CHEM127 - Study Guide Midterm #1
CHEM127 - Study Guide Midterm #1 Chem 127
Popular in General Chemistry for Agriculture and Life Science I
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This 18 page Study Guide was uploaded by Aenea Mead on Thursday October 13, 2016. The Study Guide belongs to Chem 127 at California Polytechnic State University San Luis Obispo taught by Jennifer Retsek in Fall 2016. Since its upload, it has received 10 views. For similar materials see General Chemistry for Agriculture and Life Science I in Chemistry at California Polytechnic State University San Luis Obispo.
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Date Created: 10/13/16
MASTER VOCABULARY Abso rption: W hen an atom takes in energy, resulting in an electron jumping to a higher energy level. Am plitude (a): Vertical height of the crest of a wave Relates to intensity / brightness Atom: Smallest unit of a chemical element Atom ic Orbital: 3D regions in space where there is a high probability of finding an electron Compound: two or more elements combined Diamagnetic: species with all electrons paired Element: one type of atom Emission: When an electron jumps to a lower energy level, and thus releases energy. Frequency (ν): number of cycles that pass through in a given period of time ● Units: 1 or s −1 = Hertz (Hz) s Group: column in periodic table Ion: An atom that has gained or lost electrons and therefore has a charge Cation: Positively charged ion Anion: Negatively charged ion Isoelectronic: t wo atoms / ions with the same number of electrons and therefore the same electron configuration Isotope: An atom of an element with a different number of neutrons Mole: amount of substance containing the number of atoms as exactly 12 grams of carbon-12 ● Called Avogadro's number ● Equals 6.022 10 * 23 Molar Mass: mass (g) per mole of a substance Paramagnetic: species with at least one unpaired electron Period: row in periodic table Quantum Number: describes an electron's orbit Wavelength (λ): Distance between two set points on a wave (ie. Crest to crest) ● Units: meter ______________________________________________ DIMENSIONAL ANALYSIS Dimensional analysis is a convenient way to convert quantities from one unit to another. It would be used in an instance where two volumes need to be added, but they are in different units. Dimensional analysis would help to put them in the same unit, so that they can then be added. It is like how fractions cannot be added when they have different denominators, so in order to add them a common denominator needs to first be found. Dimensional analysis could also be used to solve more complex problems such as determining the density (in g/cm )of a solid when only the diameter (in nm) of each atom and its mass (in amu) is known ahead of time. Above is the what and why, now here is the how: Question: How far (in miles) will a person run during a 5 kilometer race? Solution #1:km 0.62mi = 3.1 mi /race race * 1 km 5km 1mi Solution #2race * 1.61 km = 3.1 mi /race **Note: You want your old units to cancel diagonally so that you end with the new units. 100cm 1cm **Note: It does not matter which way you write the conve1mior 0.01m as long as they are on the correct side of the fraction line. ***Tip: It can be helpful when organizing your dimensional analysis to put the beginning and end units at the top of the page. It may not be necessary for an problem like the above example, however when the questions get more complex it is a good way to stay organized and headed in the right direction. 7 Question: A city uses 1.2 x 10 gallons of water /day. How many liters per hour needs to be pumped from the lake in order to supply the city? (1 gal = 3.785L) gallons/day → liters/hour 7 1.2*10 gal 1 day 3.785L Solution: day * 24 hours * 1 gal = 1,892,500L/hr **Note: Notice how days and gallons cancel diagonally in order to end with L/hr. **Note: Oftentimes a conversion factor will be needed (in this case gallons to liters), which will always be provided in a classroom setting. Question: A tank of gas holds 25.8 gallons. If the density of gasoline is 0.9315 g/mL, find the mass (in lbs) of the fuel in a full tank. (1 gal = 3.785L, 1g = 0.002205lb) Solution:25.8gal 3.785L 1000mL 0.9315g 0.002205lb = 200.6lb/tank tank * 1gal * 1 L * 1 mL * 1 g ______________________________________________ BUILDING BLOCKS OF CHEMISTRY The Atom: ● Smallest unit of an element ● Electrically neutral ● Nucleus is surrounded by electrons Subatomic Particles Particle Symbol Location Charge Mass* Proton p+ Inside nucleus 1+ “heavy” Neutral n0 Inside nucleus 0 “heavy” Electron e− Outside nucleus 1- ~0 *The exact masses of the three particles are not relevant, what matters is that the proton and the neutral weigh roughly the same, and the mass of the electron, while existent, is so miniscule by comparison that it can be disregarded. Atomic number (z): ● Equals the number of protons ● identifies the element ● if atom is neutral: # protons = # electrons Mass Number (A): # protons + # neutrons Rather than writing out an element’s full name, whether it is electrically neutral, as well as how many neutrons it has, scientists have come up with an easier way. The element symbol is large and in the middle, the charge is written in the upper right corner, the mass in the upper left corner, and the atomic number in the bottom left corner. The atomic number, though, is not necessary because it gives the same information as the element symbol. Both ways to write it are accepted. The diagram above is an example of how a cation of sodium-23 might be written. Atomic Mass on the Periodic Table: You might have noticed that the atomic mass on the periodic table is not a whole number, which may seem odd considering the mass of each individual atom will always be a whole number (you cannot have partial protons or neutrons). The mass on the periodic table is a reflection of the mass and the percent composition of all the different isotopes of each element. How to determine the atomic mass of an element: mass of isotope 1 * abundance of isotope 1 + mass of isotope 2 abundance *f isotope 2 *abundance is in decimal form (i.e. if abundance is 40%, it is represented in this formula by 0.4) Units of Atomic Mass: The mass on the periodic table has two different units. Depending on what you are trying to do with it, you can use either. - amu / atom - Grams / mol The molar mass of a substance can be found by adding the atomic mass of each atom in a compound together. Ex. Atomic mass of O i2 32g because one mole of oxygen weighs 16 grams, and since O has two oxygen atoms per molecule, then one mole of O will equal two moles of oxygen 2 2 atoms and therefore the 16 grams is multiplied by two. Blackbody radiation: When a substance is heated up, the excess energy will be emitted in the form of light. Composition of a Wave Classical Wave Theory (old): A wave’s energy depends on its intensity (amplitude) New Theory by Einstein (current): a wave’s energy depends on its frequency ● Einstein also theorized that a wave is not a single unit, but rather a stream of particles (photons) ● This led to the theory of wave-particle duality, which states that light is both wavelike and particle-like Quantum Mechanics now states that emitted light energy from atoms can only come in discrete (quantized, not continuous) values. Imagine many packets of light (called photons) moving within every wave. Because of this theory we can calculate the energy of a photon in a wave using the following: E = hν = energy in joules of one photon h = 6.626*10 −34 (Planck’s constant - no units) The two equations above can be manipulated to form: E = hc ν Relationships to Note: - As frequency increases, wavelength decreases (inversely related) - As energy increases, frequency increases and wavelength decreases - Amplitude has no relation and instead has to do with the brightness / intensity of the light Photoelectric Effect The diagram on the right shows that when photons in light waves with enough energy hit metal, electrons are ejected off the metal. The photons each need to hit with a minimum energy to emit an electron (varies depending on the element), but any excess energy the photon has beyond that gets translated to kinetic energy of the electron. The kinetic energy of the electrons can be harnessed for various uses. This is how we get solar energy from solar panels. Higher Intensity = more photons Higher Energy = short wavelength / high frequency The kinetic energy of an electron can be determined using the following equations: KE = E −Φ p KE = m v1 2 2 e Φ = Work Function/Binding Energy (energy required to emit an electron) E = the kinetic energy of the photon p m e mass of electron (9.11 1* −3kg) v = velocity (meters / second) KE = kinetic energy (Joules) ____________________________________________________ MODELS OF THE ATOM Bohr’s Model of the Atom: Bohr theorized that an atom does not radiate energy while in energy level (stationary state). Instead an atom absorbs or emits energy when an electron changes energy levels. Energy of a Photon = ΔE| atom|(energy = energy in other equations we know) | | N is called uantum number. Quantum numbers are: ● Integers (not zero or negative) ● Tell relative radius of orbit ● Tell energy level ● As n increases, energy levels get closer together Below is a diagram of the shape of different orbital. With each new energy level an electron reaches, the number of orbitals is double that of the previous energy level. ______________________________________________ ORBITAL ENERGY The more energy an electron has, the less stable it is. What affects an electron’s energy? 1) Charge - having a positive charge makes an ion more stable because it has more protons in the nucleus which pull the electrons in closer 2) Size of orbital - the smaller the orbital, the lower the energy because the electrons are pulled in closer to the nucleus and therefore more stable 1 s 2 s 3 s ... 1 p 2 p 3 <p... 3) Shape of Orbital - different orbitals have different levels of penetration due to their shape 4) Shielding - electrons in a higher energy level are shielded from the full force of the nuclear pull by electrons in lower energy levels ○ Z eff= effective nuclear charge (net positive charge experienced by the valence electrons → Z eff= Z - SS being the number of shielding elec)rons) ○ Electron feels less attraction = less stable = higher energy ______________________________________________ THE ELECTRON CONFIGURATION ● Represent distribution of electrons within energy levels and sublevels of atoms ● Quantum numbers → atomic orbitals → location of e → physical properties of atoms How to Write an Electron Configuration: 1) Apply Aufbau principle: shart at H, and for each addition to the configuration add one proton and one electron to the lowest energy sublevel not already filled (aufbau loosely translates to “top to bottom”) 2) Hund’s Rule: when orbitals of equal energy are available, the lowest energy configuration has the highest number of unpaired electrons with the same spin 3) Pauli Exclusion Principle: no two electrons can have the same four quantum numbers in the same atom Electron Configuration Notation: # nl Ex: 2p 2 n = energy level, l = sublevel (letter), # = number of e in sublevel Shorthand Configuration: instead of writing out every electron in every orbital, write the previous noble gas in parentheses, then record the remaining electrons Element Electron Configuration (Noble Gas) Shorthand Configuration Calcium (Ca) 1s 2s 2p 3s 3p 4s 2 [Ar] 4s 2 2 2 2 10 5 2 10 5 Iodine (I) 1s 2s 2p 3s ... 5n 4d 5p [Kr] 5n 4d 5p 2 2 6 2 2 14 4 2 14 4 Tungsten (W) 1s 2s 2p 3s ... 6n 4f 5d [Xe] 6n 4f 5d Orbital Diagram: - Pictorial representation of an electron configuration - One box per orbital (n, l, m )l - Grouped by sublevel (3d, 4s, 5f...) − - Arrow for e shows its spin Types of Electrons: ● Valence electrons - ○ Electron in the outermost energy (highest n value) ○ Include unfilled d or f orbital (even if not in highest n) ● Core Electrons - ○ All lower levels of n ○ Usually all accounted for in noble gas shorthand IONS Cations Anions Positive charge (+) Negative charge (-) Lose electrons Gain electrons The electron configurations of ions are written based on the number of electrons they have. They are identical to the configurations of atoms with the same number of electrons (called being isoelectronic). 2 2+ + Ex. The configuration 1s is the configuration for helium, Be , Li and other ions. Transition Metal Ions: - When writing the electron configuration for these, keep in mind that ns is filled after (n-1)d because ns it is farther from the nucleus 2+ - Almost all transition metals form X ions because they lose their outermost s electrons Fe: Fe :+ Fe :+ 2 6 5 6 [Ar] 4n 3d [Ar] 3d [Ar] 3d MAGNETIC PROPERTIES Paramagnetic Diamagnetic Species with at least one unpaired electron All electrons are paired Attracted to a magnetic field Slightly repelled by a magnetic field Examples: 2 2 Ti →[Ar] 4s 3d There are two unpaired electrons in the 3d suborbital so Ti is paramagnetic. Ti →[Ar] 4s All electrons in the 4s orbital are paired (and all other orbitals) so Ti2+ is diamagnetic. ______________________________________________ PERIODIC TRENDS The periodic table is organized by increasing energy, so the most recent electron added has the highest energy. 1) ATOMIC SIZE: - Determine the radius of an atom by measuring nucleus to nucleus between two bonded identical atoms, then divide that measurement by two - Down a group: the atomic radius increases because there are more inner electrons shielding the outer electrons, so they are not pulled in as tightly - Across a period: the atomic radius decreases because all outer electrons are in the same energy level, however there is a higher nuclear charge to pull them in tighter - Shielding has no effect when electrons are added to the same energy level - Zeff increases across a period 2) IONIZATION ENERGY: - Ionization energy is the amount of energy required to remove the outermost electron from an atom in a gaseous state - Always p+sitive ( >0 ) + 2+ - X g X +eg= IE− 1 X g X +eg= IE − 2 - Down a group: Ionization energy decreases because shielding is greater (and therefore distance from nucleus) so it takes less energy to remove an outer electron - Across a Period: Ionization energy increases because the Z increases (from the eff additional protons, but no extra shielding because all new electrons are in the same energy level) and therefore it takes more energy to remove an electron 3) ELECTRON AFFINITY: - An atom’s electron affinity is the theoretical change in energy if that atom were to gain an electron while in a gaseous state - Electron affinity can be positive or negative (but is usually negative). This is related to how many valence electrons an atom has and whether adding or subtracting an electron would get it closer to a full valence shell. - The more negative the electron affinity is, the more an electron is wanted. A negative affinity means energy is released when an electron is gained. To keep things as least complicated as possible, electron affinity will be referred to in terms of its magnitude. - Down a group: electron affinity decreases (becomes less negative) because the electron being added is farther from the nucleus and therefore less energy is released - Across a period: electron affinity increases (becomes more negative) because the nuclear charge increases / the atomic radius decreases and therefore more energy is released 4) ION SIZE - Cations are smaller than their original (“parent”) atom because there are the same number of protons pulling on fewer electrons and therefore each electron experiences a stronger pull - Less electrons means less repulsion between them, resulting in outer electrons moving closer to nucleus - Anions are larger than their original (“parent”) atom because there are the same number of protons pulling on more electrons and therefore each electron experiences a weaker pull - More electrons means more repulsion between them, resulting in outer electrons moving farther from the nucleus - Down a group: the ion size increases because there is more shielding from the inner electrons and therefore the outermost electrons are farther from the nucleus - Across a period: the ion size decreases, then increases, then decreases again. This has to do with whether the ion has gained or lost electrons, which results in ions in the some period to have differences in their highest energy level and therefore the ion sizes vary. Since there is no clear trend, the relative size of an ion can be found by comparing the ratio of protons to electrons in isoelectronic species. + 2+ 3+ Na Mg Al # of protons 11 12 13 # of electrons 10 10 10 Relative Size largest middle smallest Exact Size (do NOT 102 pm 72 pm 54 pm need to know how to find this) *Note: this table is an example of the first decrease across a period − 2− − P S Cl # of protons 15 16 17 # of electrons 18 18 18 Relative Size largest middle smallest Exact Size (do NOT 212 pm 184 pm 181 pm need to know how to find this) *Note: this table is an example of the second decrease across a period **Note: A good tip when trying to understand periodic trends is to relate everything back to the size of the radius. Nearly all trends can be explained when thought about in terms of whether an atom has a bigger or smaller atomic radius. ***Note: All of what is written in these notes is important, otherwise it wouldn’t be here. What is highlighted in yellow, however, are the things I feel are the most important to know. It is also important to note that it is not important to memorize the highlighted equations because they will likely be provided for you, but rather it is important to understand the theories behind them and why they exist.
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