Test 2 Study Guide
Test 2 Study Guide CH 101
Popular in General Chemistry
Popular in Department
This 16 page Study Guide was uploaded by Rebecca de la O on Sunday October 2, 2016. The Study Guide belongs to CH 101 at University of Alabama - Tuscaloosa taught by Jared Allred in Fall 2016. Since its upload, it has received 96 views.
Reviews for Test 2 Study Guide
worth it , keeps you studying
Report this Material
What is Karma?
Karma is the currency of StudySoup.
You can buy or earn more Karma at anytime and redeem it for class notes, study guides, flashcards, and more!
Date Created: 10/02/16
TEST 2 STUDY GUIDE Chapter 4 Pauli Exclusion Principle - “No two electrons in an atom can have the same four quantum numbers” - So if two electrons are in the same orbital (same n, l and m qualtum number) then each will have a different spin quantum number (+1/2 or -1/2) Shielding/Penetration - Shielding is when an electron blocks another from fully experiencing the effect of the nuclear charge o If an element has a nuclear charge of 3+ and two electrons in the electron cloud, a nearby electron will only feel a 1+ charge from the nucleus because the inner electrons block the rest of the charge Effective nuclear charge - Due to shielding, the electron described above would only feel an effective nuclear charge of +1 3+ (from the nucleus) + 2- (from the two electrons) = 1+ - Penetration is when an outer electron enters the region occupied by the inner electrons o The electron described above will feel the full 3+ charge of the nucleus if it enters the part of the electron cloud where the other electrons are because they are no longer shielding it Aufbau Principle - The electron configuration of an atom is built by systematically adding electrons to the lowest energy sublevel and filling orbitals - Electrons fill orbitals of lower energy first (s then p then d then f) Hund’s rule - “When filling degenerate orbitals fill them singly first, with parallel spins” - Analogy: students on a school bus must fill every seat (every orbital in that n number) before doubling up Electron configurations - Electrons occupy orbitals to minimize the energy of the atom - Lower energy orbitals fill before higher energy orbitals o Order: 1s 2s 2p 3p 4s 3d 4p 5s 4d 5p and so on - Orbitals can hold no more than 2 electrons each (Pauli exclusion principle) o When two electrons occupy the orbital, their arrows point in opposite directions Atoms and ions o For regular atoms: o When you write the electron configuration of a positive ion, remove one electron for every positive charge o When you write the electron configuration of a negative ion, add one electron for every negative charge Don’t need to derive d electron cases o You do need to be able to use a given d-block e. configuration to get the ionic configuration. Valence electrons, determining the # and configuration. - Valence electrons are important in in chemical bonding because they are the easiest to lose or gain - For main group elements, the valence electrons are those in the outer most principal energy level (s and p) - For transition elements, valence electrons are those in d or f (if not filled) and s and p - Elements in the same column/group have the same number of valence electrons o Which is why they have similar chemical properties - Core electrons are those in complete, filled principal energy levels - Ex. How many valence electrons doe Silicon have? 2 2 6 2 2 Si 1s 2s 2p 3s 3p 4 valence electrons in the n= 3 principle energy level Ex. How many valence electrons does Germanium have? Ge 1s 2s 2p 3s 3p 4s 3d 4p 4 valence electrons (3d level is filled) Arrangement of the periodic table Groups – the columns; can also be called families; run from top to bottom Rows- run across the table; can also be called periods Blocks- the table is divided into four blocks corresponding to the filling of the four quantum sublevels (s, p, d, f) o Main group elements make up the s and p block; transition elements make up the d block; inner transition elements make up the f block o The number of columns in a block indicates the number of electrons that can fill that sublevel (s = 2; p= 6, d=10, f=14) o The group/column number of an element indicates how many valence electrons it has o The period/row number of an element indicates the highest n number Types of elements: metals, metalloids, and non-metals - Metals lie on the lower left side and middle of the table o Good conductors of heat and electricity o Lose electrons easily (want to achieve noble gas configurations) o Elements with 1 or 2 valence electrons are the most reactive metals - Nonmetals lie on the upper right side of the table o Varied properties o Bad conductors of heat and electricity o Gain electrons easily to achieve noble gas configurations o Elements with 6 or 7 valence electrons are the most reactive nonmetals - Metalloids lie on the zig zag that divide metals and nonmetals o Exhibit metal and nonmetal properties o Semiconductors of electricity Group relationships - Noble gases (Group 18 or 8A) – mostly unreactive due to filled orbitals - Alkali metals (Group 1 or 1A) – have outer electron configuration of ns ; 1 have electron configurations that are 1 electron more than that of a noble gas; easily and violently lose that one electron during reactions - Alkaline Earth Metals (Group 2 or 2A) – have an outer electron 2 configuration of ns ; two electrons more than that of a noble gas; easily react to lose those two electrons - Halogens (Group 17 or 7A) – have an outer electron configuration of ns 2 5 np ; are one electron away from being a noble gas configuration; react easily to gain that electron Ion formation - Atoms gain or lose electrons to form ions in order to achieve stable electron configurations like those of noble gases - Metals tend to form positive ions (cations) by losing electrons Nonmetals tend to form negative ions (anions) by gaining electrons - Alkali metals tend to form cations with a 1+ charge Alkali earth metals tend to form cations with a 2+ charge Halogens tend to form anions with a 1- charge Be aware of the elements that form only one type of element (shown below), and what those typical charges are. It’s better if you can derive these instead of memorizing. Periodic trends Size of atoms o Determined by its atomic radius- the radius of an atom when it is bonded to another atom o Increase to the left across rows and increase down the columns o Peaks at the alkali metals and reaches a minimum at each noble gas o Reason: atomic radius is determined by its valence electrons, so elements lower on the periodic table have valence electrons further from the nucleus, increasing their size. Why atomic size decreases to the right of the table is determined by the inward pull of the nucleus; explained by effective nuclear charge. o Transition metals are the exception; their size stays constant because the number of electrons in their highest orbital (d) remains constant because lower orbitals (s) are changed first Size of ions o Radius of a cation is much smaller than its neutral counterpart An electron is lost, so electron-electron repulsion and shielding decreases and the protons are able to pull the remaining electrons towards the nucleus (thereby increasing Zeff) o Radius of an anion is much larger than its neutral counterpart Since additional electron occupies an outer orbital, electron-electron repulsion and shielding increases which pushes the electrons further apart. Electrons now outnumber the protons in the ion, so the protons can’t pull the extra electrons as tightly toward the nucleus; this results in decreased Zeff Effective nuclear charge o “the net charge experienced by an electron” o Zeff = Z (actual nuclear charge) – S (charge experience after shielding) o Affected by shielding o According to Coulombs law, the attraction between a nucleus and an electron increases with an increase magnitude of nuclear charge o Effective nuclear charge increases to the right of the table and up the columns, resulting in a smaller atomic radius because the attraction between the nucleus and the outer electrons is stronger Ionization energies o “Energy required to remove an electron from an atom in its gaseous state” Removing an electron from an atom requires enough energy to overcome the magnetic pull of the positive charge of the nucleus. Always positive o Increases up a column and increases right across the rows IE increases up a column because there are less electrons shielding the outer electrons from the pull of the nucleus. Therefore, it requires more energy to out power the nucleus and remove an electron IE increases to the right because the larger the effective nuclear charge means the stronger the nucleus is holding onto the electron and the more energy it takes to release an electron o IE increase relative to effective nuclear charge Highest ionization energies are the noble gases because they all have high effective nuclear charges due to their filled orbitals and require a high amount of energy to destroy that stable configuration Elements in the left corner have a low ionization energy because losing an electron allows them to have the noble gas configuration. Therefore, it requires less energy to remove one of their valence electrons 1st ionization energy trends o 1 ionization energy is the energy required to remove one valence electron nd 2 ionization energy is the energy required to remove a second valence electron 3 ionization energy is the energy required to move a third valence electron Electron affinity. What does it mean? - The energy associated with an element gaining an electron in its gaseous state; the atom's likelihood of gaining an electron - Exothermic- The electron affinity is positive Endothermic – electron affinity is negative - Increases (more negative= more exothermic) to the right of the table o The greater the negative value is, the more stable the anion is - No trend down columns Chapter 4 Practice Problems 1. Write electron configurations for the following: a. P 2- b. S c. Zn 3+ 2. Arrange these elements according to decreasing atomic size: Na, C, Sr, Cu, Fr 3. Why are cations smaller and anions larger than their respective atoms? (More than 1 may be correct) a. Cations are larger than their respective atoms because of increased electron-electron repulsion b. Anions are smaller than their respective atoms because of increased effective nuclear charge c. Cations are larger than their respective atoms because of decreased electron-electron repulsion d. Anions are smaller than their respective atoms because of decreased effective nuclear charge - - 2+ 4. Which of the following are isoelectronic: F , Cl , Ca , Ar 5. Arrange these elements according to increasing metallic character: Li, S, Ag, Cs, Ge 6. List the ions from smallest to largest: Se , Zr , Na , Mg, Rb , Br , K + - + 7. Arrange these elements according to increasing negative electron affinity: Ba, F, Si, Ca, O Chapter 4 Practice Problem Answers 2 2 6 2 3 1. a. 1s 2s 2p 3s 3p 2 2 6 2 6 b. 1s 2s 2p 3s 3p c. 1s 2s 2p 3s 3p 4s 3d 2 7 2. Fr, Sr, Cu, Na, C 3. C & D: Cations are formed when an electron is lost. When this occurs there are less electron-electron repulsions and there is a greater net nuclear attraction per electron. So, the newly formed ion becomes a more condensed version of its neutral atom. Anions are formed when an electron is gained. When this occurs there are more electron- electron repulsions and there is a lower net nuclear attraction per electron, causing them to push each other away and spread out, so the atom to becomes larger. 4. Cl-, Ca2+, Ar all have 18 electrons; therefore, they are isoelectronic (F- has 10 electrons) 5. Li, S, Ge, Ag, Cs 4+ + + - 2- 6. Zr <K <Rb <Mg<Br <Se : Ionic radii shorten with increasing positive charge and lengthen with increasing negative charge, and thus, anions are larger than cations 7. Ba, Ca, Si, O, F Chapter 5 Ionic vs. covalent bonds - “The bond that forms between a metal and a nonmetal is an ionic bond” o Because metals like to lose electrons, the metal will transfer one or more electrons to the nonmetal (which tends to gain electrons). o The metal becomes a cation; the nonmetal becomes an anion o The oppositely charged atoms now attract each other (according to Coulombs law) and form an ionic compound o Arrangement: lattice structure (alternating cations and anions) - “The bond … between two or more nonmetals is a covalent bond” o Nonmetals have high IE so neither atom transfers electrons; instead they share o Shared electrons interact with the nuclei of both atoms, lowering their potential energy (according to Coulombs law) o Covalently bonded atoms form molecules molecular compounds o Arrangement: set number of particles interacting and sharing electrons in a unique arrangement - Cute picture to help you understand ionic and covalent bonds: One atom (the The atoms are cation) is giving up sharing their an electron to electrons so both another (the anion) can have stable configurations Empirical vs. molecular formulas - “An empirical formula gives the relative number of atoms of each element in a compound” - “A molecular formula gives the actual number of atoms of each element in a molecule of a compound” Naming compounds - Binary ionic compounds Contain only two different elements o Metal forms 1 kind of ion Name = name of metal (cation) + [base name of nonmetal (anion) + “ide”] Ex. CaO calcium oxide Calcium (cation) followed by the base of oxygen (anion) which is ox, then the ending -ide Ex. CaBr 2 calcium bromide Calcium (cation) followed by the base of bromine (anion) which is brom, then the ending -ide o Metal forms more than 1 kind of ion Same formula as above (name of metal (cation) + [base name of nonmetal (anion) + “ide”]) but the name of the metal is followed a roman numeral Roman number indicates the charge of the metal 2+ Ex. Fe3+ Iron (II) Fe Iron (III) To determine the charge and roman numeral of the metal, remember the sum of all the charges (cation and anion) must be zero (neutral) Ex. CrBr3 the charge of chromium must be 3+ to cancel out the 3 Br anions that have a charge of 1- You know Br has a charge of 1- because it wants to gain 1 electron to have a stable configuration like a noble gas 3+ Cr chromium (III) CrBr3--> chromium (III) bromide o Polyatomic ions Polyatomic ions are “ions composed of two or more atoms” If the ion is in its most stable form, name is base of element name + “ate” If the ion loses 1 oxygen, name is base of element name + “ite” If the ion loses 2 oxygen, name is (prefix “hypo”) + (base) + (suffix “ite”) If the ion gains 1 oxygen, name is (prefix “per”) + (base) + (suffix “ate”) Ionic compounds with polyatomic ions are named the same way as other ionic compounds but use the name of the polyatomic ion Ex. NaNO sodium nitrite 2 + Contains the cation sodium Na and the polyatomic - anion NO n2trite Ex. FeSO 4 iron (II) sulfate Contains the cation iron with a 2+ charge and the 2- anion SO 4 Ex. NH 4O 3ammonium nitrate Contains polyatomic cation ammonium and polyatomic anion nitrate See Table 5.4. You should know these names, except for acetate and bicarbonate and its analogs. Use the rules you have learned about ‘ate’, ‘ite’, etc. to generate other names not listed in the table. Even if you don’t see them on the exam, these will come back over and over again, so if you don’t learn them now, you will have problems later o Hydrates Contain a specific amount of water molecules Name them like you would other ionic compounds but add “prefix-hydrate” where the prefix is the number of water molecules - Binary molecular compounds Can’t always determine the formula for a molecular compound from it’s the elements because the elements may form several different molecular compounds Composed of two or more nonmetals The first element is the more metal-like one (toward the left and bottom) Name = (prefix + name of first element) + (prefix + base name of second element + “ide”) If there is only 1 atom of the first element, we omit the prefix mono Ex. NO 2 nitrogen dioxide Ex. N2O dinitrogen monoxide Ex. P4S10 tetraphosphorous decasulfide Molar mass of a compound o An elements molar mass is the mass in grams of one mole of its atoms o Molar mass is numerically equivalent to the formula mass or molecular mass o (Number of atoms of 1 element x atomic mass of 1 element) + (Number of atoms of 2 nd element x atomic mass of 2 ndelement) o Ex. C H6O12 6 (6 x 12.01) + (12 x 1.01) + (6 x 16.00) = 180.18 amu - Calculate mass percent of atom or polyatomic ion, within a compound. o The mass percent of an element is that element’s percentage of the compounds total mass massof thatelement∈1mol of thecompund o Mass percent of element = massof 1mol of thecompound x 100% o Ex. Calculate the mass percent of Cl in CCl F 2 2 2 x molarmass of Cl Mass percent of Cl = molar mass of CCl2 F 2 100% Molar mass of CCl 2 2 (12.01) + (2 x 35.45) + (2 x 19.00) = 120.91 g/mol 2 x 35.45g/mol Mass percent of Cl = 120.91g/mol x 100% Mass percent of Cl = 58.64% - Use masses of constituent atoms to find the empirical formula Ex. Calculate the empirical formula of a compound with 24.5 g of nitrogen and 70.0 g of oxygen 1) Get the given mass of each element present in the compound. This may be given directly, or you may have to calculate it from mass percentage 24.5 g of Nitrogen 70.0 g of Oxygen 2) Convert each mass into moles by using the appropriate molar mass of each element as a conversion factor 1mol N 24.5 g N x = 1.75 mol N 14.01gN 1 molO 70.0 g N x 16.00 gO = 4.38 mol O 3) Write a pseudo formula using the number of moles N1.75 4.38 4) Divide the subscripts by the smallest subscript to get the smallest whole number N 1.75 4.38 1.75 O 1.75 NO 2.5 5) If the subscripts are not all whole numbers, multiply by the smallest whole number possible in order to get all the subscripts to be whole numbers N 1 x 2 2.5 x 2N 2 5 The empirical formula is N O2 5 Note, this may not be the molecular formula. To get the molecular formula, we would need the molar mass of the entire compound.
Are you sure you want to buy this material for
You're already Subscribed!
Looks like you've already subscribed to StudySoup, you won't need to purchase another subscription to get this material. To access this material simply click 'View Full Document'