Chapter 4 Notes
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This 9 page Class Notes was uploaded by Rebecca de la O on Friday September 23, 2016. The Class Notes belongs to CH 101 at University of Alabama - Tuscaloosa taught by Jared Allred in Fall 2016. Since its upload, it has received 10 views.
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Date Created: 09/23/16
Chapter 4: Periodic Properties of the Elements Periodic Trends - Mendeleev was a chemist who proposed periodic law and created the first periodic table o The Periodic Law: elements with similar properties recur in a regular pattern when arranged in order of increasing mass o Organized known elements in rows (left to right) by increasing mass so similar properties fall into the same vertical columns o His periodic table allowed him to predict the existence of elements that hadn’t even been discovered yet (like gallium and germanium) because there were gaps where elements with certain densities should go Eka means the one beyond or the next one in the family (used it as a place holder name for the elements that were still unknown) - Later on, a physicist named Moseley put the elements in order of atomic number instead of mass, which resulted in better correlation between properties - Periodic properties (density, mass, radii, electron affinity, ionization energy) can be predicted based on where the atoms are on the table o Density and radii increase as you move down the column because the mass increases more than its volume (due to additional protons and neutrons) Ex. Platinum is one of the densest metals and is found near the bottom Density of a solid is determined by the density of the atom and how close the atoms are packed together o Structure of an atom determines its properties - The periodic table is divided into main group elements, transition elements, and inner transition elements o Main group elements are the first two columns on the left and last six on the right. Their properties are very predictable based on their position in the table o Transition (middle 10 columns in the middle) and inner transition elements (the two rows that come out of the table) tend to be less predictable o Columns = family or group elements Group 1 – alkaline metals Group 2- alkaline earth metals Group 17- halogens Group 18- noble gases o Row= periods - Since a family of elements has similar properties, we expect similar electronic structures Electron configuration - An electron configuration shows the particular orbitals that electrons occupy in an atom - In a ‘hydrogen-like’ atom (atoms with 1 only electron), all Ψ are possible (infinite solutions) but only 1 is ever occupied at a time o Same energy = same state (degenerate states) o Ground state (lowest energy state): n l m l m s 1 0 0 +1/2 1 0 0 -1/2 o Electron configuration of Hydrogen: Orbital diagram of Hydrogen: Similar to the electron configuration but symbolizes the electron as an arrow and the orbital as a box The arrow symbolizes the electrons spin Electrons occupy the lowest energy state available (1s in this case) - Helium (not ‘hydrogen-like’ because it has more than 1 electron) has 2 electrons o Electron configuration and orbital diagram of Helium: o Pauli exclusion principle= 2 electrons cannot occupy the same 4 quantum numbers in the same atom. Each orbital can have a maximum of two electrons with opposing spins - Multiple electron systems o Orbitals in a principal level of multi-electron atoms are not degenerate; they depend on the value of l Energy of s < p < d <f when n number or l number increases, energy increases More energy is less stable −Rh 2 E n n you want to make the resulting value as small as possible less energy means more stability q q 1 2+ 1 Csalomb’s law: E= r 4πE 0 (for charged particles) 1 q= charges r = distance between particles = 4πE 0 constants States that the potential energy of two charged particles depends on their charges and distance from each other As q increases (more protons in the nucleus) the absolute value of energy decreases As r decreases (electrons are closer to the nucleus), absolute energy increases (closer to the nucleus = more stable) Sign matters: large negative = very stable; large positive = very unstable Like charges (both positive) that are close together have high potential energy tend to repel each other, moving towards lower potential energy Unlike charges (one positive one negative) move closer together The magnitude of the potential energy depends inversely on r. Also, it increases as the charges of the particles increase (an electron with a 1- charge is more attracted to a nucleus with a +2 charge than one with a +1 charge Shielding: the repulsion of one electron by other electrons Any one electron can experience both the positive charge of the nucleus (which attracts it) and the negative charge of other electrons (which repulse it) Ex. Lithium has 3 electrons (in the 1s orbital, and one in the 2s orbital) Therefore, its nucleus has a 3+ charge (3 protons) The net charge that the outer electron feels is 1+, not 3+ because it is shielded by the core electrons (3-2=1) Penetration: if the outer electron enters the region occupied by the inner electrons, it experiences a greater nuclear charge and therefore less energy Ex. If the third electron in Lithium enters the electron cloud of the 1s electrons (penetrates the electron cloud) it will feel the 3+ charge, no the 1+ that was caused by shielding. Smaller l value penetrates core orbitals more efficiently Ex. 2s penetrates more efficiently than 2p How to write electron configurations: l (# of subshells) n (# of 0 1 2 3 Starting in the left corner with 1s, shells) fill in by moving from diagonally 1 1s down leftward (the way the 2s 2p arrows are pointing). 2 3 3s 3p 3d 4s 4p 4d 4f So the arrows displayed on the 4 left would read: 5 5s 5p 5d 5f 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 6p 6d 6f 6s 6 Note: (This will not be tested though) d orbitals are weird because the s orbital isn’t completely filled. This is because d orbitals can fit 10 electrons and half-filled shells are pretty stable - Chromium and copper are exceptions. Chromium places 1 electron in the 4s sublevel and 5 electrons in the 3d sublevel to produce two half- filled sublevels, which are particularly stable. Copper places 1 electron in the 4s sublevel and 10 electrons in the 3d sublevel to produce a half-filled s sublevel and a completely filled d sublevel. This is because 6 electrons) half-filled 2. Determine the principal energy level (n) of the atom (n = 2 for carbon (it is in the second period) 3. Assign electrons to the sublevels using the chart a. 1s will be filled first, with the maximum of 2 electrons. You still have four electrons left. b. 2s will be filled next, with the maximum of 2 electrons. You still have two electrons left. c. 2p will be filled next, with the maximum of 2 electrons. You don’t have any electrons left now. 4. Write the full electron configuration: 1s 2s 2p 6 a. Notice that the superscripts (the electrons in each sublevel) add up to the total number of electrons in the atom (2 + 2 + 2 = 6) Examples: 2 He has 2 electrons 1s Remember!!! Li has 3 electrons 1s 2s 1 F has 9 electrons 1s 2s 2p 5 s orbitals can hold 2 2 2 6 2 2 electrons Si has 14 electrons 1s 2s 2p 3s 3p p orbitals can hold 6 Sr has 38 electrons 1s 2s 2p 3s 3p 4s 3d 2 104p 5s 2 electrons d orbitals can hold 10 electrons Shortcut: 1. Determine what the previous noble gas is from looking at the periodic table (Ex. the noble gas which comes before carbon is helium) 2. Identify the portion of carbon’s electron configuration that is the same as helium’s electron configuration (Ex. Helium, which has 2 electrons, has an electron configuration of 1s2) 3. Put the chemical symbol of the noble gas in brackets and continue with original electron configuration to make the abbreviated electron configuration (Ex. [He]2s 2p ) - To determine which element is represented by a given electron configuration, either locate the last occupied sublevel on the periodic table and count down the period to the correct number of electrons or calculate the atomic number by adding up the exponents and the atomic number of the noble gas referenced - Elements in the same group are ‘isoelectronic’- same number and type of valence electrons - Valence electrons- outermost electrons (ones on the farthest orbital). Everything else is inner - Isotopes have the same number and kind of electrons as their counterparts because the only thing that changes is the number of neutrons; protons and electrons stay the same Hund’s Rule - When filling orbitals for the ground state (most stable), put all arrows in the same direction first and fill all equivalent orbitals before pairing You could not put an arrow up and an arrow down in a box until you fill all three boxes with at least 1 arrow - Diamagnetic means there are no unpaired electrons in the atom Paramagnetic means there is at least one unpaired electron in the atom Ions - Ions are charged particles - Atoms become ions if you change the number of electrons o If an atom gains electrons, then the number of electrons is greater than the number of protons so the atom has a negative charge- it is called an anion. o If an atom loses electrons, then the number of electrons is less than the number of protons so the atom has a positive charge- it is called a cation. - Whether atoms lose or gain electrons is determined by the number of valence electrons and how close they are to being as stable as a noble gas (having a filled valence shell) o if it has less than 3 it will lose electrons o if it has 4 it will share the electron o if it more than 4 it will gain - Alkali metals lose 1 electron to be isoelectronic with the nearest noble gas Alkaline-earth metals lose two electrons be isoelectronic with the nearest noble gas (Mg 2+= lost two electrons) Halogens gain one electron to be isoelectronic the nearest noble gas (Li = + lost 1 electron) Aluminum loses 3 electrons Nitrogen gains 3 electrons - Electron configurations examples: 1 + Na: [Ne] 3s loses 1 electron Na : [Ne] F: [He] 2s 2p gains 1 electron F : [He] 2s 2p or [Ne] - Transition metals form ions that aren’t the same as noble gas electron configurations o Electrons are removed from the highest n value orbital, even if that doesn’t correspond with the order of filling 2+ o Ex. Fe: [Ar] 4s 3d loses 2 electrons Fe : [Ar] 3d6 2 3 3+ 2 V: [Ar] 4s 3d loses 3 electrons V : [Ar] 3d Periodic Trends - Atomic Radius o distance between the nucleus of two atoms o Van der Waals Radius: space between unbonded atoms Covalent Radius: space between neutral, bonded atoms Ionic Radius: space between ionized atoms o trend is explained by a highly charged nucleus (many protons) pulling electrons in closer o also explained by the effective nuclear charge trend Atomic radius increases as you move down a column Atomic radius decreases as you move right across a row - Effective Nuclear Charge o This is the net positive charge from the nucleus experienced by an atom. Zeff Z – S Z eff effective nuclear charge Z= atomic # S = # core electrons those not in the valence shell) o It increases as you move up and/or right across the periodic table, resulting in a stronger attraction between the outermost electrons and the nucleus and smaller atomic radii. - Ionic Radii o Cations are smaller than their neutral atom (K radius < K radius) 2- Anions are larger than their neutral atoms (O radius > O radius) o They get smaller/larger depending on the number of protons they have to pull in the electrons - - Ionization Energy o Energy required to remove an electron from an atom or an ion in the gaseous state o Always positive because taking an electron requires putting in energy (endothermic) o Energy required to move the first electron is the ‘first ionization energy’ (IE1) Energy required to move the second electron is the ‘second ionization energy’(IE 2 Energy required to move the third electron is the ‘third ionization energy’(IE ) 3 o The energy required to move 2 or more electrons is the sum of all their IE’s o IE decreases down a column because a higher n number means more orbitals, meaning the electrons are farther from the nucleus, and thus held less tightly IE increases right across a period because electrons experience a greater effective nuclear charge - Electron Affinity o Energy change from gaining an electron in the gaseous state o Usually negative because energy is released when an electron is gained (exothermic) o Trends are not as typical as with other properties For most main group elements, EA increases (becomes more negative across a row) Most groups don’t experience a trend but Alkali metals (Group 1) tend to decrease (EA becomes more positive) down the column - Metallic Characters o Metals are good conductors of heat and electricity, are malleable (can be pounded into different shapes), are ductile (can be made into wires), are shiny, and tend to lose electrons in a chemical reaction Nonmetals are very varied in their properties As you move down a column, metallic character increases As you move right across a row, metallic character decreases
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