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Chem 111 Lecture Notes Part 1

by: Caroline Hurlbut

Chem 111 Lecture Notes Part 1 Chem 111

Marketplace > Colorado State University > Chemistry > Chem 111 > Chem 111 Lecture Notes Part 1
Caroline Hurlbut
GPA 3.7

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These are the notes for the first half of Chem 111, covering about 5 weeks of lectures.
General Chemistry I
Dr. Kerry MacFarland
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This 21 page Bundle was uploaded by Caroline Hurlbut on Friday January 22, 2016. The Bundle belongs to Chem 111 at Colorado State University taught by Dr. Kerry MacFarland in Fall 2015. Since its upload, it has received 40 views. For similar materials see General Chemistry I in Chemistry at Colorado State University.


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Date Created: 01/22/16
Basics of Matter States of Matter • matter - anything with mass that occupies space • atoms - smallest building blocks of all matter • molecules - 2 or more atoms joined in specific geometric arrangements • 3 physical states of matter — solid (s): definite volume and shape — liquid (l): definite volume but not shape — gas (g): no definite volume or shape, fills container, highly compressible • particulate nature of matter responsible for 3 phases of matter —in a solid, particles are close together and organized —in a liquid, particles are close together and disorganized —in a gas, particles are far apart and free moving Phase Transitions melting - solid to liquid • • freezing - liquid to solid • vaporization - liquid to gas • condensation - gas to liquid • sublimation - solid to gas (absorbing energy) • deposition - gas to solid (releasing energy) Classes of Matter • pure substance - constant composition and 1 type of atom or molecule — element: 1 type of atom, simplest form of substance — compound: 2 or more elements in fixed proportions • mixture - 2 or more substances in variable proportions — homogeneous: uniform throughout — heterogeneous: can be separated into different substances Properties of Matter • physical - can be observed without changing composition of substance • chemical - observed by changing composition —burning —rusting —color change —release of gas Calculations in Chemistry SI Units and Prefixes • prefixes to memorize —1 nm=10^-9 m —1 um=10^-6 m —1 mm=10^-3 m —1 cm=10^-2 m —1 km=10^3 m —1 mL=1 cm^3 Uncertainty in Measurements • uncertainty is measured to the nearest marked value and estimated one more decimal place • measurements can be any combination of precise and accurate • precision - how close repeated measurements are to each other • accuracy - how close experimental measurement is to the true value Significant Figures • answers should have no more sig figs than the value in the problem with the least number of sig figs (weakest link) • sig figs always carried out at the end of a problem, not in the middle • multiplying/dividing —answer has same number of sig figs as measurement with fewest sig figs • adding/subtracting —answer has same number of decimal places as measurement with fewest decimal places • rules —always start counting sig figs at the first numeric (non zero) value —zeroes, when used as placeholders for decimals, are never significant —zeroes in between 2 numeric values are always significant —the farthest zero to the right of a decimal place is always significant —zeroes at the end of a whole number value are never significant • exact numbers do not limit sig figs —counted —exact conversions Problem Solving and Unit Conversions • Ex. You need 325 cm of wire that sells for $0.15/ft. How much does wire cost? —make sure you have like units when you do the math —convert 325 cm to in: 325 cm/2.54 cm= 127.952756 in —convert $0.15/ft to $/in: ($0.15/ft)x(1 ft/12 in)= $0.0125/in —multiply answer by 127.952756 to cancel out in: 127.952756 in x $0.0125/in=$1.6 (smallest value in problem is $0.15/ft which has 2 sig figs, so the final answer should have 2 sig figs) Atoms, Isotopes, and Elements Temperature Conversions • T(deg. C)=5/9[T(deg. F) - 32] • T(K)=T(deg. C) + 273.15 Structure of an Atom and Subatomic Particles • nucleus - tiny, dense center of the atom, makes up most of mass, contains protons and neutrons • electrons contained outside nucleus in orbitals • proton —p+, relative charge of +1 —mass=1.0 amu • neutron —n0, no relative charge —mass=1.0 amu • electron —e-, relative charge of -1 —mass=0.0005 amu • atomic number (Z) - number of protons —all atoms of an element have same atomic number • mass number (A) - number of protons + number of neutrons • atoms of one element can have different numbers of neutrons — isotope: atom of same element with different number of neutrons • we can represent a nucleotide (specific isotope of an element) like this: —X represents element symbol —Arepresents mass number, placed in top left next to X —Z represents atomic number, placed in bottom left next to X • cation - atom with more protons than electrons, resulting in a positive charge (X+) • anion - atom with more electrons than protons, resulting in a negative charge (X-) • monatomic ions - ions containing one atom Periodic Table • organized in order by atomic number • across=rows or periods • down=columns or groups or families • similarities in members of a group —column 1 (alkali metals) and column 2 (alkaline earth metals)=most reactive —column 7 (halogens) and column 8 (noble gases)=least reactive • main group elements are groups 1A-8A transition elements/metals are in the middle 10 columns • • inner transition elements are at the bottom, elements 57-70 and 89-102 • there are nonmetals and metalloids on the far right of the periodic table Masses and Moles Weighted Average ex. pizza party • —7 people (70%)…….1 slice —2 people (20%)…….2 slices —1 person (10%)…….3 slices • calculate weighted average: (0.7x1)+(0.2x2)+(0.1x3)=1.4 slices Average Atomic Mass • ex. Si 14 protons, 28.09 amu=avg atomic mass • isotopes of Si —Si 28: 27.97693 amu —Si 29: 28.976495 amu —Si 30: 29.973770 amu • natural abundance —92.23% Si 28 —4.67% Si 29 —3.10% Si 30 • calculate avg atomic mass as weighted average: 0.9223(27.97693)+0.0467(28.976495)+0.0310(29.973770)=28.09 amu Molecular and Formula Mass • molecular mass - mass of a molecule • formula mass - mass of a formula unit of an ionic compound • formula unit tells you the simplest ratio of elements in a compound (ex. 1 Na +1 Cl=1 formula unit of NaCl) • sum of atomic masses of atoms in a molecule/formula unit (chemical formula) • ex. H2O —2 H atoms=1.0078x2 —1 O atom=15.9994 —1.0078(2)+15.9994=18.015 amu Avogadro’s Number • Avogadro’s number=6.02214x10^23=1 mole of something • used as a conversion factor, relates particle units (like atoms or amu) to measurable units (like g) • ex. How many moles of Cu in a penny contains 2.4x10^22 Cu atoms? — 2.4x10^22 Cu atoms x 1 mol Cu = 4.0x10^-2 mol Cu 6.022x10^23 Cu atoms Molar Mass molar mass - mass of 1 mole of something • • use avg atomic masses from periodic table, but in g instead of amu (1 mol of amu=1 g) • ex. Molar mass of He is 4.003 g/mol, 6.022x10^23 mol of He=4.003 g • use molar mass to convert g to mol of something Types of Energy Potential and Kinetic Energy • electrostatic interactions (attractions and repulsions) and gravity act on particles like protons and electrons • gravitational force between 2 objects —F=G(mass 1)(mass 2)/d^2 • electrostatic force between 2 charges —F=k(charge 1)(charge 2)/d^2 • potential energy is higher when objects are farther apart • we choose PE=0 for different situations —electron of a hydrogen atom is taken away completely from the atom so that they are so far apart they don’t influence each other • kinetic energy is the energy of movement • kinetic energy has inverse relationship with potential energy Nature of Light andAtomic Spectra Nature of Light • light has wave properties — wavelength (lambda): distance between crests (usually in nm) — frequency (v): number of wave crests per second that pass a stationary point (usually in Hz) —wavelength = speed of light (c=3.00 x 10^8 m/s) / frequency (v) • electromagnetic spectrum - spectrum of light organized by wavelength (lowest to highest) —radio (750 nm), microwave, infrared, visible, ultraviolet, x-rays, gamma rays (400 nm) —frequency and wavelength are inversely proportional —certain colors on the atomic spectra are emitted when light reaches certain temperatures (higher energy=further color of light on the spectrum) • light has particle properties — Quantum Theory: energy is quantized (occurs in fixed, whole number quantities) and is not continuous —smallest quantity of energy is a quantum — a quantum of electromagnetic energy (light) is a photon (can be absorbed or emitted) —energy of a photon (E) = Planck’s constant (h=6.626x10^-34 J/s) x frequency (v) —energy of a photon is directly proportional to frequency Electron Transitions in Atoms • Neils Bohr determined there are certain amounts of energy an electron can have and modeled this by placing the electrons in orbitals around the nucleus (Bohr model) —levels of energy can be calculated by E = -2.178 x 10^-18 J (1/n^2) where n is an integer from 1 to infinity —levels can be represented by placing calculated values in an energy level diagram (lowest to highest energy) • the ground state of an electron is the lowest energy level • when an electron absorbs energy, it moves to an excited state • to move down a level, an electron emits energy Orbitals and Electron Configurations Part 1 Orbitals and Subshells • quantum mechanical model - energy of electron is quantized, electron is matter and has particle and wave properties • orbitals - regions in an atom where an electron is likely to be found, orbitals identified by principal quantum number ’n’ (specifies principal shell of orbital) and subshell letter • 4 subshells, each with distinct number of orbitals and 3D shape — s subshell (n=1+): 1 orbital (2 electrons max), spherical shape — p subshell (n=2+): 3 orbitals (6 electrons max), dumbbell shape — d subshell (n=3+): 5 orbitals (10 electrons max), double dumbbell shape — f subshell (n=4+): 7 orbitals (14 electrons max), complex dumbbell shape Electron Configurations of Atoms • Aufbau principle - to build up; electrons are placed in lowest energy orbitals available • ex. H has 1 electron and is in the 1s shell, so the electron configuration is written 1s^1 (superscript on orbital=number of electrons) • Pauli exclusion principle - an orbital can hold up to 2 electrons with opposing spins • spin - fundamental property of electrons, can be up or down • an orbital diagram gives a visual representation of electron configurations —arrows used to represent opposite spin electrons: ↿and⇂ —boxes organized by energy level —max two electrons per orbital box: ⥮ • ex. He has 2 electrons and is in the 1s shell, so the configuration is written 1s^2 • ex. Li has 3 electrons and is in the 2s shell, so the configuration is written 1s^2 2s^1 • condensed electron configurations are written in noble gas notation as a shortcut —ex. configuration for Li: [He]2s^1 • valence electrons are in the outermost occupied (valence) shell, filled valence shell makes an element stable • core electrons are in filled inner shells Orbitals and Electron Configurations Part 2 Electron Configurations of Ions • subshells are organized on the periodic table —1Aand 2Aare s-block —3A-8A(excluding He) are p-block —3-12 are d-block —lanthanides and actinides are f-block • start with electron configuration for atom and add/remove electrons accordingly —ex. Na^+ —> Na - 1 electron —> [Ne] • alkali metals lose 1 electron to gain electron configuration of a noble gas • it is easier for halogens and other nonmetals to gain electrons than lose them, so they will still have a noble gas configuration —ex. Cl^- —> [Ar] or O^2- —> [Ne] • metals tend to form cations and nonmetals tend to form anions • Mg^2+, Na^+, F^-, and Ne are isoelectronic, meaning they have the same number of electrons Sizes of Atoms and Ions • sizes of atoms and ions are measured in terms of the atomic radius (distance from center of nucleus to boundary of surrounding electron cloud) or the metallic radius (atomic radius of metals) • trends in atomic size —atomic radius tends to decrease going across rows and increase going down columns —as n increases, distance between electron and nucleus increases —as atomic number (Z) increases, there are more protons to attract electrons —electrons in filled inner shells will shield valence electrons from full charge of nucleus • as you move down a group, electrons are added to different shells, increasing size • as you move across a row, electrons are added to same shell, decreasing size • trends in ionic size — more electrons (- charge) makes an ion bigger and less electrons (+ charge) makes an ion smaller More Periodic Trends Ionization Energy • photoelectric effect - many metals emit electrons when light of a high enough frequency is shined on them —if energy imparted is higher than frequency threshold, extra energy turns into kinetic energy of electron being emitted • conservation of energy - extra energy (higher frequency photons)—>kinetic energy, metal becomes cation • ionization energy - amount of energy required to remove 1 mole of electrons from 1 mole of ground state atoms/ions • first ionization energy (IE1) - amount of energy needed to remove 1 mole of electrons from 1 mole of atoms — ex. Na—>Na+ + e- —IE1=change in E, expressed in kJ/mol • trends in first ionization energy — effective nuclear charge (Z eff.) increases going across a period, so IE1 increases —Z eff. decreases going down a group, so IE1 decreases • successive ionization energy - amount of energy needed to successively remove 1 more electron from an atom (IE2, IE3, IE4, IE5, etc.) • electron affinity - energy change when 1 mole of atoms gains 1 mole of electrons —measure of how attracted an atom is to an electron —ex. Cl has high electron affinity because it only needs 1 more for full valence shell —electron affinity is highest for halogens and lowest for alkali metals Bonding and Naming Part 1 Ionic Bonds • ionic bonds form when electrons are transferred between a metal and a nonmetal —metals tend to form cations and nonmetals tend to form anions • electrostatic potential energy —Coulomb’s law: PE proportional to (q1 x q2)/d —q=charges of ions, d=distance between them —more attraction (higher PE) for larger charge magnitudes and shorter distances —PE increases when unlike charges get closer and decreases when like charges get closer —PE decreases with distance up to a certain point of inflection (see diagram) • ions form 3D array called crystal lattice • lattice energy - energy released when 1 mole of ionic compound forms from its free ions (in gas state) —depends on size and charge of ions —greater attraction as charges increase and sizes decrease —related to Coulomb’s law Naming Ionic Compounds • binary ionic compounds - ionic compounds with two elements • compound must be neutral (+ charges balance - charges) • when writing ionic compounds, the charges of the ions cross, so each element in the formula gets the other’s charge —ex. Mg^2+ and F^-1 —> MgF2 • naming binary ionic compounds containing a metal that forms only one cation (main group metals) —(name of cation) + (base name of anion + -ide) —ex. NaCl—>sodium chloride, MgS—>magnesium sulfide • naming binary ionic compounds containing a metal that can form more than one cation (transition metals exceptAg^+ and Zn^2+) —same rule as single cation compounds, but specify charge of metal with roman numeral —ex. Fe2O3—>iron (III) oxide, CuBr—>copper (I) bromide, CuS—>copper (II) sulfide Bonding and Naming Part 2 Molecular Bonds • covalent bonds form when orbitals overlap and electrons are shared between nonmetals Naming Molecular Compounds • (prefix + name of element) + (prefix + base name of element + ide) • prefixes —mono (1) —di (2) —tri (3) —tetra (4) —penta (5) —hexa (6) —hepta (7) —octa (8) —nona (9) —deca (10) • if prefix of first element is mono, skip it • when naming covalent compounds, list elements in order from left to right (and if necessary down to up) on periodic table 1st element 2nd element • ex. CO—>carbon monoxide, PI3—>phosphorus triiodide, SO2—>sulfur dioxide Lewis Structures Intro • Lewis dot diagrams are used to represent valence electrons of an element • number of dots assigned to an element depends on number of valence electrons • placement of dots around element symbol should be as evenly distributed as possible, with one dot on each side before pairing dots • octet rule - atoms will gain, lose, or share electrons to reach 8 valence electrons • exceptions to octet rule —duet rule for hydrogen (only needs 2 electrons to fill 1s shell) —boron often has 3 bonds and less than an octet —expanded octet Lewis Structures for Ionic Compounds • rules —cation written first, anion written second —brackets around nonmetal —dots placed around nonmetal element in brackets • Na has 1 valence electron, Cl has 7 • Na transfers 1 electron to Cl • Lewis structure for NaCl: Na^+[Cl]^- • MgO: Mg^2+[O]^2- • CaF2: Ca^2+2[F]^- Lewis Structures for Molecular Compounds • overlapping orbitals is a single covalent bond or sigma bond —ex. H2: single bond H—H • electrons shared between elements can be represented with bonds —ex. H2O: dots between H’s and O are single bonds, each single bond=1 electron pair —H—O—H —add 4 dots around O to satisfy octet rule • dots can be bonding pairs (bonded to another element) or nonbonding pairs (solitary electron pair) • steps for writing molecular Lewis structures —add up valence electrons (add/subtract electrons for polyatomic ions) —place atoms with single bonds A. central atom=element with greatest bonding capacity B. if same bonding capacity, choose element with lower electronegativity —complete octets by adding dots to elements (start with surrounding atoms) until the number matches the total number of valence electrons —if valence electrons run out before octets are complete, convert lone pair(s) to bonding pair(s) with double or triple bonds —if extra valence electrons, add lone pairs to central atom (expanded octet exception) • ex. O2 —6x2=12 valence electrons —O—O —dots run out, so use double bond —O=O, 4 dots around each O to complete octet Electronegativity and Lewis Structures cont. Intro/Bond Polarity • electronegativity - relative ability of atom in a bond to attract shared electrons • trends —higher EN—> greater attraction —EN increases—> Z eff. increases (across a row) —EN decreases—> Z eff. decreases (down a group) —EN of H goes between B and C —F is most electronegative element • the larger the EN difference, the more polar the bond is • non polar bonds have same EN Lewis Structures for Polyatomic Ions • group of atoms, covalent bonds within polyatomic ion, overall + or - charge • ex. OH- —add up valence electrons including charge (6 for O + 1 for H + 1 for charge=8 valence electrons) —same procedure as Lewis structures for covalent molecules —include brackets around Lewis structure with charge Resonance • resonance structures - molecule can have more than one Lewis structure —resonance structures must have at least one double or triple bond —all atoms must be in same arrangement and satisfy octet rule —different arrangement of bonding pairs (and lone pairs) —resonance represented by double headed arrow between structures —actual structure is average of equivalent resonance structures • resonance stabilization - delocalized (spread out among atoms) electrons decrease PE and make atom more stable Formal Charge • formal charge - unreal charges used to choose more important resonance (or Lewis) structure —charge an atom would have if all electrons were shared equally • calculating formal charge —# valence electrons in free atom - # electrons assigned to atom in structure (# unshared electrons + 1/2 # shared electrons) —ex. H2O A. draw Lewis structure B. H—1-1=0 C. O—6-6=0 D. H—1-1=0 —if molecule is an ion, formal charges for all atoms must add up to value of charge • when using formal charge to compare structures, the best Lewis structure has (in order of importance): —1st: all formal charges=0 —2nd: most formal charges=0 or closer to 0 —3rd: negative formal charge on more electronegative element Naming Ionic Compounds with Polyatomic Ions • important polyatomic ions —ammonium: NH4+ —hydronium: H3O+ —acetate: CH3COO- —carbonate: CO3 2- —cyanide: CN- —nitrate: NO3- —nitrite: NO2- —phosphate: PO4 3- —sulfate: SO4 2- —sulfite: SO3 2- —hydroxide: OH- • oxoanions - polyatomic ions with different numbers of oxygens • ex. NaOH—sodium hydroxide • ex. copper (I) sulfate—Cu2SO4 • ex. ammonium phosphate—(NH4)3PO4 VSEPR and Hybrid Orbitals Molecular and Electron Geometry • valence shell electron pair repulsion theory (VSEPR) assumes that each atom in a molecule will have a geometry that minimizes repulsion between valence electrons and is used to predict geometry of molecules from the number of electron pairs around central atoms • the steric number (SN) is the number of atoms bonded to central atom + number of lone pairs on central atom • molecule shape varies with number of bonded atoms • types of geometry — linear: 2 bonded atoms with no lone pairs (SN=2) — bent: 2 or 3 bonded atoms with 1 lone pair (SN=3) — trigonal planar: 3 bonded atoms in a single plane with no lone pairs (SN=3) — trigonal pyramidal: 3 bonded atoms with 1 lone pair (SN=4) — tetrahedral: 4 bonded atoms in 3D space with no lone pairs (SN=4) • SN=3 is always trigonal planar electron geometry • SN=4 is always tetrahedral electron geometry • bond angles —SN=2: 180˚ —SN=3: 120˚ —SN=4: 109.5˚ • sigma bonds occur when orbitals overlap • orbitals can hybridize to form new orbitals and molecule arrangements • hybrid orbitals —SN=4: 1 s and 3 p—>4 sp3 orbitals —SN=3: 1 s and 2 p—>3 sp2 orbitals —SN=2: 1 s and 1 p—>2 sp orbitals • hybrid orbitals are formed so that valence electrons are distributed evenly —ex. CH4: SN=4, tetrahedral —each H has 1 valence electron in 1s —C has 4 valence electrons, 2 in 2s and 2 in 2p —C hybridizes to form 4 sp3 orbitals with 1 electron in each so that they can sigma bond to the 4 H atoms • pi bonds are formed when 2 lobes of one orbital overlap 2 lobes of the other involved orbital in a parallel manner —a pi bond and a sigma bond form a double bond (1 unhybridized p orbital) —2 pi bonds and a sigma bond form a triple bond (2 unhybridized p orbitals)


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