic energy increases Potential Energy  Energy associated with the position of a system of objects in a field  You can’t have potential energy without a field (gravitational, electric)  Sometimes referred to as “stored energy”  When the ball moves toward the ground, the potential energy decreases  When the ball falls to the ground, the total energy stays the same. Energy can be transferred (from one object to another) and transformed.  The distorted electron cloud creates a separation of charge 1.4 Atom Interactions London Dispersion Forces  Present between ALL molecules (neutral species)  Caused by fluctuations of electron density in the molecule (or atom)  Adjacent molecule- gets induced dipole Instantaneous Dipole – temporary difference in charge distribution within an atom    This is an attractive interaction but as they move closer- they repel! When two atoms move closer together o Total Energy- stays the same – Energy is being transferred to other atom oKinetic Energy- increases – Energy associated with motion o Potential Energy- decrease – they have less space to move  When you add thermal energy (raise the temp) o The kinetic energy of the molecules increases, causing them to move faster, collide and/or vibrate with more energy  How do the He atoms “know” the temperature is rising?  o The energy is transferred from other atoms that have collided with thte walls of the container, that were directly heated. 1.5 – Intermolecular Forces and Bonds  As the atoms move closer, the PE decreases  At the potential minimum, the system is most stable. o Attractive force=repulsive force  As the atoms move even closer, the PE increases  The depth of the well tells you how strong the interaction between the atoms is (and how much energy would be needed to get atoms out of the well) (Horizontal)  The internuclear distance tells you how far away from each other the atoms are, when they are most stable. (vertical) Van der Waals radius = ½ the distance between atoms at the potential minimum London Dispersion Forces  Increase with size of particle (number of electrons)  Increase with surface area  Part of a range of intermolecular forces (between particles)  Can be between atoms or molecules (intermolecular)  Bonds- more permanent – stronger harder to break Formation of covalent bonds  When two H atoms approach – they are attracted much more than two He atoms. – they form a covalent bond.  When the two H atoms collide- they form a new chemical species – a hydrogen molecule  When you heat H2 up the molecules move faster and will eventually separate. Breaking bonds requires an input of energy (to the system) - this can be by adding thermal energy – Increase in kinetic energy Ex: heating H to 6000k breaks the bond- which requires energy. Forming bonds releases energy (to the system) – this moves energy from the system to the surroundings Chapter 2: Light is a wave Electromagnetic Radiation  Radio waves, microwaves, infrared, ultraviolet, Xrays, Gamma Rays  High – Low frequency – Gamma Rays (y), X rays, UV, Visible, IR, Micro, Radio  High – Low wavelength- Radio, Micro, IR, visible, UV, X ray, Gamma  Light is a “wave”  Wavelength  (m) – distance from peak to peak  Frequency  (Hz – s ) number wave fronts per second  Amplitude- height of peaks (intensity)  C =   o Where c = velocity (3.00 x 10 ms for light) Determine the wavelength (in nm) of an X-ray with a frequency of 4.2 18 x 10 Hz. Properties of Waves: Interference OR Waves in phase can produce constructive interference. Properties of Waves: Diffraction Waves are bent, or diffracted, when they encounter an obstacle or slit with a size comparable to their wavelength. Thus, when a wave passes through a small opening, it spreads out. Particles, by contrast, do not diffract; they simply pass through the opening. Properties of Waves: Diffraction Patterns Energy of light  Short wavelengths (and higher frequency) are “Higher in Energy” Light is a Particle Photoelectric effect  When light shines on a metal surface electrons are emitted – creating a current  The light is transferring energy to the metal surface- where it is transformed into kinetic energy that give the electrons enough energy to “leave” the atoms in the metal  There is a threshold frequency below which no electrons are ejected – no matter how bright or intense the light is.  IF light were a wave – increasing the intensity should increase the energy- and eject electrons  Energy is transferred as a particle (photon) that has a definable energy.  E = h v  One photon ejects one electron  If the photon does not have enough energy – no electron is ejected o Einstein – photoelectric effect- doesn’t depend on intensity – just frequency E/m radiation is a particle (and a wave)  Planck E = h v for the energy of a photon (h= 6.626 x 10 -3J.s) 18 –1 o What is the energy of a photon of frequency 4.0 x 10 s o What is the wavelength of a photon of energy 6.2 x 10 –8J  Can be described as either a particle or a wave  These are models- not reality!! 2.2 – Spectra, Bohr Model, Heisenberg, Schrodinger Quantum Numbers and Orbitals Visible spectrum– light from the sun (white light), can be separated by a prism  Light from one particular element doesn’t contain the colors of the spectrum – it only has a few wavelengths  Spectra show light only of specific wavelengths/energies – the spectrum of an element is the same whether that element is on Earth, the Sun, or in the Galaxy. Niels Bohr  Explained emission and absorption spectra by involving discrete energy levels o Characterized by quantum numbers (n)  Electrons move in ORBITS around the nucleus o So- the energies of electrons in atoms are quantized  Photons of electromagnetic energy are emitted from atoms as electrons move from one energy level to another Bohr Model (only works for hydrogen) Matter is a wave (and a particle)  DeBroglie all matter has wave like properties – can calculate wavelength   = h/mv  Electrons – wavelength  about the size of atom –affects properties −31 o Note: Mass of an electron is 9.1×10 kg o h = 6.626 x 10 –34J.s o 1J = 1Kg.m /s2 2 6  Q1. What is the wavelength of an electron moving at 2.65 x 10 m/s Quantum Mechanics  Heisenberg Uncertainty Principle o Can’t measure accurately both the position and the momentum (or energy or velocity) of a small particle (electron)  Schrodinger – Quantum Mechanics 2 o Wave equation - Wave function  - Probability  Quantum Numbers and Atomic Orbitals Atomic Orbitals- Regions of high probability for finding electrons Quantum Numbers- allow us to understand the arrangement of electrons in an atom, n, l, m,land m s N is the principal quantum number, it can have values of 1, 2, 3, 4. It describes the electron shell, or energy level, of an atom L – the angular momentum quantum number, describes the subshell Ml – Magnetic quantum number, describes the specific orbital (or cloud) within that subshell Ms- The spin quantum Number, describes the spin of the electron in each orbital.  Each atomic Orbital can only contain 2 electrons  The position and energies of electrons in atoms can be describes by atomic orbitals  We can think of quantum numbers as a set of descriptors for electrons in an atom  Electrons in the same orbital have opposite spins, Ms = +1/2 or -1/2 2.3 – Quantum Numbers and Atomic Orbitals Aufbau Principle: Electrons occupy lowest level orbitals first Hund’s Rule: Electrons occupy orbitals of the same energy singly (they don’t pair up until they have to – people on a bus) Core and Valence Electrons  Core – is the last noble gas + any full d shell (transition metals) (a closed shell of electrons is very stable)  Valence- are the electrons that are higher in energy outside the closed shell Periodic Trends Atomic radius – Half the distance between two nuclei  Atomic radius increases down a group. Ionization Energy Decreases  Makes sense: heavier element s, more particles in the nucleus and more electrons  Atom radius across a row decreases. Ionization Energy: The energy required to remove an electron from an atom As atomic radius increases- Ionization Energy Decreases


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