Chem 101 Exam #1 Study Guide
Chem 101 Exam #1 Study Guide Chem 1010
Popular in Chemistry 101
verified elite notetaker
verified elite notetaker
ESS 210 001
verified elite notetaker
Char Paul psychnstatstutor
verified elite notetaker
verified elite notetaker
verified elite notetaker
Popular in Chemistry
This 10 page Study Guide was uploaded by Josh Trac on Tuesday January 5, 2016. The Study Guide belongs to Chem 1010 at Clemson University taught by William Wallace in Winter 2016. Since its upload, it has received 80 views. For similar materials see Chemistry 101 in Chemistry at Clemson University.
Reviews for Chem 101 Exam #1 Study Guide
Report this Material
What is Karma?
Karma is the currency of StudySoup.
You can buy or earn more Karma at anytime and redeem it for class notes, study guides, flashcards, and more!
Date Created: 01/05/16
Fall 2015_CH1010_Dr. Kreider-‐Mueller CH1010 Exam #1 Study Guide For reference see “Chemistry: An Atoms -‐focused Approach” by Gilbert, Kirss, and Foster Chapter 1: Matter and Energy, An Atomic Perspective • Know the following SI base units of measure and their abbreviations (Table 1.2 in text): *Mass (kg) *Temperature (K) *Time (s) *Length (m) *Amount of substance (mol) • You need to know (memorize) the following prefixes for multiples of SI Units o Mega, kilo, deci, centi, milli, micro, and nano (Table 1.1 in your textbook). o Be able to do the following: § Know the abbreviations for each of the prefixes 3 § Know the conversion factors (ex. 1 g = 10 mg) § Be able to perform these types of conversions in word problems • Know the equations for converting between Kelvin & Celsius (we will not provide you with these equations on the exam!) K = ˚C + 273.15˚ ˚C = K – 273.15˚ • Know what the difference is between SI base units and derived units and how to perform conversions with derived units (ex: convert 3 m to 3 cm ) 3 o Know the following conversions for volume: 1 dm = 1 L 1 cm = 1 mL o Know the equation for density (g/mL) (we will not provide you with this equation!) and how to use it in a word problem. Practice solving for density given a mass and volume. Practice solving for mass given a density and volume. Practice solving for volume given a density and mass. Recognize how changes in mass and/or volume affect density. o Unit for energy is Joules (J). 1 J = (kg•m )/s • Know how to work with scientific notation. Given a number, be able to represent it in scientific notation. Given a number in scientific notation, be able to represent the number in its standard form. (See Appendix 1 in your textbook) • KNOW THE RULES OF SIGNIFICANT FIGURES: 1) All nonzero digits are significant 2) Zeros in the middle of a # are significant § Ex: 2.9905 has 5 significant figures 3) Zeros at the beginning of a # are not significant § Ex: 0.025 has 2 significant figures 4) Zeros at the end of a # and after a decimal point are significant § Ex: 245.0 has 4 significant figures 5) Zeros at the end of a # and before a decimal point may or may not be significant § Ex: 34,200 may have 3, 4, or 5 significant figures § In this case it is best to use scientific notation to indicate the proper number of significant figures 1 of 10 Fall 2015_CH1010_Dr. Kreider-‐Mueller • Know the rules for significant figures when performing addition & subtraction, as well as multiplication & division. • What is the difference between accuracy and precision in measurements? • Know how to round numbers: o If the 1st digit you remove is less than 5 round down by dropping it and all following digits. o It the 1st digit you remove is 5 or greater, round up by adding 1 to digit on the left. • PRACTICE CONVERTING BETWEEN ONE UNIT AND ANOTHER. We have done countless examples in class, on quizzes, in the lecture slides, on HMWK assignments etc. You need to be comfortable with conversion problems. You may need to perform several conversions in any one problem. • Be able to distinguish between solids, liquids, and gases at all 3 levels of representation: symbolic (H O),2 particulate (one molecule of H O)2 and macroscopic (1 g of 2 O) o Be able to recognize graphical representations of molecules o The properties of matter are related to its molecular level structure • Understand that chemistry is an experimental science and depends on accurate, precise, and reproducible measurements o Scientific Method: an approach to acquiring knowledge based on observation of phenomena, development of a testable hypothesis, and additional experiments that test the validity of the hypothesis. o Hypothesis: a tentative and testable explanation for an observation or a series of observations o Scientific Theory: A general explanation of widely observed phenomena that has been extensively tested. o Scientific Law: a concise and generally applicable statement of a fundamental scientific principle • Matter can be divided into two principle classes (Figure 1.5 in textbook): o Pure substances: Matter that cannot be separated into simpler matter by a physical process § Element: a pure substance that cannot be separated into simpler substances by an chemical process • Atom: the smallest particle of an element that retains the chemical characteristics of the element § Compound: a pure substance that is composed of 2 or more elements linked together in fixed proportions and that can be broken down into those elements by some chemical process. (A sample of a given compound may contain 1 molecule or several molecules of that compound.) • Molecule: a collection of atoms chemically bonded together o Mixtures: a combination of pure substances in variable proportions in which the individual substances retain their chemical identities and can be separated from one another by a physical process 2 of 10 Fall 2015_CH1010_Dr. Kreider-‐Mueller § Homogeneous Mixture: a mixture in which the components are distributed uniformly throughout and have no visible boundaries or regions § Heterogeneous Mixture: a mixture in which the components are not distributed uniformly, so that the mixture contains distinct regions of different compositions Introduction to Energy • Matter exists in 3 phases or physical states: solid, liquid gas • Matter can be transformed from one physical state to another as its temperature is raised or lowered: o solid è liquid (melting; endothermic) o liquid è gas (vaporization; endothermic) o solid è gas (sublimation; endothermic) o gas è liquid (condensation; exothermic) o liquid è solid (freezing; exothermic) o gas è solid (deposition; exothermic) • Understand the difference between an exothermic and endothermic process o Exothermic process: energy flows from a system into its surroundings o Endothermic process: energy flows from the surroundings into the system • Be able to identify the system and the surroundings for phase changes. Where does the energy (heat) flow? From the system to the surroundings? Or from the surroundings to the system? • Understand the difference between heat and temperature: o Heat: a flow of energy from one object or place to another due to differences in the temperatures of the objects or places o Thermal energy: the portion of the total internal energy of a system that is proportional to its absolute temperature o Temperature: a measure of thermal energy • What is kinetic energy? KE = ½mv (where m = mass, v = velocity) • What is potential energy? • Why is thermal energy classified as kinetic energy rather than potential energy? Chapter 2: Atom, Ions, and Molecules • Know the following Laws: o Law of Mass Conservation: Mass is neither created nor destroyed in chemical reactions o Law of Definite Proportions: Different samples of a pure chemical compound always contain the same proportion of elements by mass o Law of Multiple Proportions: Elements can combine in different ways to form different chemical compounds, whose mass ratios are simple whole-‐number multiples of each other. 3 of 10 Fall 2015_CH1010_Dr. Kreider-‐Mueller • Know the various points in Dalton’s Atomic Theory: 1) Matter is composed of small particles called atoms. 2) Atoms of the same element are identical in shape and mass, but differ from the atoms of other elements. 3) Atoms of an element cannot be changed into atoms of a different element by chemical reactions. Atoms cannot be created or destroyed in chemical reactions. 4) Atoms of different elements may combine with other atoms in fixed, simple, whole number ratios to form compounds. A given compound always has the same relative number and kind of atoms. ∗ Note: we now know that atoms can be further subdivided into protons, neutrons, and electrons (but not by chemical processes). We also know that atoms of some elements vary in their masses and densities (isotopes). Dalton’s 2 postulate stated above is not correct! • What were the main conclusions from these experiments (What did we learn from them? Section 2.1 in textbook) o Thomson’s cathode ray experiment o Millikan’s Oil drop experiment o Rutherford’s Gold Foil experiment • Understand our current model of the atom: o Nucleus: the positively charged center of an atom that contains nearly all the atom’s mass o Proton: a subatomic particle, present in the nucleus of an atom, that has a relative charge of 1+ and a mass number of 1 o Neutron: An electrically neutral (uncharged) subatomic particle with a mass number of 1 o Electron: a subatomic particle that has a relative charge of 1− and essentially zero mass • Be able to relate the symbols and names for the elements in the first four rows of the periodic table (elements hydrogen through krypton) • Dimitri Mendeleev published a table that is considered the forerunner of the modern periodic table of elements. • The modern periodic table arranges the elements in order of their atomic numbers • What is a period of elements? All of the elements in a row of the periodic table. • What is a group or family of elements? All of the elements in a column of the periodic table. Elements in a group have similar properties! • Know how to use the periodic table to look up the atomic number (Z) and Mass Number (A) of different elements. • Know how to read/write a chemical symbol for a given element. Where to we put the mass number? Where do we put the atomic number? o Atomic number (Z): the number of protons in the nucleus of an atom 4 of 10 Fall 2015_CH1010_Dr. Kreider-‐Mueller o Mass Number (A): the number of nucleons in an atom § Nucleon: a proton or neutron in a nucleus • Know how to calculate the mass number (A) A = Number of protons (Z) + number of neutrons • Understand the difference between Mass Number and Average Atomic Mass. o Average Atomic Mass: the weighted average of masses of all isotopes of an element, calculated by multiplying the natural abundance of each isotope by its mass in atomic mass units and then summing the products. The average atomic mass is not a whole number! o Isotope: atoms of an element containing the same number of protons but different numbers of neutrons. o Be able to identify the information given by the symbol for the isotope of an element o Be able to calculate atomic masses from relative abundances of isotopes and vice versa. • Know where the following groups are located in the periodic table, and be familiar with a few of their properties (recognize that there are multiple ways to label groups in the periodic table): o Transition Metals o Inner Transition Metals § Lanthanides § Actinides o Main Group Elements § Main Group nonmetals • Halogens (Group 7A or Group 17) • Noble gases (Group 8A or Group 18) § Main Group Metals • Alkali Metals (Group 1A or Group 1) • Alkaline-‐Earth Metals (Group 2A or Group 2) • How do metals, semimetals, and nonmetals differ? Where are they located in the periodic table? What are the basic properties of each? Which ones are good conductors of electricity? • Be able to recognize the difference between ionic solids and discrete molecules o Ion: An atom or molecule that has a positive or negative charge § Cation: A positively charged ion § Anion: A negatively charge ion o Ionic solid: a solid consisting of monatomic or polyatomic ions held together by ionic bonds o Be able to assign charges to the ion derived from MAIN GROUP elements (Figure 2.10 in textbook) 5 of 10 Fall 2015_CH1010_Dr. Kreider-‐Mueller • Interpret chemical formulas & calculate the formula mass of a compound. • Understand that in our macroscopic world, quantities of substances contain extremely large numbers of atoms, ions, or molecules! We use the mole (the chemist’s dozen) to express quantities of substances. The mole relates macroscopic quantities of substances, such as their masses expressed in grams, to the # of particles they contain. o Mole: An amount of a substance that contains Avogadro’s number of particles (atoms, ions, molecules, or formula units) o Avogadro’s number (N ): 6.022A × 10 . The number of carbon atoms in exactly 12 grams of the carbon-‐12 isotope o Molar Mass: the mass of 1 mole of a substance. o Be able to convert between the mass of a substance and the amount (# of moles) and vice versa. Chapter 3: Atomic Structure, Explaining the Properties of Elements • What is radiant energy? What is the electromagnetic spectrum? o Electromagnetic Radiation: any form of radiant energy in the electromagnetic spectrum • Know the order of the different regions in the electromagnetic spectrum (Figure 3.1 in textbook). In order of increasing wavelength: Gamma rays < X-‐rays < UV < Visible < IR < Microwaves < Radio waves 6 of 10 Fall 2015_CH1010_Dr. Kreider-‐Mueller • Know the different parts of a wave: amplitude, peak, trough, wavelength, frequency • How are frequency and wavelength related? c = νλ where c = speed of light (2.998 × 10 m/s) ν = frequency (Hz or s ) λ = wavelength (m) • Understand the development of the current model of the electronic structure of atoms o Understand what the Balmer-‐Rydberg equation tells us (YOU DO NOT NEED TO USE/MEMORIZE THIS EQUATION FOR THE EXAM!) o What is a continuous spectrum? What does it look like? o What is an emission (line) spectrum? What does it look like? § Atomic emission spectra: characteristic patterns of bright lines produced when atoms are vaporized in high-‐temperature flames or electrical discharges. o What is an absorption spectrum? What does it look like? § Atomic Absorption spectra: characteristic patterns of dark lines produced when an external source of radiation passes through free, gaseous atoms. o Use the concepts of energy levels and orbitals to explain the occurrence of emission and absorption spectra o Understand the development of the dual wave/particle nature of the electrons. § The photoelectric effect: the release of electrons from a material as a result of electromagnetic radiation striking it. • Understand the experimental setup • Be able to solve word problems involving the photoelectric effect 7 of 10 Fall 2015_CH1010_Dr. Kreider-‐Mueller o Work function (Φ): the amount of energy needed to dislodge an electron from the surface of a material o Threshold Frequency (ν ): the mi0imum frequency of light required to produce the photoelectric effect • Does the intensity of a beam of light affect the number of electrons ejected from a metal? If so, how? § What does it mean when we say energy is quantized? • Quantized: having values restricted to whole-‐number multiples of a specific base value • Quantum: the smallest discrete quantity of a particular form of energy • Photon: a quantum of electromagnetic radiation § Know De Broglie’s Hypothesis and how to use it in a word problem λ = h/(mv) −34 where h = Planck constant (6.626 × 10 J•s) m = mass v = velocity § Know what Heisenberg’s Principle tells us (YOU DO NOT NEED TO USE/MEMORIZE THIS EQUATION FOR THE EXAM!) • The principle that one cannot simultaneously know the exact position and the exact momentum of an electron • Know how to use Quantum Numbers: o Quantum Number: one of four related numbers that specify the energy, shape, and orientation of orbitals in an atom and the spin orientation of electrons in the orbitals. o What is an orbital? § Regions in an atom where the probability of finding an electron is high 2 § Defined by the square of the wave function (ψ ) o How do you determine possible values of n, l, m, l m ?s (section 3.9 of textbook) o What do each of the different quantum numbers tell us about an orbital? o What letters correspond to each value of l? o Know what a shell is, and which quantum number it refers to § Shell: orbitals with the same value of n are in the same shell o Know what a subshell is, and which quantum number it refers to § Subshell: orbitals with the same values of n and l are in the same subshell • Be able to identify an orbital (s, p, d) based on it’s shape and orientation on a 3D set of axes • What is a node? How many nodes are in a p orbital? How many are in a d orbital? o Node: a location in a standing wave that experiences no displacement 8 of 10 Fall 2015_CH1010_Dr. Kreider-‐Mueller • Electron configurations: o Know how to write and identify electron configurations using the standard (expanded) form and shorthand versions. o Be able to use/generate an orbital-‐filling diagram o Know the following Principles & Rules: § Pauli Exclusion Principle: No 2 electrons in an atom can have the same values of their 4 quantum numbers § Hund’s Rule: If 2 or more orbitals with the same energy are available, one electron goes in each until all are half full. The electrons in the half-‐ filled orbitals all have the same value of their spin quantum number. § Aufbau Principle • Lower energy orbitals fill before higher energy orbitals. • An orbital can hold only 2 electrons, which must have opposite spins (Pauli Exclusion principle). • If 2 or more degenerate orbitals are available, 1 electron goes into each until all are half-‐full (Hund’s Rule) o Know that with increasing energy, the gap between one shell and the next shell decreases. (Ex. there is a large energy gap between 1s and 2s orbitals, but a smaller energy gap between 2s and 3s orbitals) o Be able to explain anomalous electron configurations for Cr and Cu § Cr: [Ar]3d 4s § Cu: [Ar]3d 4s 1 o Recognize the relationship between an element’s electron configuration, the number of valence electrons, and its Group Number. Elemental families have similar electron configurations § Core Electrons: Electrons in the filled, inner shells in an atom or ion that are not involved in chemical reactions § Valence Electrons: electrons in the outermost occupied shell of an atom having the most influence on the atom’s chemical behavior § Group 1A atom: [Noble Gas] ns 1 2 § Group 2A atom: [Noble Gas] ns 2 4 § Group 6A atom: [Noble Gas] ns np § Group 7A atom: [Noble Gas] ns np 2 5 § How do electron configurations help us explain why metals tend to form cations and nonmetals tend to form anions? o Be able to write/identify electron configurations for ions § s block elements: form monatomic cations by losing all the outer-‐shell electrons leaving their ions with the electron configuration of the noble gas immediately preceding them in the periodic table. • Ex: Na = [Ne]3s Na = [Ne] § p block elements: form monatomic anions by gaining enough electrons to completely fill its valence-‐shell p orbitals. The ion formed has the 9 of 10 Fall 2015_CH1010_Dr. Kreider-‐Mueller electron configuration of the noble gas at the end of its row in the periodic table 2 5 − 2 6 • Ex: Cl = [Ne]3s 3p Cl = [Ne]3s 3p or [Ar] § d block elements: the electrons in orbitals with the highest n value ionize first . When forming cations, transition metals lose outer-‐shell s electrons first, followed by d electrons • Ex: Ni = [Ar]3d 4s Ni = [Ar]3d • Ex: Fe: [Ar]3d 4s 6 2 Fe : [Ar]3d 6 Fe : [Ar]3d 5 • Trends to know (and be able to explain the trend…think about Z ): eff o Effective Nuclear Charge (Z ): the effattraction toward the nucleus experienced by an electron in an atom; the positive charge on the nucleus reduced by the extent to which other electron in the atom shield the electron from the nucleus o Trend for atomic radius of neutral atoms (Figure 3.34 in textbook) o Ionic radius (Figure 3.35 in textbook) § When you form a cation, the radius shrinks (ex: the ionic radius of Na is smaller than the atomic radius of Na). Why? − § Whey you form an anion, the radius expands (ex: the ionic radius of Cl is larger than the atomic radius of Cl). Why? o Ionization Energy + − § M + Energy → M + e M + Energy → M + e M + Energy → M + e …and so forth § Look at Figure Table 3.2 in your book (Think about when/why there are large jumps between successive ionization energies) § Be able to make predictions about ionization energies for various elements o Electron Affinity § Which group in the periodic table has the highest (most negative) electron affinity? Which group(s) in the periodic table have the lowest (values close to zero) electron affinity § Look at Figure 3.37 in your book § Be able to make predictions about electron affinities for various elements § Group 1A elements have low electron affinities so they tend to lose their ns electron § Group 2A elements have low electron affinities so they tend to lose their 2 ns electrons § Group 7A elements have large electron affinities so they tend to gain one electron, adopting the electron configuration of the neighboring noble gas § Group 8A elements are inert, and don’t generally gain or lose electrons 10 of 10
Are you sure you want to buy this material for
You're already Subscribed!
Looks like you've already subscribed to StudySoup, you won't need to purchase another subscription to get this material. To access this material simply click 'View Full Document'