CHM103 Exam 1 Study Guide
CHM103 Exam 1 Study Guide 103
Popular in Chemistry for Life Sciences I (Lecture)
verified elite notetaker
Popular in Chemistry
This 10 page Study Guide was uploaded by askcch on Friday September 16, 2016. The Study Guide belongs to 103 at University of Miami taught by Elliot Atlas in Fall 2016. Since its upload, it has received 42 views. For similar materials see Chemistry for Life Sciences I (Lecture) in Chemistry at University of Miami.
Reviews for CHM103 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: 09/16/16
CHM103 Class otes eek (8/228/26) ________________________________________________________________________________ From revious ections N/A ________________________________________________________________________________ Chapter 1: atter and Measurements 1.1 Chemistry: The entral Science Chemistry: The study of nature, properties, and transformations of matter Matter: The physical material that makes up the universe; anything that has mass and ccupies space Property: A characteristic useful for identifying a substance Physical vs chemical properties Physical: observed without chemical reaction & without changing the chemical composition Color, size, shape, volume, density, temperature, boiling point/melting oint, etc. Chemical: observed when substance undergoes chemical change or reaction Toxicity, oxidation, flammability, heat of combustion, radioactivity, etc. Physical vs chemical change Physical: change in state, size, formation/separation of mixtures Chemical: hange in chemical composition 1.2 States of Matter Solid, iquid, Gas 1.3 Classification Matter Mixture vs Pure substance Mixture: A blend of two or more substances, each of which retains its chemical dentity Homogeneous: A uniform mixture that has the same composition throughout can’t be easily separated) Heterogeneous: A nonuniform mixture that has regions of different composition Pure substance: A substance that has a uniform chemical composition throughout Element: e.g. Na, H, e, a Compound: .g H2O, NaCl 1.4 Chemical Elements and Symbols 118 elements have been identified; 91 occur naturally 1.5 Chemical Reactions: Examples of Chemical Change Reactants → Product(s) E.g. Ni (s) + 2HCl (aq) → NiCl2 (aq) + H2 (g) CHM103 lass Notes Week 2 8/299/2) ________________________________________________________________________________ From revious ections 1.11.5 Definition and examples of matter, physical and chemical properties/change Chemical lements and symbols, chemical eactions ________________________________________________________________________________ Chapter 1: atter and Measurements 1.6 Physical Quantities: Units and Scientific Notation The ysteme International d’ Unites SI) Mass ilograms (kg) Length → eters (m) Volume cubic eters (m^3) Temperature → kelvins (K) Time → seconds (s) Electric Current → mpere (A) Substance Amount → mole (mol) Speed m/s Density → /cm^3 Scientific Notation 7650 .65 10^3 0.00215 = 2.15 x 10^3 1.7 easuring Mass, Length, and Volume Mass: a measure of the amount of matter in an object Weight: a measure of the gravitational force that the earth or other large body exerts on n object Length: the SI unit meter is way too big for usual lab measurements Volume: the amount of space occupied by an object 1.8 easurement and Significant igures Every set of measurements has a degree of uncertainty to it,thus significant figures re pplied Accuracy: how close to observed value is to the “true” value Precision: the reproducibility of values; does it produce the same value every time? Two ypes of rrors for measurements: Random error: randomly too high/low Systematic rror: consistently too high/low 1.034 → 4 sig. fig. 0.0738 → 3 sig. fig. 7.360 → 4 sig. fig 7600 → 2 sig. fig. 1.9 Rounding off Numbers Rule 1: During multiplication or division, the answer can’t have more sig. fig. than the original numbers (the one that has the least sig. fig.) Rule 2: During +/,the answer can’t have more digits after the decimal point than the original numbers (the one that has the least sig. fig.) 1.10 Problem Solving: Unit Conversions and Estimating Answers Factorlabel method Starting Quantity x Conversion Factor = Equivalent Quantity Conversion factors re umerically equal to e.g. 26.22mi x (1km/0.6214mi) 2.20km e.g. 1yard/s → (x)mi/h? x ≈ 2mi/h CHM103 lass Notes Week 3 9/59/9) ________________________________________________________________________________ From revious ections 1.61.10 Units and Scientific Notations, Significant Figures and Rounding, Unit Conversions and Estimating Answers ________________________________________________________________________________ Chapter 1: atter and Measurements 1.11 Temperature, eat, and Energy Energy: the capacity to do work or supply heat 1000cal = 1kcal, 1000J = 1KJ, 1cal = 4.184J, 1kcal = 4.184KJ Heat (cal) = Mass (g) x Temp Change (Δ℃) x Specific Heat (cal/(g*℃) Temperature: ℉, ℃, ℃ and K are the same size (^ 1℃ = ^ 1K), but DIFFERENT in starting point 0℃ = freezing point of water, 273.15℃/ 0K = Absolute Zero 32℉ = 0℃, 212℉ = 100℃, ^ 1℃ ^ .8℉ 1.12 Density & Specific Gravity Density: mass per unit volume → solid (g/cm ) → liquid (g/ml) Hydrometer measures pecific gravity ________________________________________________________________________________ Chapter 2: Atoms and the eriodic Table 2.1 Atomic heory nd he Structure of Atoms Atomic heory: 1. All matter is composed of atoms 2. The atoms of a given element differ from the atoms of all other elements 3. Chemical compounds consist of atoms combined in specific ratios. Only whole atoms an combine 4. Chemical reactions change only the way atoms are combined in compounds Protons: carry a positive lectrical charge Neutrons: have a mass similar to that of a proton, but are electrically neutral Electrons: very small in mass, carry a negative electrical charge All atoms are neutral and have o net charge Number of protons = Number of electrons, to balance out Atomic mass unit (amu) is the unit for describing the mass of an atom 1amu = 1.660539 x 10^24 g = 1/12 wt of 1 C12 atom The protons and neutrons are packed densely together, forming the nucleus Opposite lectrical charges attract ach other Like charges repel each other 2.2 Elements and tomic Number Atomic Number is the number of protons in atoms of a given element Mass Number is the sum of the protons and neutrons in an atom 2.3 Isotopes and tomic Weight Isotopes are toms ith dentical atomic numbers but different mass umbers A specific isotope s epresented by The isotopes of most elements do not have distinctive names The mass number A is given after the name of the element (e.g. Uranium235, Carbon12) Most naturally occurring elements are mixtures of isotopes CHM103 Class N otes eek (9/129/16) ________________________________________________________________________________ From revious ections 1.111.12, 2.12.3 Temp, eat, Energy, Atoms, Isotopes ________________________________________________________________________________ Chapter 2: Atoms and the eriodic Table 2.4 The Periodic able Metal: A malleable element, with lustrous appearance, good conductor of heat & electricity (can be worked ith/molded) Non metal: Poor conductor of heat & electricity Metalloid: Properties in between metal and nonmetal Elements in a group (vertical column) have similar properties 2.5 Some Characteristics of Different Groups Group 1A lkali etals Lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and rancium (Fr) Shiny, soft metals with ow melting points React with water to form products that re ighly alkaline Never ound in nature n ure tate due to heir igh reactivity Group 2A lkaline arth etals Beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), nd radium (Ra) Lustrous, silvery metals Less reactive han Alkali etals 1A) Never found in nature n a pure state Group 7A Halogens Fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At) Colorful and corrosive non etals Found in nature only in combination with other elements, such as with sodium in table alt (Sodium Chloride) The group name, halogen, is taken from the Greek word als meaning salt Group A Noble ases Helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn) Colorless gases Labeled the noble” gases because they on’t react He, Ne, and Ar don’t combine with any other elements. Kr and e combine with very ew 2.6 Electronic Structure of Atoms The properties of the elements are determined by the arrangement of electrons in heir atoms Electrons are not perfectly free to move. They are restricted to certain energy alues, or uantized Shell: A grouping of electrons in an atom according to energy Shell number > Electron capacity: 1>2, 2>8, 3>18, 4>32 Within the shells, electrons are further grouped into subshells (s, p, d, f) Within each subshell, electrons are grouped into orbitals 2.7 Electron Configurations The exact arrangement of electrons in an atom’s shells and subshells Rule 1: Electrons occupy the lowest energy orbitals available. This is complicated by “crossover” of energies above the 3p level Electrons fil orbitals from the lowestenergy orbitals upward Rule 2: Each orbital can hold only two electrons, which must be of opposite spin Rule 3: Two or more orbitals with the same energy are each halffilled by one electron before any one orbital is completely filled by the addition of the second electron A shorthand using noble gas configurations is very useful for large atoms 2.8 Electron Configurations and the Periodic Table A valence shell is the outermost electron shell of an atom A valence electron is an electron in the valence shell of an atom 2.9 Electrondot Symbol an atomic symbol with dots placed around i to indicate the number of valence electrons
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'