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by: Jomary Arias

CH 1010 FINAL EXAM General Chemistry 1010

Jomary Arias
GPA 3.0

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This study guide contains material for the cumulative final exam
General Chemistry 1
Dr. Ava Kreider-Mueller
Study Guide
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This 10 page Study Guide was uploaded by Jomary Arias on Monday April 25, 2016. The Study Guide belongs to General Chemistry 1010 at Clemson University taught by Dr. Ava Kreider-Mueller in Winter 2016. Since its upload, it has received 43 views. For similar materials see General Chemistry 1 in Chemistry at Clemson University.


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Date Created: 04/25/16
CH 1010: General Chemistry Dr. Ava Kreider-Mueller FINAL EXAM CHAPTER 1: Key Terms-  Steps of the Scientific Method: 1. Ask a Question 2. Background Research on the topic 3. Create a Hypothesis 4. Test your Hypothesis 5. Analyze your data, and draw a conclusion 6. Report all findings, check if hypothesis was correct  Hypothesis: a testable explanation  Scientific Theory: a explanation which has been repeated tested  Scientific Law: a statement of a fundamental principle that can be applied to a topic.  Meter: distance traveled by light emitted from a helium-neon laser through a vacuum  Matter: anything that has a mass and occupies space o 3 States of Matter are: 1. Solid 2. Liquid 3. Gas  Weight: measure of force that gravity exerts on an object. Weight can change since gravity can change but mass stays the same  Mass: quantity of matter in an object 3 Most Common Units of Temperature: 1. Fahrenheit, Fº 2. Celsius, Cº 3. Kelvin, Kº [Both Celsius and Kelvin have the same temperature interval] If you are finding Kº use the following equation: Kº = Temp. in Cº + 273.15 If you are finding Cº use the following equation: Cº = Temp. in Kº - 273.15  Dimensional-analysis method: used to convert from one unit to the next 3  Volume: amount of space occupied by an object (m ) CH 1010: General Chemistry Dr. Ava Kreider-Mueller FINAL EXAM  Density: mass of a substance per unit volume. It is temperature dependent. Use the following formula: Density = Mass (g) / Volume (mL or cm ) 3  Joule (J): units used to measure energy  Accuracy: when an experimental value and true value match  Precision: when repeated measurements of the same variable agree  Significant Figures: total number of digits recorded for a measurement. It is used to indicate uncertainty in a measurement. The rule is to use all the digits you know to be true plus one extra digit you can estimate  Atom: smallest particle of an element  Molecules: a collection of atoms chemically bonded together  Chemical Bond: a force that hold 2 atoms in a molecule together  Pure Substance: matter CAN NOT be separated into simpler matter by a PHYSICAL PROCESS  Mixture: combination of pure substance in different proportions. CAN BE SEPERATED FROM ONE ANOTHER  Physical Process: a transformation of a sample of matter that does not alter the chemical identity of any substance in the sample  Element: pure substance, CAN NOT be separated into smaller particles  Law of Constant Composition: principle that ALL samples of a particular compound always contains the same elements combined in the same proportions Intensive Property Extensive Property -Temp. -Length -Boiling point -Volume -Density -Mass -Color -Energy -Concentration -Weight 4 Rules to Determine a Significant Figure: 1. Zeros in the MIDDLE of a number are significant 2. Zeros at the BEGINNING of the number are NOT significant 3. Zeros at the END of a number and AFTER a decimal point are significant 4. Zeros at the END of the number and BEFORE decimal point MAY or MAY NOT be significant. When Multiplying & Dividing: - Final answer CAN NOT have more significant figures than either of the original numbers CH 1010: General Chemistry Dr. Ava Kreider-Mueller FINAL EXAM When Addition & Subtraction: - Final answer CAN NOT have more digits to the RIGHT of the decimal point than either original numbers. CHAPTER 2:  Protons (+): charge that is equal in magnitude to that of an electron  Neutrons: No charge  Electrons (-) : Negative charge atom  Nucleus: small central core of the atom, contains protons and neutrons  Reactants: compound undergoing change  Product: compound generated in the reaction  Atomic Number (Z): number of protons in an atom’s nucleus  Mass Number (A): number of protons + number of neutrons  Isotopes: atoms with identical atomic numbers but DIFFERENT mass numbers  Atomic Mass: mass of a specific atom  Atomic Weight: weighted average of atomic masses of the element’s naturally occurring isotopes. To find the Atomic weight use the following formula: (Mass of Element)(Abundance of Element) + (Mass of Element)(Abundance of Element)  Molecular Mass: the mass in amu of one molecule of a molecular compound  Ion: an atom or molecule that has a (+) or (-) charge  Ionic bond: results from the complete transfer of 1 or more electrons from 1 atom to another. o Cations: (+) charged ions o Anions: ( - ) charged ions CHAPTER 3 Electromagnetic Waves: Wavelength Amplitude (A) Frequency (v) Intensity of radiant (ʎ ) energy Distance from one Height of the wave, # of wave peaks that Proportional to A^2. peak to the next. measured form the pass a given point per Measured in Units = middle of the peak unit time. Units = s^-1 meters or Hz. CH 1010: General Chemistry Dr. Ava Kreider-Mueller FINAL EXAM How To Measure the Speed of a Wave: Use the following equation: Wavelength * Frequency = Speed ʎ (m) * v ( s^-1) = c (m/s)  Speed of Light- is noted as ( c ) and is basically the rate of travel of all radiant energy in a vacuum. - This value is always a constant value of ( 2.998 * 10^8 m/s) - All electromagnetic radiation moves at the same speed. How to Calculate Wavelength: Use the following equation: c / v# = ʎ# In this equation you would divide the speed of light ( 2.998 * 10^8 m/s) / by the given frequency  to give you the wavelength. How to Calculate Frequency: C / ʎ# = v# In the equation you would divide the speed of light/ by the given wavelength to get  Frequency. - Frequency and Wavelength are INVERSELY related. o Longer wavelength = Lower Frequency o Shorter Wavelength = Higher Frequency Electromagnetic Spectrum: continuous range of wavelengths and frequencies. CH 1010: General Chemistry Dr. Ava Kreider-Mueller FINAL EXAM Gamma Rays < X Rays < Ultraviolet < Visible Light < Infrared < Microwave < Radio (Shortest Wavelength) (Longest Wavelength) (Highest energy) (Lowest Energy)  Light appears in the form of wavelengths ( ʎ ). A source can emit a single wavelength such as a laser. - Monochromatic radiation: contained only a single wavelength - Radiation can also contain many different wavelengths. Ex) Light bulbs - Spectrum: is produced when radiation from a source is separated into its different wavelengths.  White Light- continuous distribution of wavelength spanning entire spectrum. As a beam of white light passes through a prism, the wavelengths separate into different component color. - Different wavelengths travel through glass at different rates Ex) Red Light- travels at 780 nm, long wavelength Violet Light- travels at 380 nm, short wavelength CHAPTER 4 Key Terms:  Ionic Bond- results from the complete transfer of 1 or more electrons from 1 atom to another, the formation of 2 charged particles (cations and anions)  Ionic Solid- cations & anions pack together in a regular way  Electrostatic Potential Energy (E )elenergy a charged particle has because of its position relative to another charged particle.  Lattice Energy ( U) – energy released when 1 mole of an ionic compound forms from its free ions in the gas phase. -As Radius of a Cation ↓ the lattice energy ↑ -As the Radius of an Anion ↓ the lattice energy ↑ -As the Charge of the Cation ↑ the lattice energy ↑  Polyatomic ions- charged, covalently bonded groups of atoms  Covalent Bond- two atoms share 2 electrons  Molecule- a unit of matter that results when 2 or more atoms are joined by covalent bonds  Bond length- the optimum distance between nuclei, where net attractive forces are maximized and the molecule is stable (lowest energy formation)  Nonpolar covalent bond: a bond characterized by an even distribution of charge; electrons in the bonds shared equally by the two atoms CH 1010: General Chemistry Dr. Ava Kreider-Mueller FINAL EXAM  Polar covalent bond: bond resulting from an unequal distribution of bonding pairs of electrons between atoms, results in a partial charge on each of the atoms  Electronegativity: relative measure of the ability of an atom in a bond to attract electrons to itself when bonded to another atom.( ↑ as you move UP a group)( ↑ as you move across a period or from left to right) Nonpolar Covalent Polar Covalent Ionic Compound Bonds between atoms Bonds between atoms whose Bonds between atoms whose with the SAME or similar electronegativity differ by less electro negativities differ by electronegativity (∆x≤ than 2 units (0.4< ∆x< 2.0) more than 2 units 0.4) (∆ x 2.0 )  Chemical Formula: lists the symbols of the elements involved in a given compound; using subscripts to indicate the # of atoms of each element.  Structural Formula: shows the bonds between atoms; tells us how the atoms are connected and gives more info than a chemical formula  Molecular model: give a 3D representation of the molecule  Octet Rule: Atoms of main group elements make bonds by gaining, losing, or sharing electrons to achieve a valence shell containing 8 electron, or four electron pairs, can be expanded.  Lewis Structure: 2D representation of the bonds and lone pairs of valence electrons in an ionic or molecular compound  Single bond: bond that results when 2 atoms share 1 pair of electrons  Lone pair: a pair of electrons that is not shared  Bonding pair: a pair of electrons shared between 2 atoms  Resonance: characteristic of electron distributions when 2 or more equivalent Lewis structures can be drawn for one compound  Resonance Structure: one of two or more Lewis structures with the same arrangement of atoms but different arrangements of bonding pairs of electrons. - Only differ in placement of the valence shell electrons - Connections b/w atoms remain the same - Pairs of electrons that are spread out among atoms are said to be delocalized  Resonance Stabilization: stability of a molecular structure due to delocalization of its electrons - May occur in polyatomic ions - Polyatomic Anions ADD the appropriate # of valence electrons - Polyatomic Cations SUBTRACT the appropriate # of valence electrons Main Group Metals: CH 1010: General Chemistry Dr. Ava Kreider-Mueller FINAL EXAM Cations ( + ) Anion ( -) Group 1A elements form 1 + cation Group 6A elements form 2- anions (6-8 = -2) Group 2A elements forms 2 + cations Group 7A elements form 1 – anion (7-8 = -1) Group 3A elements form 3 + cation Group 8A elements DON’T form ions (8-8 = 0) Melting an Ionic Solid: - Requires energy, ions begin to move with more freedom - Cations and Anions separate - Compound will change phases from solid  liquid - As lattice energy ↑ the melting point ↑ Ionic Bond Covalent Bond - High melting solids - Low melting solids, liquid, or - Vast 3D network of ions gasses - Opposite charges attract - Strong, attracts forces between different molecules is weak - Little energy required to overcome forces between molecules How to Draw Electron-Dot Structure: - 1. Determine the # of valance e 2. Arrange the symbols of the elements in a pattern that shows how their atoms are boned together & then connect them with single bonds. Place elements with large bonding capacity in the center and terminal atoms surrounding it. Electron Dot Structures: - 1 line indicates  single covalent bond (2 shared electrons) - 2 lines indicate double bond ( 4 shared electrons) - 3 lines indicate  triple bond ( 6 shared electrons) (Multiple bonds are shorter & stronger than single bonds because more shared electrons are holding them together) Ex) Cl 2,ngle bond CH 1010: General Chemistry Dr. Ava Kreider-Mueller FINAL EXAM Each atom in Cl m2lecule gains a noble-gas configuration with 8 valence electrons, so it obeys the octet rule CHAPTER 6 Terms Topics [Key Terms]  Gases - particles are independent of one another, feel little attractive force, and are free to move about randomly.  Liquids- particles are strongly held together by attractive forces which are strong enough to hold the particles in close contact while letting them slide over one another.  Solids- the particles are held by attractive forces which hold the particles in place.  Intermolecular forces- act between molecules to hold them together at certain temperatures  Van der Waals forces- several different types of intermolecular forces, including dipole-dipole forces, London dispersion forces, & hydrogen bonds. Contain partial charges.  Ion-dipole forces- act b/w ions and molecules  Net Force- sum of many individual interactions  Viscosity- measure of a liquid’s resistance to flow  Surface Tension- resistance of a liquid to spread out & increase its surface area  Solvent- component of solution that is present in the larger # of moles. Ex) water  Solute- component in a solution other than the solvent. A solution may contain one or more solutes. Ex) NaCl  Solubility- maximum quantity of a substance that can dissolve in a given volume of solution [Ion- Dipole Force] - NOT one of the Van der Waals forces - Result from interactions b/w an ion & the polar charges of a polar molecule - Contain full charges and the ion-ion attraction is so strong that they create an ionic bon [3 Types of Intermolecular Forces] 1. London Dispersion 2. Dipole-Dipole Forces 3. Hydrogen Bonds [London Dispersion] CH 1010: General Chemistry Dr. Ava Kreider-Mueller FINAL EXAM - Present in ALL atoms & molecules, regardless of structure - WEAK attractive forces - Magnitude of force depends on the ease with which a molecule’s e- cloud can be distorted by a nearby electric field - Temporary dipoles are created as the nuclei & electron clouds interact - Shape contributes to the strength - More spread-out shape (longer Hydrocarbon chains) allow greater contract b/w molecules & give rise to higher dispersion forces - Chain length ↑, the Boiling Point ↑ - The LESS compact the molecule, HIGHER the boiling point. Chains vs. Spherical shapes allow greater contact b/w molecules and give rise to higher dispersion forces Lighter atoms Heavier atoms -Less polarizable -More polarizable - Contains only a few tightly held electrons -Contains many electrons, some less tightly held -Smaller dispersion forces and farther from the nucleus -Smaller molecules -Larger dispersion forces -Larger dispersion forces [Dipole-Dipole Forces] -Experienced by NEUTRAL, but polar molecules b/c of electrical interactions among dipoles of neighboring molecules -Forces can be ATTRACTIVE or REPULSIVE, depending on orientation -Polar molecules ATTRACT, when they orient w/ UNLIKE charges close together -Polar molecules REPEL, when they orient w/ LIKE charges close together -Net force is the sum of many individual interactions. Strength depends on the size of the dipole moments involved. The more polar the substance, the greater the strength of its dipole-dipole interaction - The larger the dipole moment-> the stronger the intermolecular forces  the greater the boiling point - WEAKER than ion-dipole forces [Hydrogen Bonds]  Attractive interaction b/w hydrogen atoms bonded to a very electronegative atom ex) O, N, F and an electron- rich region elsewhere in the same molecules or in a different molecules  O-H, N-H, and F-H are highly polar with a partial (+) charge on the hydrogen, and has a partial (-) charge on the electronegative atom CH 1010: General Chemistry Dr. Ava Kreider-Mueller FINAL EXAM  Hydrogen has NO CORE e- to shield the nucleus, & it has a small size so it can be approached closely by other molecules  Dipole-dipole interaction b/w hydrogen & unshared electron pair on a nearby atom is unusually strong  Is the primary intermolecular force that holds large molecules together  Boiling point ↑ w/ molecular weight or as you move down a group, exceptions: NH 3, H2O, and HF [Viscosity]  Ease in which individual molecules move around when intermolecular forces are present  Temperature dependent  Small nonpolar molecules experience only weak intermolecular forces, & have low viscosities  Larger polar substances have stronger intermolecular forces, & have higher viscosities  Stronger intermolecular forces have higher boiling and melting points [Surface Tension]  Resistance of a liquid to spread out & increase its surface area, caused by the difference in intermolecular forces  Temperature dependent  Molecules at the surface experience attractive forces only on one side  Molecules in the interior are surrounded and are pulled equally in all directions  Surface tension is higher in liquids that have stronger intermolecular forces. Ex) Mercury Miscible Immiscible Liquids that are mutually soluble in any Liquids that have limited solubility in each other, proportion, completely dissolves won’t dissolve “LIKE dissolved LIKE”  POLAR solutes tend to dissolve in POLAR solvents  NON-POLAR solutes tend to dissolve in NON-POLAR solvents


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