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Chemistry 1- CHEM 1331 Exam 1 Review

by: Alexis Clowtis

Chemistry 1- CHEM 1331 Exam 1 Review CHEM 1331

Marketplace > University of Houston > Chemistry > CHEM 1331 > Chemistry 1 CHEM 1331 Exam 1 Review
Alexis Clowtis
GPA 4.0

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About this Document

Review of the material covered for the first exam, condensed and geared towards the material given by the review questions.
Fundamentals of chemistry
Thomas Teets
Study Guide
Chemistry, Exam 1, review
50 ?




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This 52 page Study Guide was uploaded by Alexis Clowtis on Monday February 22, 2016. The Study Guide belongs to CHEM 1331 at University of Houston taught by Thomas Teets in Spring 2016. Since its upload, it has received 185 views. For similar materials see Fundamentals of chemistry in Chemistry at University of Houston.


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Date Created: 02/22/16
Chemistry 1331 Exam 1 Review Chapters 1-3 Review Chapter • Units/Conversions • Prefixes • Measurement • Temperature • Density • Matter Units/Conversions • Use dimensional analysis to convert • 1in=2.54cm • Any with prefixes (Ex. 1 kilometer=1000 meters) Prefixes Measurement • Certain digits: digits not estimated • Uncertain digits: digits must be estimated due to limited precision of measuring device • True Value: accepted value for a measurement • Accuracy: how close a measurement is to the true value • Precision: how well multiple measurements match with each other Temperature Density • Mass of a substance per unit of volume • Can be used to ID a pure substance • Can be used to separate mixtures Matter • Matter= anything that has mass • Pure substance= matter with constant composition • mixture:=matter with variable composition, 2+ pure substances combined physically • Elements= cannot be decomposed to simpler things/substances physically or chemically • Compounds= pure substances with 2 or more elements Chapter 1 • Scientific Method • Early Chemical History • Dalton’s Atomic theory • Subatomic particles • Picture of the atom • Isotopes • Ions Scientific Method Early Chemical History • The Greeks: • Tried to explain chemical changes • Proposed 4 fundamental substances- Earth, fire, water, air • First to propose atoms as fundamental building blocks of matter as fundamental particles • The Alchemists: • “pseudoscience”, spanned 2000 years • Tried to change base metals (ex lead) to gold • Robert Boyle: • First “chemist” • Quantitative measurement on gases • Proposed the concept of elements • Joseph Priestly: • Discovered oxygen, careful studies of chemical reactions Three Laws, One Theory 1. Mass conservation- Antoine Lavoisier • In a chemical reaction, the mass of reactants=mass of products • Matter is neither created nor destroyed 2. Definite Composition- Joseph Proust • A given compound always contains the same proportion of elements by mass, regardless of source mass %= (mass of element)/(mass of compound) X 100 (Mass fraction is part before X100) 3. Multiple Proportions • If 1 gram of Element A combines with Element X to form 2 different compounds, the masses of X that combine with A are small, whole number ratios of each other - A1+X1 α A2+X2 Dalton’s Atomic Theory • Substances are composed of tiny particles called atoms which cannot be decomposed further • Atoms cannot be converted into different atoms- elements cannot be interconverted • Atoms of a given element have the same mass and other properties; distinct from other elements • Chemical compounds are formed when atoms of different elements combine in a specific ratio Guy-Lussac’s Observations: at = temp and pressure, gases combine in chemical reactions in fixed, whole number ratios Avogadro’s Interpretation: at the same temp and pressure, equal volumes of different gases hae the same number of particles Subatomic Particles The electron(first subatomic particle discovered): Cathode Ray Experiments- JJ Thompson • Cathode rays are repelled by negative pole of an electric field • Electrons could be produced from many different materials • Concluded all atoms have these electrons • Determined charge to mass ratio (couldn’t get absolute charge or mass) Oil Drop experiment- Milikin • Falling negatively charged oil drops stopped by an electric field (induce electric field and charged particles) • Charges on oil drops occur in regular intervals • Determined the mass and charge of an electron (using this info and stuff from cathode ray) • Didn’t add info about structure, just reaffirmed existence of an electron Gold Foil Experiment (Rutherford): Most particles went straight through but some were deflected at very large angles or reflected completely (came straight back at source); The positive charge and most of the atom’s mass is in a very small area called the nucleus  Nuclear picture of atom Early Picture of Atom Plumb-pudding model (Thomson)- atoms consist of electrons held in place by diffuse cloud of positive charge; electrons are randomly embedded in this positive charge Atomic Symbols X: atomic symbol= 1 or 2 letter abbreviation for element, first letter is capitalized (know symbol-name for first 36 elements Z: atomic number= number of protons in the nucleus A: mass number= number of protons+ number of neutrons Isotopes • Atoms of a given element that differ in the number of NEUTRONS (and hence the mass is different)  Dalton was wrong! Ions Neutral atoms must contain = #s of protons and electrons to balance charge Formation of ions is achieved through removing or adding ELECTRONS (NEVER PROTONS because that would change the identity of the atom itself) –Just mess with the cloud Add one electron for each negative charge, remove one for each positive charge Chapter 2 • Quantum Theory • Properties of a wave • Quantization of energy • Wave-particle duality • Atomic line spectra • Bohr Model • Schrodinger Equation • Heisenberg Uncertainty Principle • Radial Probability Distribution • Quantum Numbers • Atomic orbitals • Pauli exclusion Principle • Electron Energy and configurations • Hund’s Rule • Atomic Radius, Ionization energy, Electron Affinity Quantum Theory • Electromagnetic wave (“radiation”): energy propagated through space by electric and magnetic fields that oscillate in intensity as they travel • Ex. Visibile light, x-rays, radio waves, microwaves Properties of a wave • Frequency- number of cycles per second(ν) • Wavelength- distance between 2 corresponding points on a wave(λ) • Amplitude- height of crest=intensity (A) Quantization Energy • Blackbody radiation- energy is quantized in whole-number multiples of hν (introduced by Max Planck) • Photoelectric effect (Einstein)- Light consists of particles called photons 2 ℎ • de Broglie wavelength-???? = ???????? so for matter λ???????? Wave-particle Duality • Matter and energy both have dual particulate and wave properties • Electrons exhibit wave-like behavior • Diffraction-scattering of light by a regular army of points or lines • Electrons are diffracted by crystals Atomic Line Spectra • Continuous spectrum (white light)- visible wavelengths • Hydrogen line spectrum- only a few sharp lines observed, each corresponds to distinct wavelength Bohr Model of the Atom • Derived for hydrogen atom • Atom’s energy does not change while electron moves in an orbit • Quantum numbers- associated with the radius of an electron’s orbit Eαn^2 • Ground state (n=1)- lowest energy orbit • Absorption- electron moves to higher energy • Emission- electron moves to lower energy • Ionization of hydrogen atom (see equations) *this property only applies to 1 electron systems* Schrodinger Equation • Ĥψ=Eψ • Can define for any bound particle • Describes electron as standing wave in terms of x,y,z coordinates Heisenberg Uncertainty Principle • The reason why the Schrodinger equation can’t tell where electron exactly is • Its impossible to simultaneously know the position and momentum of a particle ℎ ∆???? ∗ ????∆???? ≥ 4???? Quantum Numbers L=0: S orbital L=1: p orbital L=2: d orbital L=3: f orbital Atomic Orbitals • Node: region where electron has 0 probability of existing • Radial node: a certain r from the nucleus where the probability of finding an electron is 0 (# of radial nodes=n-L-1) • Nodal Plane: a 2D plane where the electron has no probability of existing Pauli Exclusion Principle • No 2 electrons in an atom can have the same set of 4 quantum numbers • Cannot have a pair of spin up or a pair of spin down, have to have 1 th spin up 1 spin down so the 4 quantum number wont be the same Electron Energy Configurations Hund’s Rule • When degenerate orbitals are filled, first put 1 electron in each orbital, all with the same spin BEFORE pairing two electrons in the same orbitals to maximize the number of unpaired electrons • Only with degenerate subshells, not S subshells: p,d,f because minimizes electron/electron repulsions Atomic Radius, Ionization energy , Electron Affinity Chapter 3 • Predicting which ions are formed • Ionic formulas • Energetics • Born-Haber Cycle • Lattice energy • Covalent bonds • Electronegativity • Bond polarity • Bond strength • Lewis Structures Predicting which ions are Formed • Anions- add electrons to the partially filled p shell of main group ions • S-block cations: remove the outermost s electrons • Cations are always smaller than the parent atom • Anions are larger than the parent atom (because electrons added and e-/e- repulsion increases Rule: A-group elements (1A, 2A, 5A, 6A, 7A) lose or gain electrons to form ions with the same number of electrons as the nearest noble gas (8A) Ionic formulas • Must be balanced charges • Subscripts must have smallest whole numbers possible Energetics 1. Ionization - Ionization energy if removing electron, electron affinity if adding - Put them together and add IE1 and EA1 to get ΔE - Formation of gas-phase ions is unfavorable 2. Ion attraction: electrostatic interaction of oppositely charged ions releases energy (negative number) -ΔE is also called ΔH -Gas phase= X+(g) + Y-(g)  XY(g) -Solid phase= X+(g) + Y-(g)  XY(s) ΔE=ΔElattice<0 Born-Haber Cycle allows us to find lattice energy Born-Haber Cycle Eformation Esublimation Edissociation Electron affinity Ionization energy Possible vaporization energy Lattice Energy • Larger ions- smaller lattice energy, Δelatt is less negative • Higher charge- higher lattice energy (quite a large effect)- more negative Covalent Bonds • Bond strength=bond dissociation=bond energy= energy required to break one mole of a bond in the gas phase • Bond length= distance between nuclei of 2 bonded atoms (inversely related to bond energy) • Single bond: 1 shared electron pair, Bond order=1 • Double bond: 2 shared electron pairs, bond order= 2 • Triple bond: 3 shared electron pairs, bond order=3 • HIGHER BOND ORDER=SHORTER BOND=STRONGER BOND Electronegativity • Same trend as electron affinity (closely related) • Noble gasses not included in trend because don’t have a defined electronegativity because they don’t form bonds • LR higher increase than top to bottom Bond Polarity • When atoms of 2 different electronegativity bond (basically just 2 different atoms bonding) the bonding pairs of electrons are NOT shared equally  indicate polar bonds • Partial negative charge on MORE electronegative atom Lewis Structures 1. Find total number of valence electrons 2. Draw all single bonds 3. Distribute remaining electrons as lone pairs 4. If the central atom does not have an octet, form 1+ multiple bonds 5. Determine formal charges Resonance Structures • When the central atom has multiple bonds (double, triple) it is possible to put the double bonds in different positions • Better to have formal charge on more electronegative (outer) atom Rules for Ionic Compounds Metal first (cation) Nonmetal second (anion) Take the name of the metal for the cation For the anion, take the root and then add –ide ending Subscripts= lowest whole number ratio *NO prefixes for IONIC compounds* For TRANSITION metals, exact same convention but roman numeral to denote the cation’s charge For common name, higher charge –ic suffix, lower charge –ous suffix (instead of roman numeral form) Polyatomic Ions Covalent Compounds • Binary Covalent Compounds- covalent compound with ONLY 2 different atoms • First atom with lower group number • If in the same group, higher period comes first • Ide ending on second element • Prefixes tell how many of each atom • formula is NOT reduced to lowest whole number ratio of subscripts Acids Acid= anion + H+ Binary acid: H+ with a monoatomic ion (hydro) + (anion root) + (ic) acid Oxoacid: H+ paired with polyatomic “oxoanions” (root name) + (ic/ous) acid -ate anion= -ic acid -ite anion=-ous acid Equations to Know Ephoton=hν=hc/λ C=λν −18 2 1 1 ∆???? = −2.18????10 ∗ ???? (????2− ????2) (ionization of a one-electron atom) ???? ???? Number of orbitals= n if only given n, 2L+1 if you’re given L Number of electrons in subshell= double the number of orbitals (because 2 electrons in each subshell) Number of nodal planes= L Number of radial nodes= n-L-1 Heisenberg Uncertainty Principle: ∆???? ∗ ∆(????????) ≥ ℎ 4???? X: uncertainty in position- how precise could we be m: mass v: uncertainty in speed h: planck’s constant Formal charge= #valence electrons-#bonds atom forms-#nonbonding electrons


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