Chem 105 Lecture Notes pt 2 PDF

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9/24/2019 ❖ Bring Periodic Table from now on! ●

Periodic Properties- predictable properties based upon the location of an element on the periodic table ● Arranged by increasing atomic number ● HORIZONTAL ROWS - periods ● VERTICAL ROWS - groups Certain groups that we need to know, 1A is the Alkali Metals 2A is the Alkaline Earth Metals 7A is the Halogens 8A are the Noble Gases Elements in the same groups display similar properties due to their electron configurations. Electron configurations tell us how we show the particular orbitals that electrons occupy in an atom. Ground state electron configuration - lowest Energy electron configuration Pauli’s Exclusion Principle ● No two electrons will have the same set of quantum numbers ● The spin (msub (s)) is -½ or +½ ● +½ is up arrow ● -½ is down arrow ● Each orbital can only be occupied by two electrons Orbitals with the same energy are degenerate FOR example: 2s + 2p Hunds Rule ● When filling orbitals of some Energy, each orbital must get 1 electron before we start pairing. ○ Carbon has electron configuration of 1s2 2s2 2p2 ○ Draw the arrows and lines underneath 1s on bottom then 2s then 2p

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Aufbau ○ Lowest energy orbitals are filled first ○ s->p->d->f The period number corresponds to n Groups 1,2,13-18 are known as the MAIN GROUP ELEMENTS 1A, 2A, 3A-8A Group 3-12 are transition metals Group 1+2 have the s block Group 13-18 have the p block Group 3-12 have the d block

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Transition metals -- period #-1 Ti atomic # is 22, 22 electrons The last electrons in Ti will be in a 3d orbital 1s2 2s2 2p6 3s2 3p6 4s2 3d2 Cl + Ca

Ca 20 e 1s2 2s2 2p6 3s2 3p6 4s2 LOOK at highest energy level and its s and p orbitals [Ar] 4s2 Cl 17 e 1s2 2s2 2p6 3s2 3p5 Must consider last two configuration values for valence electrons Cl has 7 valence electrons Electrons are stripped from 4s orbital before 3d orbital Ti has 4 valence electrons Group 1 has 1 valence electron Neon has 10 electrons. 1s2 2s2 2p6 [Ne] 3s2 3p5 (for chlorine) ***Make sure to fill all lines before spin pairing electrons*** DRAW Stair step line on periodic table under boron ends on 5 d Metals are left of the stair step line. They form cations. Nonmetals are right of the stair step line. Group 1 forms 1+ ions Group 2 forms 2+ ions. Ex. Calcium has 2 valence electrons, so it gives up 2 so it can have a filled outer shell. When losing electrons, redo electron configuration accordingly [Ar]4s2 3d10 4p5 valence electron = 7 because 4s2 +4p5 the 2+5 only consider the highest energy level… Points of stability with electron configuration. Full s + p orbitals for a given energy level Full d orbitals Half full d orbitals Cr has electron configuration of [Ar] 4s1 3d4

Metals are good conductors of electricity and heat Malleable (flatten), and ductile (pull into wire) They have a tendency to lose electrons to form cations. Most stable groups are the noble gases. Nonmetals are poor conductors. The physical state of nonmetals can vary. They have a tendency to gain electrons and form anions. They gain enough electrons to get to the nearest noble gas electron configuration. Periodic Trends Atom size-- Atomic size will increase as you go down a group.This is because we start to fill orbitals that are farther away from the nucleus. Left to right across a period: atom size DECREASES, this is because as you add more electrons, you are also adding another proton so there is a higher effective nuclear charge which allows protons to pull electrons closer to the nucleus. F, S, Ba Smallest to largest Isoelectronic: same electron configurations Anions will be larger than normal atoms Ionization energy: the energy required to remove an electron from an atom or ion. Mg Mg2+ always For an element, M+ a certain amount of energy IE decreases as we go down a group because the e- are farther away from the nucleus. As we go from left to right on a period, IE increases as a general trend because you get to nonmetals who want to gain electrons. Chapter 4 Chemical bonds: force that holds atoms together Two major types of bonds: Ionic bonds - metal + nonmetal (electrons are transferred) The attraction is held together by the charged ions Covalent bonds - results from the sharing of electrons 2 nonmetals

Empirical formula- smallest whole number ratio of the atoms within a molecule Molecular formula - the actual number of atoms of each element in a compound C4H8 ->CH2 Molecular to empirical formula CCl4 Structural formula-shows the connectivity of the atoms within the molecule Lewis Dot structure: The first atom is normally the central atom Octet rule - want to have 8 valence electrons around all the atoms in molecule (except H which is 2 electrons) Octet rule really only applies to C, N, O, F There are many exceptions to the rule. Exceptions: 1. Nonmetals in period 3 follow the octet rule if they are not central atom of a molecule. a. The central atom is the atom that the others are bonded to in the molecule 2. Skeletal structure 3. Fill in the octet around the outer atoms

When nonmetals period 3 or lower on the period table Ex: period 4-and greater Are the central atom, may exceed the octet rule. THis is because of the vacant D orbitals ICl3 B, Be, and Al May form stable compounds with fewer than 8 electrons around central atom BH3 10/3/2019 Nomenclature Ionic compound Metal + nonmetal Cation+Anion There is NH4+ which is the ammonium ion. Ionic compounds - Brittle solids - High melting point - NaCl mp=801C - Solid form - do not conduct electricity - Liquid form or in aqueous solutions - conduct electricity

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In liquid form, the compound dissociates into its individual ions - This is what allows the conduction of electricity because of the ions Naming Ionic Compounds Mg when it forms an ion is always 2+. F when it forms an ion is always 1Charges must cancel out Expressed has smallest whole # ratio MgF2 Mg2+ FCa2+ O2CaO Cross over method Group 1 is always +1 Group 2 is always +2 Al is always +3 Zinc is always +2 Zinc is at 3d10 but it will give up its 4s electrons so it can keep the d orbital full Ag is always +1 Metal + anion(-ide) Na + Cl Sodium Chloride Oxygen → oxide Sulfur → sulfide Nitrogen → nitride MgF2 Magnesium Fluoride Halogens → -1 Oxygen Sulfur → -2 Nitrogen → -3 Aluminum Oxide Al2O3 FeCl2 -- iron (II) chloride FeCl3 -- iron (III) chloride With ions that have multiple charges Metal (Roman numeral that indicates the charge) + anion (-ide) TiO2 Titanium (IV) oxide Binary compounds

Polyatomic ions Composed of two or more elements List of ions-a.)Acetate C2H3O2- or CH3CO2b.)Carbonate CO3 2c.) bicarbonate HCO3 1d.)Nitrite NO2 1e.)Nitrate NO3 1f.)Phosphate PO4 3g.)Hypochlorite ClO 1h.) Chlorite ClO2 1i.) Chlorate ClO3 1j.) perchlorate ClO4 1k.)Sulfite SO3 2l.) sulfate SO4 2m.)Cyanide CNn.) hydroxide OHo.) ammonium NH4 + Examples: KNO2 Potassium Nitrite AlSO4 Al2(SO4)3 Barium Nitrate Ba(NO3)2 NH4Cl Ammonium Chloride Prefix + element name + prefix +second element with the (-ide) ending Mono =1 Di = 2 Tri = 3 Tetra = 4 Penta = 5 Hexa = 6 CO2 Always list in order it appears in Formula Carbon Dioxide Omit mono from name for carbon

H2O Dihydrogen Monoxide P2O5 Diphosphorus Pentoxide Formula Mass Molecular Mass or Molecular Weight Molar mass - Mass of one mole of a compound H2O = 2(1.01) + 1(16) = 18.02g/mol For every one mole of water, there is 2 hydrogen atoms and 1 mole of oxygen How many molecules are in 19.3 grams of C8H10?...