Chemistry Midterm Review 2024 PDF
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St. Mary's School
2024
Emma Gibbons
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This document is a chemistry midterm review for the Fall 2024 semester, focusing on topics such as Dalton's atomic theory, the discovery of the electron, the oil drop experiment, and Rutherford's gold foil experiment. The document also contains short answer questions and examples of chemical calculations.
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C hemistry Midterm Review FALL SEMESTER 2024 E mma Gibbons [email protected] Atomic Structure of Matter Dalton’s Atomic Theory: Dalton’s Atomic Theory is made up of 5 main postulates: 1. Each element is composed of tiny indestructible particles called atoms...
C hemistry Midterm Review FALL SEMESTER 2024 E mma Gibbons [email protected] Atomic Structure of Matter Dalton’s Atomic Theory: Dalton’s Atomic Theory is made up of 5 main postulates: 1. Each element is composed of tiny indestructible particles called atoms 2. Atoms of a given element are identical. Atoms of different elements are different 3. Compounds are formed when atoms of different elements are combined. A given compound always has the same relative numbers and types of atoms 4. Chemical reactions involve the reorganization of atoms, but the atoms themselves are not changed during the process 5. Atoms cannot be divided any further Incorrect Law of Multiple Proportions-When 2 elements combinewith each other to form more than one compound, the weights of one element that combine with a fixed weight of the other are in a ratio of small whole numbers 1.116𝑔𝐶 𝑙 2 0.558𝑔𝐶 𝑙 EX: 1.00𝑔𝐶 𝑙 = 1 𝑜 𝑓 1𝑔𝐶 𝑙 1.116 Cl is double the amount of 0.558g Cl Law of Definite Proportions-A given chemical compoundalways contains the same elements in the exact same proportions by mass 5.33𝑔𝑐 𝑎𝑟𝑏𝑜𝑛 EX: 1.00𝑔ℎ 𝑦𝑑𝑟𝑜𝑔𝑒𝑛 2 atoms have the same mass ratio 1 Discovery of the Electron and the Nucleus J. J. Thomson-discovered the electron using a cathoderay tube experiment and made the “plum pudding” model of the atom Cathode ray tube experiment: 1. Cathode would send electrons to the Anode 2. This would cast an electron beam 3. The direction of the beam was based on how the magnets were positioned and their charge Thomsons observations: Beam travels from the cathode to the anode: SHADOW Beam is made of particles (mass): PADDLE SPINNING Particles has a negative charge: DIRECTION OF MAGNET BEAM Plum Pudding Model: Showed that the Electrons had a negative charge 2 Robert Millikan-discovered how to determine the mass and charge of an electron with the help of his oil drop experiment Oil Drop Experiment: How it works: 1. The oil drops picked up static charge and were suspended between two charged plates 2. Millikan was able to observe the motion of the oil drops with a microscope 3. Found that the drops lined up in a specific way between the plates, based on the number of electric charges they had acquired. Lead to discovering charges and the mass of an electron along with how to find an electron’s fundamental charge Ernest Rutherford-discovered the proton using hisgold foil experiment. He also created the “solar system” model of the atom Gold Foil Experiment: 3 Rutherford’s conclusions: 1. Most alpha particles passed through the gold foil which suggested that atoms are mostly empty space 2. Some alpha particles were deflected slightly, indicating that there is a concentrated positive charge within the atom 3. A few alpha particles were deflected back toward the source, providing evidence that the concentrated positive charge is in a small, densely packed region which Rutherford termed as the nucleus Solar System Model: He knew this model was wrong: ○ Electrons can’t orbit around a positive nucleus without caving in Scientific Method-the process of objectively establishing facts through testing and experimentation 4 Right-hand rule: 1. Make an L-shape with your thumb and index finger and middle finger. 2. Point your thumb perpendicular to your thumb and index finger 3. Point your index finger in the direction of the positive charge 4. Your middle finger=the direction of the magnetic field 5. Thumb=direction of the moving charge; a. Positive, charge is where the thumb point b. Negative, charge is opposite the thumb point Short Answer Questions: Elemental Symbol Nuclide Symbol Isotope=variant of an element/different number ofneutrons Ions=an atom that has a negative of positive electricalcharge Rules: 1. Atomic number=number of protons in the nucleus 2. Protons+neutrons=mass number 3. Number after the element=its mass; oxygen-12 mass=12 4. Mass number-Atomic number=number of neutrons 5. Charge is based on the relationship between protons and electrons Isotope and Ion Table EX (things to find are in blue): 5 Isotope name Symbol # of p+ # of e- #of n0 mass number 4 helium-4 𝐻𝑒 2 2 2 2 4 40 calcium-40 20 𝐶𝑎 20 20 20 40 184 tungsten-184 74 𝑊 74 74 110 184 210 lead-210 82 𝑃𝑏 82 82 128 210 237 neptunium-237 93 𝑁𝑝 93 93 144 237 70 𝐺𝑎 31 31 39 70 gallium-70 31 95 molybdenum-95 𝑀𝑜 42 42 42 53 95 Ion name Symbol # of p+ # of e- charge potassium K+ 19 18 +1 calcium Ca2+ 20 18 +2 aluminum Al3+ 13 10 +3 oxide O2 - 8 10 -2 fluoride F– 9 10 -1 copper (II) Cu2+ 29 27 +2 Relative and Average Atomic Mass: Relative Atomic Mass=Mass given to you on the periodicrounded Average Atomic Mass=The Mass found between differentisotopes of an element Number closest to average atomic mass=most abundant ○ Hydrogen-1 is most abundant because average atomic mass is 1.008 6 Round to the 3 decimal place ○ EX:23. 985(0. 7870) 24. 986(0. 1013) 25. 983(0. 1117) ←Percent . 9023011 = 24. 310 ←AverageAtomic Mass ○ 18. 876195 + 2. 5310818 + 2 Electronic Structure of the Atom Light and Spectroscopy Light waves: Each color relates to a different energy level ○ Red=Low ○ Purple=high ○ This is also infrared→UV; with UV being the highest and not visible and infrared being the lowest. The energy is also related to the wavelengths ○ The farther apart the wavelengths are the lower the amount of energy ○ The closer together the wavelengths are the higher the amount of energy 7 To sum it up close together waves=high energy=colors like purple or UV light; far apart waves=low energy=colors like red or infrared light Spectroscopy: It shows the energy of e-’s is quantized The longer the arrow=the more energy there is If the arrows are going down the energy is emitting; If the arrows are going up the energy is absorbing Dark light=Absorption; Light light=Exmission The higher the energy levels are the closer the orbits are together; So the further up you get the closer the lines are together Particle Nature of Light Photoelectric effect and Double Slit Experiment: Double Slit experiment-A wave of light was shined at a board that had two slits in it; however it came together and formed one line not 2 like expected ○ When it is shined the light goes to dark and light lines across the board=This is the interference pattern ○ Shows how it is wave and not particles because it formed one line due to the interference pattern of the wave 8 ○ Photoelectric Effect-Is when light shines on metal; electrons are ejected from said metal ○ This also shows how light has both wave-like and particle like properties ○ Line spectrums: There are 3 different line spectrums that show different things; continuous, emitting; absorption ○ Continuous line spectrum-Nothing breaks up the spectrum; it shows all the wavelengths of light in a certain range ○ Emission spectrum-Shows the light that is emitted off of it; The light that is left on the spectrum is the light that is emitted and in between the lines is dark space of the light that isn’t emitted ○ Absorption spectrum-Shows the light that is absorb; The light that is absorbed is left in the visible spectrum and so the light that is emitted is a dark space on the spectrum 9 ○ The Bohr Model and the deBroglie Hypothesis The Quantum Model of an Atom was discovered by Neils Bohr It views electrons as waves not particles deBroglie Hypothesis: It stated that electrons can be described with both waves and particles ○ This is important because it helps prove the Line Experiment ○ It also proves other quantum models and experiment The Quantum Mechanical Model Quantum Mechanical Model: Views the electrons inside of an atom as waves This then causes electrons to reside in a specific orbital or region of space within the atom This is also how electrons are in certain “shells” in the atom 10 Schrodinger-In his model electrons don’t travel in sharp, distinct orbits but rather orbitals or regions of space in the atom Heisenberg-Came up with an Uncertainty Principle: There is unpredictability in measuring variables in a particle ○ It limits the amount you know for example you can’t accurately measure the speed and position Short Answer Questions: Electron Configurations: To write an Electron Configuration you can use subshell notation or orbital notation Subshell notation: Amount of electrons each sublevel holds ○ S-2 D-10 P-6 F-14 When reading the rows for configurations you start 1s→2s→2p ○ For s and p you just read the row number ○ For row d you start at 3 ○ For row f it starts at 4 Certain levels can only hold a certain number of electrons ○ s=2 ○ p=6 ○ d=10 ○ f=14 Subshell notation is when you use the subshells to name the electron ○ Ex: Si-1s22s22p63s23p2 ○ Noble Gas configuration-you use the noble gas symbol that is right before the element you are naming and then do the notation as normal Ex:Si-[Ne] 3s23p2 Orbital notation: 11 Using orbital notation is when all the rules and exceptions mainly come in play Aufbau Principle-You starting filling the orbitals that have the lowest energy ○ Low energy→high energy Hund’s Rule-You must fill a complete orbital before moving on to fill the others ○ You can’t leave an orbital empty Pauli Exclusion-The orbitals can not be parallel and must spin different ways Aufbau Exception-This is where you add or subtract electrons to the electron configuration. You would add or subtract to either get the full amount of electrons or half of it ○ Ex:4s24 p5→ 4s14p6 ○ The exception would just impact the elements charge: Taking away electrons=Positive charge Gaining electrons=Negative charge ○ The exception is used because it makes the electrons more stable whether they are fully filled or half filled Orbital notation-Is based on the electrons orbits but is basically the same as electron configuration ○ You find the electron configurations but use lines and arrows to fill it 12 ○ Each subshell has a curtain number of lines which is half the amount of electrons it can hold S-1 P-3 D-5 F-7 The Periodic Table Historical Overview Dimitri Mendeleev and Lothar Meyer riginal Tables ordered by atomic mass O Dimitri Mendeleev and Lothar Meyer independently produced tables in 1869. Both arranged elements into groups with similar chemical and physical properties. Mendeleev predicted the existence of elements which were yet to be discovered. He left gaps in his table for these elements. John Newlands: Often overlooked, although his first table was published in 1864, before Mendeleev or Meyer. Newlands proposed a “periodic law” describing the repeating pattern of elements, although his work was not widely recognized during his lifetime. Predicted unknown elements, like germanium, that were later discovered.room for undiscovered elements As Mendeleev and Meyer would also do, he grouped elements with similar properties together. Law of Octaves – chemical properties repeat every 8 elements. Groups Periods and Electron Configurations Groups and Periods: roups or families of elements are columns ↕ G Periods are rows within the table ↔ 13 Electron Configurations Find in Electronic Structure of the Atom Valence Electrons Valence electron=Electron located on the outermost electron shell of an atom. Inner electron=Electron’s on the innermost shells Short Answer Questions: Periodic Properties and Trends Shielding=The inner electron shell repelling the outerelectron shell. When there is more shielding the pull on the electrons DECREASES When there is less shielding the pull on the electrons INCREASES 14 Nuclear Charge=The number of protons in an atom When there is more nuclear charge the pull on the electrons INCREASES When there is less nuclear charge the pull on the electrons DECREASES Electron; Electron Repulsion=The more electrons thereare the more electron repulsion there is Electronegativity, ionization energy, atomic/ionic radii Electronegativity: Electronegativity=a measure of an atom's ability to attract shared electrons to itself SHIELDING The more shielding there is the lower electronegativity The less shielding there is the higher the electronegativity This is due to the fact that when there is less shielding the pull on electrons are higher; the more shielding the pull on the electrons are lower which decreases the electronegativity NUCLEAR CHARGE The more nuclear charge there is the higher the electronegativity The less nuclear charge there is the lower the electronegativity This is due to the fact that the more nuclear charge there is the higher the pull on the electrons; the less nuclear charge the less pull on the electrons More pull on the electrons=higher the electronegativity Noble Gas is an exception though because the electronegativity is 0 15 Ionization Energy: Ionization energy=energy needed to remove an electron from an atom or ion Shielding The more shielding there is the lower the ionization energy The less shielding there is the higher the ionization energy Nuclear Charge The more nuclear charge there is the higher the ionization energy The less nuclear charge the lower the ionization energy Ionic radii: SHIELDING The higher the shielding means the ionic radius gets bigger The less shielding means the ionic radius gets smaller NUCLEAR CHARGE The higher the nuclear charge means the ionic radius gets smaller The less nuclear charge means the ionic radius gets bigger 16 Atomic Radii SHIELDING As the shielding increase the atomic radius gets bigger As the shielding decreases the atomic radius gets smaller NUCLEAR CHARGE As the nuclear charge increase the atomic radius gets smaller As the nuclear charge decreases the atomic radius gets bigger Overall: Ionic radius=Atomic Radiuc 17 Chemical Bonding and Formulas Nature of the Chemical Bonds Chemical Bonds are formed when electrons are shared between 2 or more atoms There are 3 main types of bonds: Covalent; Metallic; Ionic Covalent=nonmetal+nonmetal Ionic=metal+nonmetal Metallic Bonding=Metal+metal Types of Bonding Metallic Bonding: Very strong bonds which give metals high melting and boiling points Metal form a sea of electrons ○ Delocalized (not connected to an atom or bond) valence electrons Metal Characteristics: ○ Conductivity-conducting heat and electrons ○ Malleability-they can flatten ○ Ductility-it can stretch instead of break ○ Low volatility-high boiling point and high melting point Protons stay in place and the electrons move around Shell model comes in use because of the irregularities of melting points Ionic Bonding: Formed mostly with ions 18 Metals tend to give electrons Nonmetals tend to receive electrons Covalent Bonding: Based on Electronegativity Equal # of electrons→Nonpolar covalent bond Unequal # of electrons→Polar covalent bond Short Answer Questions: Lewis Structures: 1. Find the total number of valence electrons a. EX: H2O 1(2)+6=8 b. CN- 4+5+1 (charge)=10 2. Put the least electronegative atom in the center a. NO2 Nitrogen would be in the center b. PCl3 P would be in the center c. Hydrogen always goes on the outside 3. Put two electrons between atoms to form a chemical bond a. H:Cl 4. Completer octets on the outside atoms a. H-1 Cl-7 8-2 Cl gets the rest b. Hydrogen only needs 2 electrons 5. If central atom does not have an octet, move electrons from outer atoms to form double or triple bonds 19 OR 1. ve=# of valence electrons times # of atoms 2. pe=ve/2 3. pn=# of bonds needed times # of atoms 4. bp=pn-pe 5. nbp=pe-bp The number of bonds an element wants to form is equal to the number of additional electrons it need for an octet EX:SnCl2 ve=4+7(2)=18 pe=18/2=9 pn=4(1)+4(2)=12 bp=12-9=3 nbp=9-3=6 Diatomic atoms=hydrogen, nitrogen, fluorine, oxygen, iodine, chlorine, bromine Covalent Compounds and Molecular Geometry Short Answer Questions: 20 VSEPR: Valence, shell, electron, pair, repulsion ○ Predicts shape of a molecule Atoms are surrounded by a cloud of negatively charged electrons ○ Repeal each other ○ Forces atoms as far away as possible from each other Closer together=more energy; You want the LEAST amount of energy Each electron domain=1 bond or 1 lone pair 2 bonds=2 electron domains and so on ○ Sigma bond=single bond ○ Pi bond=double or triple bond # of electron domains/# of bonds→hybridization of central atom→electron domain geography ○ Domain Geometry # of electron domains Hybridization Linear 2 sp Trigonal planar 3 sp2 Bent (1 lone pair) 3 sp2 Tetrahedral 4 sp3 rigonal pyramidal (1 T 4 sp3 lone pair) Bent (2 lone pairs) 4 sp3 Trigonal Bipyramidal 5 sp3d Octahedral 6 sp3d2 Hybridization (and Expanded Octets) Expanded Octets: 21 –More common in the d, f, g periods –Have more than 8 valence electrons –Allows 5 of more bonds Formal Charge and Resonance: Formal Charge: -figure out the best structure for a chemical compound -most likely to exist in nature Equations: Formal Charge=Valence electrons-NonBonding electrons-Bonding electrons/2 -Do this for each atom -Choose the structure with the charge closest to 0 -Ignore the signs EX: Resonance Structures 22 –different ways to draw lewis structures –From what I have seen you just move around the double bond 1. Find # of valence electrons 2. Find # of lone pairs a. LP=ve-8n divided by 2 i. N is the # of atoms that doesn’t include the central atom ii. Only works if Hydrogen is NOT involved 23
