Chemistry Exam 1 Review Guide
Chemistry Exam 1 Review Guide Chem 1180-002
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This 7 page Study Guide was uploaded by Kaelin Kneen on Sunday September 18, 2016. The Study Guide belongs to Chem 1180-002 at University of Nebraska at Omaha taught by Dr. Alan Gift in Fall 2016. Since its upload, it has received 193 views. For similar materials see General Chemistry I in Chemistry at University of Nebraska at Omaha.
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Date Created: 09/18/16
Chem 1180: Exam 1 Review Chapters 1, 2, and 6 Review Concepts Chapter 1- Definitions o Chemistry- study of properties and transformations of matter o Matter- things that have mass and take up space o Experiment- observation of natural phenomena o Hypothesis- tentative explanation of nature o Theory- tested explanation of natural phenomena o Law- fundamental theory Changes in Matter o Chemical vs. Physical Chemical- transformation into new matter Physical- change in form, not matter Properties of Matter o Physical- does not change the identity of substance when measured (intensive-independent vs. extensive- dependent) Ex. Color or Volume o Chemical- may only be observed when changing the chemical identity of a substance. Ex. Flammability or Toxicity Law of Conservation of Mass- o States that the mass of the products in a chemical reaction must be equal to the mass of the reactants. o Mass is an isolated system that is neither destroyed or created by chemical reactions or physical transformations 3 Types of Matter o Element- a substance that cannot be composed into simpler substances by chemical means, and is made up of atoms all with identical number of protons o Compound- two or more chemical elements, chemically bonded together o Mixture- two substances combined (homogeneous or heterogeneous) Temperature Conversions- o K=C˚+273.15 (Kelvin and Celsius are stationary) o C˚=5/9(F˚-32) o F=9/5(C˚+32) Precision vs. Accuracy- o Precision- how close measurements are to each other o Accuracy= how close measurements are to actual value Significant Figures- numbers of digits used to express value o Ex. 5.05 (3 sig figs) 0.08 (1 sig fig) 6.000 (4 sig figs) 430 (2 sig figs) o For (x) and (/) always do least numbers of sig figs. Ex- 3.56*4.3= 15 o For (+) and (-) sig figs are set by uncertainty of coarsest measurement. Ex- 67.678+12.3= 79.9 SI Prefixes- Density= mass/volume Dimensional Analysis Chapter 2 Parts of an atom- o Proton- positive electric charge (mass of approx. 1amu) o Neutron- neutral electric charge o Electron- negative electric charge REMEMBER! o One Angstrom= 1x10^-0 m How to Read an Elements Properties- Definitions- o Isotope- two or more of the same element that contain the same number of protons and different numbers of neutrons o Nuclide- a distinct kind of atom, characterized by the number of protons and neutrons. o 1 amu = 1 Dalton o Atomic Mass- protons + neutrons o Molecular Mass- Script of number x atomic mass of element Periodic Table Groups Molecular Compound- tightly bound atoms in a single unit Hydrocarbons- compounds only containing carbon and hydrogen (ex. Methane CH4) Anion: (-) ion vs. Cation: (+) ion Polyatomic ions- molecules that gain or lose electrons o Alkali (1+ charge) (lose 1 e) o Alkali Earth (2+ charge) (lose 2 e) o Halogens (1- charge) (gain 1 e) o Noble Gases (no charge) Ionic vs. Covalent Compounds o Ionic- metal+ nonmetal (ex. NaCl) o Covalent- nonmetal+ nonmetal (ex. HCl) Compounds are not charged Diatomic elements- (H2, N2, O2, F2, Cl2, Br2, I2) NOMENCLATURE- o Polyatomic Ions- Nitrate, Carbonate, Chlorate, Sulfate, Phosphate Per____ (1 more O) ___ate (standard) ___ite (1 less O) hypo____(2 less O) o Ex. NaNO3= Sodium Ntrate, BaSO3= Barium Sulfite o Monatomic Ions- have “ide” ending o Ex. FeCl2= Iron (II) chloride, Li2O= Lithium oxide o Molecular compounds- 1. Write element farthest from upper right first (prefix) nd 2. Add name of 2 element with “ide” ending (prefix) o Acids- one ion is attached to a hydrogen Hydro+ first syllable of the anion+ suffix (ic) o For transition metals USE ROMAN NUMERALS! o Main group metals- use roman numerals only if metal has multiple charges Ex. Mg2+= Magnesium, Pb2+= Lead(II) ALUMINUM IS ALWAYS 3+ and ZINC IS ALWAYS 2+ Chapter 6 Electromagnetic (EM) radiation has both wave-like and particle-like properties o Wave-like properties= wavelength and frequency o Particle-like properties= photoelectric effect Production of electrons when light is shone onto a material. NEED TO KNOW EQUATIONS o (νλ=c) frequency x wavelength = speed of light o (E=hν) planck’s constant x frequency = energy o (E=hc/λ) Relationship between frequency, wavelength, and energy o Energy and Frequency will move together (up or down) o Wavelength is inversely proportional to both Wavelength Red has the lowest frequency and energy but longest wavelength Violet has the highest frequency and energy and shortest wavelength Bohr Model o Depicts the atom as a small, positively charged nucleus surrounded by electrons that travel in circular orbits around the nucleus. o Why it is inaccurate- they are three dimensional (electron cloud) Absorption- as photons of light are absorbed they move into higher energy levels o Example- (N=1 -> N=2) Emission- move into lower energy levels o Example- (N=2 -> N=1) DOES NOT NEED TO BE MEMORIZED o Energy= Rutherfords (1/n final ^2 – 1/ n initial ^2) Ground state- lowest energy state of an atom Excited state- any quantum state that has a higher energy than ground state De Broglie relation-all moving particles have wavelengths o frequency is equal to planck’s constant divided by mass multiplied by velocity Atomic orbital- Region in space (outside the nucleus) where electrons are most likely to be found. Quantum Numbers- o Principal (n) n=1,2,3,… related to distance from nucleus o Azimuthal (l) (angular momentum) L=0,1,2,3,…. Related to shape of orbital o Magnetic (ml) Ml= l, l-1, l-2,…-(l-1), -l Related to orientation of orbital o Spin (ms) Ms= +1/2 or -1/2 Related to spin axis of electron S=2 electrons P=6 electrons D=10 electrons F=14 electrons Pauli Exclusion Principle- no two electrons in an atom can have the same four quantum numbers Building-up Principle- (Aufbau) scheme for filling electrons into subshells of atom o Tries to reproduce ground state of atom Hund’s Rule- orbitals of subshells fill singly before doubly (for lowest energy configuration) Degenerate- when two or more orbitals have the same energy Examples-
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