AP Chemistry To: 2013-2014 AP Chemistry students From: Big Evergreen III, a past student of o f Mrs. Moses SUBJECT : Hints / strategies / review to survive AP Chemistry.
Think categorically : know acid from base, strong from weak, metal from nonmetal, ionic from covalent, etc. Know your nomenclature , you will need it for everything. When you are taught something, learn it : many concepts are reused within other concepts so if you didn’t learn something in the first place it will hurt you later. Know your solubility rules: rules : this will help so much on the reactions part of the exam! Learn the details and the relationships: How does temperature, pressure, etc. affect a certain system? What does it mean if something is solid, liquid or gas? What are certain numbers dependent upon?
Table of Contents 1. Atoms, Molecules, and Ions........................................................................................................... Ions...........................................................................................................
2. Stoichiometry................................................................................................................................. Stoichiometry ................................................................................................................................. 3. Reactions ........................................................................................................................................ 4. Gases .............................................................................................................................................. 5. Thermochemistry ........................................................................................................................... 6. Atomic Structure & Periodicity ..................................................................................................... 7. Fundamentals of Chemical Bonding ................................................. ............................................. 8. Theories of Chemical Bonding ...................................................................................................... 9. Liquids & Solids ............................................................................................................................ 10. Properties of Solutions ................................................... .................................................... .............................................................. .......... 11. Kinetics ........................................................................................................................................ 12. Chemical Equilibrium .................................................................................................................. 13. Acids and Bases ........................................................................................................................... 14. Aqueous Equilibria ...................................................................................................................... 15. Spontaneity of Chemical Processes ............................................................................................. 16. Electrochemistry .......................................................................................................................... 17. Nuclear Chemistry & Radiochemistry ..................................................... ......................................................................................... .................................... 18. O-Chem aka Organic Chemistry ...................................................... .................................................................................................. ............................................
Atomic Structure and Periodicity Parti Part i cles and Waves Waves El ectromagnetic ctromagnetic Radi Radi ation
Much information about atomic electronic structure was obtained from studies on the interaction of electromagnetic electromagnetic radiation with matter. • Electromagnetic radiation carries energy through space & has a wavelike nature. E.g. light, x-rays, microwaves • Each wave has a characteristic wavelength and frequency. wavelength, λ: distance between wave peaks. Units: m frequency, ν: # of cycles (complete waves) that pass a point in one second. Units are hertz. 1 Hz = 1 s-1 s -1 8 In a vacuum, all electromagnetic radiation travels at a speed of 3.00 x 10 m/s. This is the speed of light, c. c = νλ (speed νλ (speed of light = frequency X wavelength) -1 units: m/s = s m frequency and wavelength are inversely proportional
Practice Problems 1. Calculate the frequency of an X ray that has a wavelength of 8.21 nm. -9 (hint: 1nm = 10 m) Step 1: Write the formula to find frequency if wavelength is given c = νλ Step 2: Manipulate the formula so that you’re you’re solving for frequency c = νλ c/λ = ν/λ ν = c/λ Step 3: Convert 8.21nm to m so that we can cancel out the units in the train-and-caboose train-and-caboose (this is kind of a given, but I just added it as a step anyway…) -9 -9 8.21 nm x 10 m = 8.21 8.21 x 10 m 1 nm Step 3: Plug in the values and chug 8 -9 ν = (3 x 10 m/s) / (8.21 x 10 m) 16 16 -1 = 3.65 x 10 1/s or 3.65 x 10 s
Now, try the rest rest on your own! 2.
3.
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Calculate the wavelength, in nm, of infra-red radiation that has a 13 -1 frequency of 9.76 x 10 s (the answer should be in nm, not m) Calculate the frequency, in hertz, of a microwave that has a wavelength of 1.07 mm (that’s right, mm not nm!) Answers:
2) 3.1 x 10 3 nm
3) 2.8 x 10 11 s-1
El ectromagne ctrom agneti ti c Spe Spectru ctru m
A "continuum" of all possible wavelengths of electromagnetic radiation Humans only see a small part of the whole EMS The region we can see is called “visible light” (our (our colors)
What you need to know from the Electromagnetic Spectrum (EMS) 1. Memorize each type of wave and know how they are arranged from lowest to highest energy (Radio, TV, Micro, Infra-red, Visible Light (ROYGBIV), UltraViolet, X, Gamma) and their frequency 2. Know that low energy corresponds to long wavelengths and low frequency, whereas high energy corresponds to short wavelengths and high frequency 3. Wavelengths in the visible region range from about 400 nm to 700 nm 4. Left -> Right is from lowest energy to highest energy Question: 1. Which light has a higher frequency: the bright red brake lights of an automobile or the faint green light of a distant traffic signal? Green light has a higher frequency than Red light, which means it has more energy
Quan tiz ti zed En er gy
1900 - Max Planck introduced Planck introduced the theory of “quantum packets of energy”. This theory states that energy can only be absorbed or released from atoms in discrete quantities or “bundles. “bundles.” He called the smallest bundle of energy a “quantum.” Thus, E is quantized, not continuous.
Just like an atom is the smallest smallest piece of an element, a “quantum” is the smallest amount of energy you can gain or lose
A “quantum” of energy (E) = hv
And since c = λ ν, A “quantum” of energy (E) = h times c
h = Planck’s constant = 6.63 x 10 hν = smallest amount amount of energy
-34
J-s
Practice Problems 1. Calculate the energy (in Joules) and the frequency (in Hertz) of electromagnetic radiation radiation that is given off by a sodium vapor lamp if the wavelength of the radiation that is 515 nm. Step 1: Write the correct formula to find frequency c=λν Step 2: Manipulate the formula so that you’re solving for frequency ν = c / λ Step 3: Calculate frequency
= 5.83 x 10
14
-1
14
sec or 5.83 x 10 Hz
Step 4: Write the correct formula to find Energy of electromagnetic radiation radiation E = hv Step 5: Use Planck’s constant and the calculated frequency to find Energy -34 . 14 -19 E = hv = (6.63 x 10 J s)(5.83 x 10 / s) = 3.87 x 10 J
Try these on your own. 2. Calculate the smallest increment of energy that can be emitted or absorbed at a wavelength of 645 nm. 3. What frequency and wavelength of radiation has photons of energy 8.23 -20 x 10 J? What type of electromagnetic radiation radiation is this (refer to the EM spectrum). Hint: find frequency first from the Energy formula(E = hv) and then plug the calculated frequency into the wavelength and frequency formula(c = λν).
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Answers:
2) 3.08 x 10 -19 J
3)2.41 x10 -6 m, infrared radiation
The Photoelectric Effect 1905 - Einstein used Plank’s theory to explain the photoelectric effect. effect. He assumed assumed light traveled in energy packets called photons. 1 photon = smallest increment of radiant energy • Energy of 1 photon: E = hν • Thus light has both wave-like wave-like characteristics characteristics (EM studies) & particle nature (Planck & Einstein) a.k.a wave-particle duality of light •More intense light would have more photons and thus eject more electrons, whereas higher frequency light would have more energy and give the ejected electrons more energy
Here are some super awesome flashcards ~ http://www.funnelbrain.com/fc-14325-line-spectrum.html
Line spectra • When white white light is passed through a prism, it separates into a continuous spectrum of all wavelengths of visible light. (ROYGBIV) • When the light from a heated element passes through a prism, a line spectrum with distinct lines is observed. Each line corresponds to a specific wavelength of visible light. • Each atom has its own unique line spectrum. BOHR MODEL MODEL ~1913 • Bohr tried to explain observed line spectra based on movement of electrons. • Niels Bohr used the "planetary model” of the atom in which electrons orbit the nucleus like planets orbiting the sun to explain the phenomenon of "line spectra" (at least for hydrogen) • charged electrons travel rapidly in orbits around the tiny + charged nucleus • Based on Planck's and Einstein's research, Niels Bohr proposed that the energy possessed by electrons was also also "quantized". "quantized". Therefore, an an electron can only be located in specific orbits (energy levels) and not just anywhere within the electron cloud 1. Electrons are contained in specific energy levels called orbits. These energy levels are quantized which means only certain energies are allowed. An e in a permitted orbit has a specific energy. Energy levels are designated by the principal quantum number, n. n = 1,2,3… n = 1 is ground state level - this is level closest to nucleus (lowest in E). The equation below shows how much energy an electron el ectron will have based on its location in the electron electron cloud (the (the energy level it it is on)
where...
E = energy of an electron -18 R H = Rydberg constant = 2.178 x 10 J n = the energy level of the electron
electrons would have quantized amounts of energy so they could either be in the first energy level (n = 1) or the second energy level (n = 2) or the third energy level (n =3 )... and so on but they could not exist between these levels. Practice Problem
Calculate the amount of energy an electron must have a) to be in the 1 st energy level of a hydrogen atom and b) to be in the third energy level of a hydrogen atom a) b)
-18 E= = = -2.18 x 10 J = -2.42 x 10-19 J = E=
2. Electrons circle the nucleus at s pecific radii. (r α n2) n2) 3. Electrons can jump from one level to another by absorbing or emitting photons of specific specific frequencies. frequencies. (e must gain energy from heat, radiation, etc. to jump to higher level.) An electron at n=1 is in its “ground” state. In order to jump to a higher energy level (an “excited” state), the electron must absorb energy in the form of photons of light. Since everything wants to be low energy, an excited electron will eventually transition from an excited state to a lower energy level (moves closer to the nucleus). To do this, the electron must release/emit energy in the form of photons of light. This emission is the cause of line spectra! The emitted energy has a specific frequency(E=h frequency(E=h ) and wavelength(c wavelength(c = λν) λν) that corresponds to a specific part of the electromagnetic spectrum. spectrum. If it falls in the visible portion of the spectrum, spectrum, we see it as a colored s pectral “line”. Mystery Mystery Solved!
Bohr's theory explains 4 observed lines in line spectra for hydrogen. Lines correspond to emitted radiation in visible portion of the EM spectrum when e jumps from 1 level to another.
This process is responsible for colors of fireworks & neon signs. Electrons are excited by heat or electricity and electrons jump to higher energy levels. Light is emitted when electrons lose energy and drop back down to lower energy levels. The colors correspond to wavelengths of emitted light waves. • Calculating ΔE (the transition energy) This tells you how much energy must be absorbed or emitted to move from one level to another 1. ΔE = Efinal – E – Einitial =
2. (-) J means that energy was released/emitted released/emitted
Practice Problem Calculate ΔE when an electron moves from n=5 to n=2. Step 1: Write out the formula you’re going to use ΔE = Efinal – E – Einitial = Step 2: Plug & Chug
-19 ΔE = Efinal – E – Einitial = = -4.58 x 10 J
How does Bohr’s model of the atom explain the concept concept of line spectra? Planck said energy is quantized Einstein said light is energy so it must also be quantized (photons) Bohr said the amount of energy an electron has must also be quantized So when excited electrons move to a level that is closer to the nucleus by emitting energy, the energy they emit is represented by ΔE = hv which shows the frequency of the emitted energy. With a known frequency, the wavelength can be calculated (using c = λν) λν) and we can use the electromagnetic electromagnetic spectrum to determine what type of energy was emitted. If it is in the visible region, we can see a color (a spectral line). Practice Problem
Calculate the frequency and wavelength of light emitted in a hydrogen atom when an electron goes from n=5 to n=2. Step 1: Find change in energy ΔE = Efinal – E – Einitial = Step 2: Find frequency
= -4.58 x 10
-19
J
ΔE = hv
14 -1 V= = 6.90 x 10 s
Step 3: Find wavelength
c = λν
-7 λ= = 4.3 x 10 m
Step 4: Convert meters to nanometers -7 4.3 x 10 m x 1 nm = 430 nm -9 10 m
***So, excited electrons transitioning from higher to lower energy levels will emit/release/lose emit/release/lose photons of specific wavelengths and frequencies that are often visible as "colored lines" (line spectra) when viewed through a prism.
These are the known transitions that can be made by electrons in a hydrogen atom. You DO NOT need to memorize them, but be aware of the scientist’s names that were used to name them.
WEAKNESS IN BOHR’S BOHR’S THEORY Bohr’s calculations fell apart when applied to atom others than hydrogen (atoms with more than one electron) Bohr did not take into consideration consideration the interactions that take place when other electrons are present o Electrons repel other electrons o Electrons “shield” each other from nuclear attraction Interactions with other electrons alter the amount of energy required by electrons to transition from one level to another
The Particles and Waves section is done Do these in order to reinforce what you just learned: 1. https://staf https://staff.rockwood.k12.m f.rockwood.k12.mo.us/grayted/ o.us/grayted/apchemistry/D apchemistry/Documents/U5% ocuments/U5%20At 20At omic%20Structure/PROBSET%201%20Transition%20Energy.pdf (only 2 problems!! problems!! ) 2. https://docs.google.com/viewer?a=v&pid=sites&srcid=bHNuZXBhbC5jb21 8YXAtY2hlbWlzdHJ5fGd4Ojc4NTlkYjFjNzNmYjUwMzc (only do questions # 6 and 48(part a- i and ii, part b- i, ii, and iii(you’re smart, I know you can think of a plausible reason for part iii!!)) 3. https://staf https://staff.rockwood.k12.m f.rockwood.k12.mo.us/grayted/ o.us/grayted/apchemistry/D apchemistry/Documents/U5% ocuments/U5%20At 20At omic%20Structure/PROBSET%205%20Test%20Prep%20Atomic%20Struct ure.pdf (do ure.pdf (do #1, 2, 3, 4, 5, 6, 7, 8, 9, 14, 15, 16, 16 , and 23(if you want but I’ll explain how to do this at the tutoring sesh) you can do the other ones too, but some of them are a bit redundant) I’ll post a review key for these questions when I get a chance. By the end of the semester, you’ll be able to to answer all the question on the review don’t let these questions intimidate you!
Quantum N umbers umbers
THE QUANTUM-MECHANICAL MODEL OF THE ATOM This is how we envision the atom today Planck noted that excited matter emits energy Einstein showed that energy acts like both a wave (it has frequency and wavelength) and a particle (bundled into packets called photons) Louis DeBroglie wondered, “If light waves can act like particles of matter, could particles of matter act like waves?” o DeBroglie discovered that electrons could behave like waves… a standing wave (like a vibrating guitar string). The particle and wave properties are are related by:
λ = λ =
where h = Planck’s constant
m = mass of particle v = velocity Erwin Schrodinger developed an equation that describes electron behavior as both a particle and a wave (the equation has a wave function, Ψ) 1. When you square square the wave function, Ψ, you get a 3-dimensional 3-dimensional probability map map that describes regions in space (the electron electron cloud) where electrons are likely to be. We call these regions “orbitals” (This relates to QUANTUM NUMBERS!!! I really liked quantum numbers obvs) 2. He also set-up and solved a series of complex equations that took into account: KE (kinetic energy) of an electron Wavelength of an electron Attraction of electron for nucleus Repulsions between electrons Heisenberg Uncertainty Principle Both the position and the momentum of an electron cannot be exactly known at the same time (ask Moses to give her example about the speed ticket for this principle, it’ll help you remember) 1. In order to known either the location or momentum of an electron, light must hit the electron and bounce back to your your eye or measuring measuring device… however, the light that hits the electron will cause it to change momentum and/or location… it’s a nonowin situation
QUANTUM NUMBERS (eep, NUMBERS (eep, i’m so excited! i loved quantum numbers last year!) Mathematical solutions to Schrodinger’s Schrodinger’s equation are associated with a set of 3 quantum numbers 1. Principal quantum number, n 2. Azimuthal quantum number, m 3. Magnetic quantum number, ml They are kind of like an “address” that tells us where each electron electron lives within a given atom Every electron within an atom has a unique set of quantum numbers
Principal Quantum Number (n) Generally referred to as the “shell” Tells the average distance from the nucleus an electron is 1. As “n” becomes larger, the radius radius that the electron can travel away from the nucleus gets larger 2. Think of “n” as similar to the Bohr energy levels 2 Each shell can hold a max of electrons equal to 2n 1. The first four shells shells can hold 2, 8, 18, and 32 electrons, electrons, respectively
Angular Momentum Quantum Number (l (l ) Generally referred to as the “sublevel” or “subshell” (i’m going to call them the sublevel) Tells the shape of the orbital
Allowed values of l of l = = 0, 1, 2, 3, … o The number of sublevels possible in each shell is equal to the value of n for that subshell rd The 3 subshell (n = 3) may contain a maximum of three sublevels o The value of l of l can can never be greater than n – n – 1 1
Principal Level / Shell (n) 1 2 3 4
Sublevels in the Atom Sublevel Number, l 0 0,1 0,1,2 0,1,2,3
Sublevel Letter s s, p s, p, d s, p, d, f
Problem: If n=3, what are the allowed values of l of l Answer: 0, 1, 2 (s, p, d)
Magnetic Quantum Number Generally referred to as the “orbital” orbital” Tells the orientation of the orbital in space Any orbital can hold a maximum of two electrons YOU MUST KNOW HOW MANY ORIENTATIONS ARE POSSIBLE FOR EACH ORBITAL o 1 for s o 3 for p o 5 for d o 7 for f The number of orbitals that a sublevel may have depends on the azimuthal quantum number, l , of the sublevel and is equal to 2l + 1 Orbitals in the Atom Sublevel Sublevel letter Number of Number of number, l orbitals, 2l 2l + + 1 electrons per sublevel 0 s 1 2 1 p 3 6 2 d 5 10 3 f 7 14 Allowed values ml = -l -l … -3, -2, -1, 0, 1, 2, 3, … l o Orbital s = 0 o Orbital p = -1, 0 +1 o Orbital d = -2, -1, 0, +1, +2 o Orbital f = -3, -2, -1, 0, +1, +2, +3
Problem: If l If l = = 2, what are the allowed values ml ? Answer: -2, -1, 0, 1, 2 Note that when l = l = 2 that corresponds to a “d” orbital. There are 5 possible orientations of “d” orbitals, right? When l = l = 2, there are also 5 possible values for ml . See the relationship? relationship?
Electron Spin Quantum Number (ms) spin” Generally referred to as the “spin” This quantum number is NOT part of a solution to Schrodinger ’s equation. It was later added to ensure that each electron (even those paired up in the same orbital) has its own unique set of quantum numbers to identify it o Pauli Exclusion’ Exclusion’s Principle: no two electrons in the same atom may have the same four quantum numbers Tells the spin of the electron within the orbital Spinning charges produce a magnetic field. In order for two electrons to exist in the same orbital, they must spin in opposite directions thereby creating opposite magnetic fields that cancel each other out. This minimizes electron-to-electron electron-to-electron repulsion and thus creates a lower energy state that is more stable Allowed values of ms = +1/2 or -1/2
Problem: Write all possible sets of quantum numbers for an electron in a 3p orbital n = 3; l = p” orbital; there are three orbitals, so m l = -1, 0, 1 l = 1 since it is a “ p” (3, 1, -1, -1, ½) (3, 1, -1, -1/2) (3, 1, 0, ½) (3, 1, 0, -1/2) (3, 1, 1, ½) (3, 1, 1, -1/2) A Summary of Schrodinger’ Schrodinger’s “Quantum-Mechanical ” Model of the Atom 1. Electrons don't just move around the nucleus of the atom in simple circular "orbits" as Bohr had predicted. Schrodinger Schrodinger developed developed complex equations equations that describe 3-dimensional regions within the atom where electrons are likely to be found. These regions are called "orbitals", and 90% of the time, electrons will be found somewhere within that region.
2. Possible solutions to Schrodinger's equations result in a set of 3 quantum numbers (n, l , ml ) that describe the size, shape, and orientation of the orbitals, respectively. They are sort of like an "address" for each electron within an atom. 3. According to Heisenberg's Uncertainty Principle, the more you know about an electron's location, the less you know about its momentum and vice versa. Schrodinger’’s Quantum Mechanical Model > Bohr’ Schrodinger Bohr ’s Model
The Quantum Numbers section is done Do these in order to reinforce what you just learned:
1. https://staf https://staff.rockwood.k12.m f.rockwood.k12.mo.us/grayted/ o.us/grayted/apchemistry/D apchemistry/Documents/U5% ocuments/U5%20At 20At omic%20Structure/PROBSET%202%20Quantum%20Numbers.pdf (only 15 questions)
2. https://staf https://staff.rockwood.k12.m f.rockwood.k12.mo.us/grayted/ o.us/grayted/apchemistry/D apchemistry/Documents/U5%20At ocuments/U5%20At omic%20Structure/PROBSET%205%20Test%20Prep%20Quantum%20Nu mbers.pdf (only mbers.pdf (only do a few problems from the first 5 pages)
3. http://www.pwi http://www.pwista.com/Midter sta.com/Midterm%20Revi m%20Review.pdf ew.pdf (there (there are A LOAD of questions on here. But just do #1-12 or just #7-12) I’ll post an answer key answer key for these questions when I get a chance. If I don’ don ’t post ‘em fast enough, bring the questions to tutoring and I’ I’ll give you the answers then. By the end of the semester you’ll semester you’ll be able to to answer all the question on the review don’t let these questions intimidate you!