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MSU / Chemistry / CEM 141 / chem 1007 study guide

chem 1007 study guide

chem 1007 study guide

Description

School: Michigan State University
Department: Chemistry
Course: General Chemistry
Professor: A. pollock
Term: Fall 2015
Tags: cem 141
Cost: 50
Name: Final Exam Study Guide
Description: Everything covered in CEM 141 this year and everything you need to know to ace the final! Stoichiometry problems along with explanations of all chapters.
Uploaded: 12/12/2015
21 Pages 16 Views 6 Unlocks
Reviews

Joelle Wisozk (Rating: )

I had to miss class because of a doctors appointment and these notes were a LIFESAVER



CEM 141 Final Study Guide


Why coffee cools down when milk is added?



Ch. 1

Claim, Evidence & Explanation/Reasoning

∙ Why coffee cools down when milk is added

o Claim: coffee cools when milk is added

o Evidence: temperature of milky coffee is between milk & coffee

o Reasoning: Higher temp. molecules collide w/lower temp. molecules, as they  collide E is transferred, hot molecules slow down & cold molecules speed up

Atoms

∙ Matter is made of atoms

∙ Smallest building blocks, can’t be split

∙ Positively charged nucleus surrounded by e- in orbitals

o Most of the mass is in the nucleus

∙ Contain protons (p+), neutrons (n0) & electrons (e-)

o They are neutral as a whole

Dalton’s Atomic Theory  

∙ Atoms are in constant motion-Brownian motion

∙ Elements are composed of small, indivisible indestructible sphere particles called  atoms


What is neutrons?



∙ Atoms of elements are different from atoms of another element

∙ Compounds are combinations of atoms of two or more elements

∙ Atoms are neither created/destroyed during a reaction (chemical reactions are  rearrangements of atoms)

JJ Thomson Experiment (cathode rays)

∙ 1st subatomic particle discovered was the e-

∙ Ray of particles emerged from cathode toward the other end (anode)  ∙ Particles were deflected by the electromagnetic fields in a direction that indicated  they were negatively charged

∙ The metal that the cathode was made of did not affect the behavior of the ray so, the  nature of the e- is independent of the metal

∙ Findings: the electron in atoms

Thomson’s Plum Pudding Model

∙ Atoms contain e- embedded in the atom  


What is Coulomb’s Law ?



∙ Problem: model didn’t justify Rutherford’s experiment (most alpha particles went  through golf foil, some deflected back) We also discuss several other topics like econ 325

Existence of electrons

∙ Claim: atoms contain e- embedded in atom

∙ Evidence: Thomson’s experiment showed us that particles could be deflected/bent by  magnetic fields

∙ Reasoning: the metal the cathode was made out of had no effect on the rays being  bent, the composition of the ray is independent of the element it came from Ernst Rutherford Experiment (gold foil)

∙ Radioactive source emitting alpha particles (He atom/nucleus) with lots of E at gold  foil

∙ Most alpha particles went straight through (some deflected back) Don't forget about the age old question of trochanter rolls

∙ The positive charges that came close to the nucleus were deflected o Evidence for: positive charge must be in the nucleus If you want to learn more check out pasonality

∙ Findings: there is a small dense positively charged nucleus in the center of atoms

Rutherford’s Planetary Model  

∙ Orbits contain e-

∙ Problem: e- aren’t stable & would emit E & orbits would collapse

Existence of Small, Dense, Positively Charged Nucleus If you want to learn more check out ksu dwl

∙ Claim: atoms contain a positive nucleus

∙ Evidence: Rutherford’s experiment showed some atoms were deflected off gold foil  and therefore had a positive nucleus

∙ Reasoning: those positive charges that came close to the nucleus were deflected,  most of the volume is where e- exist

Neutrons

∙ Last subatomic particle discovered > they are harder to detect because they have no  charge/neutral

∙ Located in nucleus, slightly heaver then p+

Model of Atom we use now

∙ Electrically neutral  

∙ Cloud of e- take up most of the space

o Cloud can be shifted, e- move

Coulomb’s Law

∙ Positive & negative (opposite) charges attract & like charges repel Don't forget about the age old question of calcium ion storage organelle

∙ Force of attraction= (q1 x q2)/ r2 

o Charges on particles (q)

o Distance between particles (r) We also discuss several other topics like who marries ino

Gravitational Forces vs. Electromagnetic Forces

*Similarity for both forces

Gravitational Forces (gravity)

Electromagnetic/Electrostatic Force

Attraction between objects that have  mass

Attraction/repulsion between objects  w/electric charge

Mediated by fields*

Mediated by fields*

Requires 2 or more objects*

Requires 2 or more objects*

Weaker

Stronger

Always attractive

Can be attractive/repulsive

Decreases as distance increases*

Decreases as distance increases*

Holds us on the earth

Increases as charge increases

Stops us from falling through the earth

Throwing/Holding a Ball:

∙ When the ball is in the air, the only force acting on it is gravitational force ∙ Force of attraction:

o Increases as mass of interacting objects increases

o Decreases as distance between objects increases

∙ When holding the ball, forces acting on it are gravity & electrostatic

Kinetic Energy (KE)

∙ KE= 12m v2 E associated with motion

∙ As ball moves toward the ground, KE increases (velocity is increasing)

Potential Energy (PE)

∙ E associated w/ the position of system of objects in field

∙ Can’t have PE without a field, or without 2 or more objects

∙ As ball moves toward the ground, PE of system (ball & earth) decreases ∙ Gravitational interaction between earth & ball:

o Cause PE to decrease

o Cause ball to fall down

∙ Total E of the system (ball & earth) stays the same

*E can be transferred from one object to another & E can be transformed (from PE to KE or  KE to PE)*

What Makes Atoms Stick Together

Solid Liquid Gas -Touching in very organized -touch, but there is more -Completely  separated

& ordered state disorder & they take the

shape of the container

London Dispersion Forces (LDF)

∙ Forces are not very strong, they are a type of intermolecular force (IMF) ∙ Caused by fluctuation of e- density in molecule/atom

∙ Between 2 atoms/molecules

∙ Partially positive side of atom being attracted to partially negative side on the other  atom

∙ As atoms approach KE increases (PE decreases) until they get too close ∙ If they get too close their e- clouds repel each other

∙ Larger e- cloud= stronger LDF

o Larger/floppier e- cloud is stronger because their protons can get closer  together

He atom Interactions (noble gas)

∙ Held together by LDFs

∙ When they hit the wall of a container total E decreases

o It transfers to the wall

∙ E transferred through collisions

∙ Increased temp = move faster = more collisions = more E = break apart

∙ Why they move toward each other:

o Attractive electrostatic force

o LDF

∙ Why they oscillate:

o Cloud of e- repel each other and LDF attracts them to each other

∙ To keep them close together:

o Lose E by bringing in 3rd atom

o E transferred by collisions to 3rd atom

∙ To form stable interaction:

o Remove E

∙ Increasing temperature:

o Increase E (move faster)

o Knew it was warmer by colliding, E is transferred from container to atoms o If enough E is transferred to the atoms they break apart

o If thermal E increases, temperature increases

He Atoms- PE Graph

Atoms Approaching/coming closer together:

∙ Electromagnetic attraction & LDF bringing them together

∙ KE increases, PE decreases

At the well:

∙ Electromagnetic attraction = repulsion

∙ Most stable

∙ Depth of well tells you:

o How strong the interaction is & how much E is needed to get the atoms out of  well

o Deeper well requires more E

o Deeper wells have higher boiling points

∙ Position of well (left/right) tells you:

o Internuclear distance- distance between atoms at the most stable point o Farther right the well is means they have a larger atomic radius

Atoms Very/Too Close Together:

∙ The e- clouds start to overlap and the negative charges repel each other ∙ PE increases, KE decreases

He vs. Xe Atoms

∙ 2 Xe atoms can’t get as close together as 2 He atoms because their e- clouds will  repel

o The LDF are larger between the Xe atoms then the He atoms because they have more e- so they are further away from the nucleus making the e- cloud larger o Xe has larger LDFs

o Xe has larger boiling point (stronger LDFs require more E to break the forces)

Breaking bonds

∙ E is absorbed (when atoms absorb enough energy the forces between them are  overcome)

Forming bonds

∙ E is released (when atoms lose E they come closer together and do not have enough  E to overcome to attractive force between them so they stick together)

H Atom Interactions

∙ 2 H atoms approach w/stronger attraction then He atoms it forms a covalent bond ∙ When something collides into H atoms, E is transferred from H atoms to that atom &  that forms a covalent bond between them

∙ 2 H atoms= H2 molecule

H2 Molecules (H-H)

∙ When stable, H2 stays in PE well

∙ When they collide and absorb enough E, H2 separate into H atoms ∙ Takes 6000 K to break 2 H atoms and 4K to break 2 He atoms

o 2 He are held together by weaker interactions (LDF)

∙ Under normal circumstances H exists as H2 diatomic molecules

o Other examples: O2, N2, F2, Cl2, Br2, I2

H @ 5K H @ 15K H @ 30K

*H @ more than 6000K looks like gas monoatomic picture

Stoichiometry Problems

Stoichiometry

∙ Allows us to calculate how much can be produced in a reaction ∙ Coefficients tell you mole ratio

Mole

∙ 6.022 x 1023 

∙ Use for mass to mole conversions

Stoichiometry Calculations:

1. Write a balanced equation

2. Draw map of where you start & how you get to the end

3. Write out calculation w/units

4. Calculate

Balance the equation: C6H14 + O2 > CO2 + H2 O

 2 C6H14 + 19 O2 > 12 CO2 + 14 H2 O

Calculate the molar mass of: H2 O

2H = 2 x 1 g/mol = 2 g/mol

1O= 1 x 16 g/mol = 16 g/mol 2+16= 18 g/mol

 ^atomic mass

Mass > Mole conversions

Ex) How many moles of (Ca( NO3¿2 ) are in 325 g?

1 Ca= 1x40 g/mol

2N= 2x14 g/mol 325 g Ca( NO3¿2 x 1 mol 

164 g = 1.98 mol Ca( NO3¿2 

6O= 6x16 g/mol

^= 164 g/mol

 How many moles of O atoms are in the Ca( NO3¿2 ?

NO3¿2 

1.98 mol Ca( NO3¿2 x  

1molCa ¿ 6mol O ¿

 = 11.88 mol O

Ex) Which is biggest? 10g C H4, or 10g C2H6 ?

10 g C H4 x 1 mol

16 g = .625 mol C H4 

10 g C2H6 x 1 mol 

30 g = .333 mol C2H6 

Ex) How many moles of NH3 would be produced if 6 moles of H2 reacted w/excess N2 ?

3 H2 + N2 > 2 NH3 6 mol H2 x 2mol NH3

3mol H2 = 4 mol NH3 

 How many grams of N2 would be required to produce 2.75 moles of NH3 ? 2mol NH3 x 28g N2 

2.75 mol NH3 x 1mol N2 Limiting Reagents

1 mol N2 = 38.5 g N2 

Ex) Find the limiting reagent starting with 3 moles C2H 4 O and 5 moles H2O to make C2H6O2.

C2H 4 O + H2O > C2H6O2 

3 mol C2H 4 O x 1 molC2H 6O2 

1molC2 H4O = 3 mol C2H6O2 * C2H 4 O is the  

limiting reagent because it can  

only make 3 mol C2H6O2 

5 mol H2O x 1 molC2H 6O2 

1mol H2O = 5 mol C2H6O2 

Ex) How many moles of NH3 would be produced if 6 moles H2 reacted with 3 moles N2 ?

N2 + 3 H2 > 2 NH3 

3 mol N2 x 2mol NH3 

1mol N2 = 6 mol NH3 

6 mol H2 x 2mol NH3

3mol H2 = 4 mol NH3 * would be produced because it is the  limiting reagent

Using the equation: CH4 + 2 O2 > CO2 + 2 H2O

Ex) How many moles of CO2 would be produced from 1 mole of O2 ? 1 mol O2 x 1 molCO2 

1molO2 = .5 mol CO2 

 What mass in grams of CO2 would be produced from 16g O2 ? 16g O2 x 1 molO2 

32 gO2 x 1 molCO2 

2mol O2 x 44gCO2 

1 molCO2 = 11g CO2

 If you had 16g O2 and 16g CH4, how much water would be produced in grams? 32 gO2 x 2mol H2O

16g O2 x 1 molO2 limiting reagent

2 molO2 x 18 g H2O

1 mol H2O = 9g H2O *because it is the  

16 gCH4 x 2mol H2O

16g CH4 x 1 molCH4 

1molCH4 x 18g H2O

1 mol H2O = 36g H2O

Percent Yield = actual amount produced (given) 

theoretical amount(calculated) x 100

 If the reaction produced 8g CO2 what would the percent yield be?

8 gCO2 

11 gCO2 x 100 = 73%

 If you had 10g CH4 and 10g O2, what is the maximum amount of CO2 that could  be produced?

10g CH4 x 1 molCH4 

16 gCH4 x 1molCO2 

1 molCH4 x 44 gCO2

1 molCO2 = 27.5g CO2 

10g O2 x 1 molO2

32 gO2 x 1 molCO2

2mol O2 x 44 gCO2

1 molCO2 = 6.9g CO2 *because it is  

the limiting reagent

Ch. 2

Electromagnetic Spectrum

∙ Log scale (2 objects 102 m apart are 100m apart)

∙ Highest Frequency (Lowest λ)

o Gamma

o X-rays

o UV

o Visible Spectrum

 Blue

 Red

 *Red has longer wavelength then blue

o IR

o Radio

∙ ^Lowest frequency (highest λ)

Light is a Wave

∙ Wavelength λ- distance from peak to peak

∙ Frequency v (Hz)- # of wave fronts per second

∙ Amplitude- height of peaks (intensity)

∙ C = λv, c=3x 108 m/s

∙ Energy increases as frequency increases (wavelength decreases)

∙ Large wavelength has small frequency & small amplitude

Ex) Determine λ (in nano meters) of an x-ray with a frequency of 3 x 1018 Hz (3 x 1 08 m/s) = (λ)( 3 x 1018 ) 1 x 1 0−10 m x 1nm

1 x 1 0−9m = 1 x 10−1nm

λ= 1 x 1 0−10 m

Diffraction

∙ When waves hit a barrier with a slit, the wave that goes through the slit is diffracted

Interference

∙ Constructive

o Two waves the same reinforce & create higher crests & troughs

∙ Destructive  

o Two opposite waves cancel out (become straight line)

Electromagnetic Radiation as a Wave

∙ Claim: E/m radiation is a wave

∙ Evidence: E/m radiation is diffracted & shows patterns of interference ∙ Reasoning: When e/m goes through a barrier w/ a slit the resulting waves are  diffracted, they can be combined either constructively or destructively

Photoelectric Effect

∙ Metals emit e- when light shines on their surface. The certain threshold frequency  must be met otherwise no e- will be ejected no matter how intense the light is. There  must be enough E (a photon) to emit an e-

∙ E is transferred as a photon, each photon has definable E

∙ One photon ejects one e-

∙ If the photon does not have enough E, then no e- is ejected

∙ E= hv, h= 6.626 x 10−34 J /s

Ex) What is the E of a photon of frequency 4 x 1018 

4 x1018 = (6.626 x 10−34 J /s )(E)

E= 6 x 1 051 J

 What is the λ of a photon of E 6.2 x 10−8J

E=hv 6.2 x10−8J

6.626 x10−34 J .s = 9.36 x 1025s−1

Λ=c/v 3 x 108m/s

9.36 x1025s−1 = 3.2 x 1 0−18 m

E/m Radiation is a Particle

∙ Claim: E/m radiation is a particle.

∙ Evidence: The photoelectric effect says that metals emit e- when light shines on their  surface

∙ Reasoning: The light is transferring energy to the e- which is transformed into KE  which gives e- enough E to be emitted. The photons of E absorbed have to be more  than the threshold or no e- will be ejected. If a certain amount of E is absorbed then  the e- will be ejected.  

Electrons in Atoms Have Quantized E Levels

∙ Different atoms contain different e- that each have specific energy levels that are  quantized. It requires the exact amount of energy the e- needs to move to the next  energy level.  

∙ Emission spectra has lines where the e- are moved from higher to lower energy levels  and emit a photon

∙ Absorption spectra have lines where the e- are moved from a lower to higher energy  level from absorbing a photon.

o Spectra are created by absorbing and emitting photons. It takes a specific  amount of energy to move and e- to the next energy level and specific amount  of energy must be emitted to move it down an energy level.

o Spectra show light only of specific energies

Electrons (Wave Properties)

∙ Energy diagrams, e- transition between E levels by absorbing or emitting photons ∙ Atoms absorb/release specific amounts of E to transition, these specific amounts emit  that specific wavelength

∙ Core are right elements which are harder to remove than valence e- bc they are  attracted to the positive nucleus, valence are outermost e- which are the easiest to  remove

Bohr

∙ e- are most likely in their orbital, only works for Hydrogen

∙ Orbits have definite energies so e- in atoms are quantized

∙ E of photons corresponds to the difference in E levels of e-

∙ If e- loses E it is de-excited and moves to a lower E orbit closer to the nucleus ∙ Atoms of elements can’t absorb/emit any λ of light (only certain wavelengths that  correspond to that element)

E Diagrams

∙ Each E level has a quantum #

∙ Higher # = more E, E levels are not orbits

∙ E- transition between E levels by absorbing/emitting photons

Effective nuclear charge

∙ Core e- cancel out the positive charge from the same number of p+ (Zeff: p+ minus  core e-)

Periodic Trends

∙ Atomic radius increases down a row, and decreases across a row left to right ∙ Ionization E: decreases down a group, and increases across the row left to right o When there are more e-, the IE is smaller bc they are easier to remove the  valence e- when there are many of them so it takes less E to remove them ∙ Zeff: increases across a row, high Zeff hold onto their e- more tightly bc their e- are  more strongly attracted to the nucleus (this is why atomic radius decreases across a  row bc they hold onto e- more tightly so they are smaller)

Ch. 3

Big Bang to Atoms

∙ Started from one and it burst into more at very high temperatures, as it cooled quarks and leptons formed, as it cooled further p+ and n0 formed, a few minutes later when  it cooled more H+, D+, He2+ and Li3+ formed through fusion

∙ Atoms from your body come from one star before the big bang

∙ The number of atoms in the universe is constant

∙ Everything in the universe is moving away from us (Doppler effect) ∙ Evidence: CMBR Cosmic Background, red shift (red on spectra has longer wavelengths showing its moving away from us)

Nuclear fusion, fission and radioactive decay

∙ Fusion (adding two nuclei together)

∙ Fission (breaking apart nuclei, forms lighter more stable nuclei +Energy) *U in equations, chain reactions

∙ Decay (emitting particles, alpha, beta or gamma)

Balance nuclear equations:

∙ Mass # on top

∙ Atomic # on bottom (number of p+ and e-)

∙ Isotopes have diff # of n0

∙ E=mc^2

∙ m=(mass defect)(amu)

Atomic

Mass of 1 > Atomic

Massof 2 +  

Atomic

¿of 1−Atomic ¿of 2¿ =

¿Element 1of 1 ¿

determine element

¿Element 2 of 2 ¿

Mass of 1−Mass of 2

Nano-particles, and larger macro scale materials  

∙ Atoms can be joined together and have emergent properties ∙ Surface area to Size ratio affects properties

∙ Boiling/melting points happen when more than one atom interact together  (interactions between particles have to be overcome)

∙ Nano, don’t have a state

Melting/Boiling Points of Diatomic Molecules vs. Diamond

∙ Hydrogen bonds to form H2, 2 e- in the bonding MO make it bond to itself ∙ He, has 2 e- in the anti bonding MO so it cancels out the 2 e- that are in the bonding  MO

∙ To break the covalent bond between 2 H atoms, enough energy has to be added to  raise e- to the anti bonding MO

∙ Covalent bonds (Diamon) have high boil/melt points, molecules (H2) have low  melt/boil points bc it is easier to break the LDFs then the covalent bonds so they are  still in the form of molecules when they are boiled it takes a lot more E to break their  covalent bonds

∙ Diamond/graphite have such high melting points because they would just disintegrate if they were melted, it also takes a lot more E to break the 3D structure of all of the  covalent bonds

Forming Bonds

∙ Electrostatic forces are there (PE) when two atoms are approaching, when their nuclei  get a nuclei in distance apart the strong nuclear force takes over and PE is no longer  acting on the atoms (during fusion)

∙ Forming a bond, E is released to the surroundings

Molecular Orbital Bond Model (metals)

∙ Combine atomic orbitals (n atomic orbitals=n molecular orbitals)

∙ Each orbital can contain up to 2 e-

∙ Lower 2 in model are bonding MO, higher 2 are anti bonding

∙ If anti bonding are filled up this cancels out the bonding between the atoms (that is  why He doesn’t form a covalent bond just LDF)  

Valence Bond Model (everything else)

∙ Atomic orbitals overlap to from bonds

∙ The greater the overlap the stronger to bond

∙ Each bond is made of 2 e-

∙ e- are localized in the bond

Metals

∙ Conduct electricity: e- can move freely around

∙ Malleable: atoms can move w/respect to one another

∙ Shiny: absorbs photons & moves to higher E level then immediately re-emits it &  moves to lower E level

∙ Interacts with many wavelengths so the metal is white/colorless, silvery  

Ch. 4

Diamond

∙ Carbon forms 4 identical bonds

∙ Sp3 hybrid orbital, tetrahedral geometry  

∙ Localized e-, sigma bonds

∙ Hard

o 3D network of strong covalent bonds

∙ High melting point

o Takes lot of E to overcome strong covalent bonds

∙ Doesn’t conduct

o Localized e- in bonds aren’t free to roam & there’s a large “band gap”

Graphite

∙ Each C attached to 3 C, forms sheets of graphite

∙ Sp2 orbitals form sigma bonds and leftover p orbital is a pi bond ∙ Slippery

o Sheets held together by weak LDFs so they slide over each other ∙ Conducts electricity

o Delocalized pi bonds over structure allow e- to move freely ∙ Shiny

o Absorbs and re-emits photons (like metals)

Carbon

∙ Bonds to C, H, O, N, S, P

∙ Emergent properties  

∙ Methane, CH4

o 4 identical C-H bonds equidistant from one another

o C has sp3 hybridization

Lewis Structures

∙ Calculate # of valence e

o H=1, B=3, C=4, N=5, O=6, F=7

∙ How many bonds it usually forms

o H=1, B=3, C=4, N=3, O=2, F=1

∙ Write skeleton using 2 e- for each bond

∙ Make sure each atom (except H) has 8 e- by adding lone pairs ∙ Add multiple bonds (double, triple) if there aren’t enough

∙ Don’t give 3D info

Isomers

∙ # of carbons > # of isomers it has

o 0, 1, 2, 3 > 1 isomer

o 4 > 2 isomers

o 5 > 3 isomers

o 6 > 5 isomers

o 7 > 9 isomers

Sigma bonds

∙ Allow free rotation of bonded atoms

Pi bonds

∙ Not freely rotatable, pi bond would break

Formal Charge

∙ F.C.= (# of valence e-)-(dots or lone pair e-)-(sticks or 12 bonding e-) ∙ NH4 (draw out Lewis structure first)

o FC (N)= (5)-(0)-(4)= +1

VESPR (Valence Shell e- Pair Repulsion)

Groups of e around  

center atom

Geometry

Bond  

Angle

Hybridizatio n

Examples

2 e-

Linear

180

Sp

CO2

3 e-

Trigonal Pyramid

120

Sp2

BF3

4 e-

Tetrahedral

109

Sp3

CH4, H2O

5 e-

Trigonal  

Bipyramid

90/120

Sp3d

PCl5

6 e-

Octahedral

90

Sp3d2

SF6

*For geometry count all groups of e- surrounding that atom (double/triple=1 e-) *For shape count groups of e- surrounding atoms but ignore lone pairs

Geometry/Shape

∙ Ignore lone pairs

∙ To figure out shape:  

o “Take off” lone pairs from geometry and the remaining structure is the shape ∙ H2O (draw Lewis Structure)

o Geometry: tetrahedral (3 bonds, 1 lone pair= 4 e- groups)

o Shape: bent (3 bonds, ignore lone pairs = 3 e- groups)

∙ CO2

o Geometry: linear (2 bonds (double/triple=1), no lone pairs= 2 e- groups) o Shape: linear (2 bonds, ignore lone pairs= 2 e- groups)

Electronegativity

∙ Ability of an element to attract e- to itself when its in a bond

∙ Increases across table left > right

∙ Decreases across table top > bottom

∙ Noble gases don’t form bonds so they can’t be polar

∙ H & C are so similar they we look at them as the same electronegativity

Polar Bonds

∙ 2 atoms with different electronegativities bond

o Unequal sharing of e

o Results in a dipole

∙ Dipole points towards the more electronegative element in the bond o Ex) H-F it would point towards F since F is more electronegative

o Ex) H-C is not polar since they have same electronegativities

∙ Shape is important!

o Dipole arrows can cancel each other out if they are facing the same way

 Ex) CO2, linear: arrows pointing left and right cancel each other out so  CO2 is nonpolar

 Ex) H2O, bent: arrows pointing upward toward O do not cancel out, they  combine to form one larger dipole so water is polar

London Dispersion Forces

∙ Weakest IMF

∙ Temporary fluctuating dipoles

∙ Only force present in nonpolar molecules

Dipole-dipole Forces

∙ In polar substances (along with LDFs)

∙ Stronger then LDFs

∙ The more electronegative element is partially negative and the less electronegative  element is partially positive

∙ Dipole-dipole forces happen when the partially negative element in one molecule is  attracted to the partially positive element in another molecule

o Ex) H-Cl, Cl of one molecule is attracted to the H of another molecule

Hydrogen Bonding (IMF)

∙ Strongest IMF

∙ Compounds with H-bonds also had dipole-dipole and LDFs

∙ Between molecules with H covalently bonded to either O, N, or F and a lone pair of e on either O, N or F

o H hydrogen bonds to the O, N or F with the lone pairs

o (O, N or F)-H {hydrogen bond} :(O, N or F)

o Ex) CH3OH

Strengths of Bonds & IMFs

Strongest >Covalent Bond

Ionic Bond

Hydrogen Bond (IMF)

Dipole-dipole Forces (IMF)

Weakest > London Dispersion Forces (IMF)

Water

∙ Anomalous properties

∙ High mp/bp & specific heat

∙ Water freezes at 0 degrees C, water boils at 100 degrees C

∙ Low vapor pressure

∙ Density of ice < density of liquid water

o Ice takes hexagonal crystal from which creates “holes”

Ionic Bonding

∙ Transfer of valence e- creates ions (each ion achieves noble gas configuration) ∙ Between nonmetals (high electronegativity) and metals (low EN) ∙ Metals lose e- to nonmetals

∙ Metals form cations +

∙ Nonmetals form anions –

∙ Ionic compounds are neutral

∙ Electrostatic forces holds ions together

NaCl

∙ Na+ is bigger then Cl

o Na has more e- so it has more repulsion which = larger size

∙ Forms colorless crystals

∙ Conducts electricity in liquid state but not solid state

Charges on Ions

∙ K > +1 S > -2

∙ Mg > +2 Al > +3

Formulas for Ionic Compounds

∙ Na+ & Cl- > NaCl

∙ Mg 2+ & O 2- > MgO

∙ Ca 2+ & Br - > CaBr2

Lattice Energy

∙ E released when ionic lattice forms from ions in gas phase

∙ Force of attraction, Coulomb’s law

∙ Strong attraction= more charge, smaller ions

∙ Weaker attraction= less charge, larger ions

Ch. 5

Temperature

∙ Measure of hotness

∙ E always moves from the hotter object to the cooler object

∙ 1 drop boiling water vs. 1 bucket boiling water = same temps. ∙ Depends on average KE of molecules

∙ T=KE

Kinetic Energy

∙ KE= 1/2mv^2 KE= 3/2kT

∙ An individual gas particle can not have a temperature

∙ If average velocity increases, temperature increases

∙ Heavier & lighter molecule = same KE

∙ Lighter molecule has higher average velocity

Thermal Energy

∙ 1 drop boiling water vs. 1 bucket boiling water = not same thermal E ∙ Depends on how much material you have

∙ Added through collisions

∙ Moves faster, translate, vibrate/rotate

Maxwell-Boltzmann Distribution

∙ 2 curves, left one is higher and short & right one is shorter but wider ∙ Higher curve: lowest temperature, lowest velocity & highest mass ∙ Lower/longer curve: higher temperature, high velocity & small mass

Gas Molecules

∙ Speed of gas

o Distributed, random amounts of E are transferred through collisions creating a  distribution of speeds

∙ When gas molecules collide they don’t stick together:

o Have a high KE, enough E to overcome the attractive interactions between  them

∙ Why wouldn’t you smell it immediately across a room if you opened container: o Takes time for molecules to move

o Collide with each other and other molecules in air

o Don’t travel in straight lines to your nose, when they collide their direction  changes

o Random movement (Brownian motion)

Heat

∙ Transfer of thermal E from one place to another by collisions

Thermal E

∙ E associated w/motion of particles

Specific Heat

∙ E it takes to raise temperature of 1 gram by 1 degree C

∙ Lower specific heat raises temperature the fastest

o Lower sp. Heat means its easier to raise the temperature  

∙ Water specific heat is high

o Lots of molecules per gram, low molecular weight

o Lots of IMFs

Phase Changes

∙ Solid > liquid > gas

o Absorbs E, breaks interactions

∙ Gas > liquid > solid

o Releases E, makes interactions

∙ Temperature doesn’t change until change of phase is complete

∙ Extra E being added is used to overcome the IMFs

State Functions

∙ Depend on initial & final stakes

∙ Don’t depend on path taken

o Ex) Elevation of campsite

∙ T, E, H, S, G, P, V or altitude

Path Functions

∙ Depend on the path taken

∙ Depend on how the change takes place

o Ex) Distance travelled

∙ Lower q, w

First Law of Thermodynamics

∙ E cannot be created/destroyed

∙ E can only be transferred/transformed

Exothermic

∙ E goes out of the system

∙ q- sign

∙ Heat goes out of the reaction and the surroundings get hot

Endothermic

∙ E goes into the system

∙ q+ sign

∙ Heat goes into the reaction and the surroundings get cold

Enthalpy ΔH

∙ Change in H or q

∙ Equation: q=m x c x change in T (Tf-Ti)

∙ Mass x specific heat x change in temp.

∙ Ex) 50g metal heated to 100 degrees C added to 50g water at 20 degrees C, highest  temp. water reaches is 27.5 degrees C, what is specific heat?

o Water q = (50g)(4.18)(7.5)=1567.5

o Metal q= (50g)(c)(72.5)=3625

o Compute: 1567.5/3625 = .43 J/g C

Entropy

∙ Disorder

∙ Measure of # of possible arrangements for a given state

∙ More possible arrangements = higher entropy

2nd Law of Thermodynamics

∙ For any change the total entropy of universe must increase

∙ You can’t get as much E back as you put in (some is lost-spread out at thermal E) ∙ Δ Stotal = Δ Ssystem + Δ Ssurroundings 

Mixing

∙ Dye mixes because it spreads out (possible arrangements increases) & entropy  increases

∙ It will never un-mix because it is more probable that they will stay mixed

Entropy & Phases

∙ Solid

o Low entropy

o Molecules are fixed in place & can’t move relative to one another (one possible  arrangement)

∙ Liquid

o Medium entropy

o H-bonds, molecules can move relative to one another not completely free  moving though

o Some possible arrangements (more then solid, less than gas)

∙ Gas

o High entropy

o No forces holding molecules together so they are free to move independently  (most possible arrangements)

Ex) Which has more entropy?

Deck of cards in order or Shuffled deck

CaO (s) + C O2 (g) or CaC O3 (s) *making a gas, entropy is  always higher

H2O (l) @ 25°C or H2O (l) @ 50°C *molecules can move faster &  some LDFs  

 are broken allowing more  

arrangements

Hot block or Cold block *entropy increases with temperature ΔH (change in enthalpy) of Phase Change

Which has higher Δ Hvap ?

C H3OH or C H3C H3 

-H bonds -LDF

-Dipole *harder to break all the bonds

-LDF

Δ Ssurroundings = ΔHsurroundings 

T

Freezing

Vs.

Melting

Δ Ssystem -

Equal to

Δ Ssystem +

Δ Hsystem -

Equal to

Δ Hsystem +

Δ Hsurroundings +

Equal to

Δ Hsurroundings -

Δ Ssurroundings +

Less than

Δ Ssurroundings _

*same magnitude just opposite directions

Freezing Water

∙ Exothermic

∙ Δ Ssurroundings depends on temperature

o Higher temp = smaller the effect on entropy

Δ Stotal 

∙ Tells us whether change will occur

o If total entropy increases the change will happen

o If it decreases it will not happen

∙ Use Gibbs free energy to measure

ΔG Gibbs Free E

Solid > Gas

 ΔG = ΔH – T ΔS *T is always positive For this process to happen ΔG must be

 (-) = (+) –(+)(+) negative so you want a  high temperature

 ^depends on how high the temp is.

Gas > Liquid

 ΔG = ΔH – T ΔS *Want T to be small so ΔG can be negative  (-) = (-) - (+)(-)

 ^system loses E when you make those interactions

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