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by: Lauren Faris


Marketplace > University of Alabama - Tuscaloosa > Science > Ch 101 > CHEM 101 STUDY GUIDE FOR TEST 2
Lauren Faris

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This study guide goes over the material in chapters 3-6 that we have gone over in class. It is everything that will be covered on the test. It also has 3 pages of worked through problems includin...
Chemistry 101 008
Dr. Bakker
Study Guide
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This 16 page Study Guide was uploaded by Lauren Faris on Friday September 30, 2016. The Study Guide belongs to Ch 101 at University of Alabama - Tuscaloosa taught by Dr. Bakker in Fall 2016. Since its upload, it has received 176 views. For similar materials see Chemistry 101 008 in Science at University of Alabama - Tuscaloosa.




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Date Created: 09/30/16
CHEM 101 STUDY GUIDE TEST 2 HOW TO By Lauren Faris Glossary  Chapter 3 Cliff Notes pg.2-4  Chapter 4 Cliff Notes pg.5-7  Chapter 5 Cliff Notes pg.8-10  Chapter 6 Cliff Notes (Beginning) pg.11  Examples of Worked Through Questions and Answers pg.12-14 1 Chapter 3 Cliff Notes  Wavelength (λ) is the distance between identical points on successive waves.  Frequency (ν) is the number of waves that pass through a particular point in 1 second (Hz = 1 cycle/s).  The speed (v) of the wave = λ x ν  Speed of light (c) in vacuum = 3.00 x 108 m/s  Energy is represented by E = h x ν where Planck’s constant (h) is: h = 6.63 x -34 10 J•s  Hertz observed that when (some) light shines on a metal surface, electrons are produced from the surface. This is called the Photoelectric effect.  One photon at the threshold frequency gives the electron just enough energy for it to escape the atom: Binding energy, ϕ  Kinetic Energy = E photon Ebinding  KE = hυ − ϕ Where (hυ) is a quantized packet of energy.  Energy = (hυ) The energy of a photon of light is directly proportional to its frequency.  E = hc/λ Or it is inversely proportional to its wavelength.  The electrons travel in orbits that are at a fixed distance from the nucleus called stationary states.  Bohr’s Model of the Atom (1913) states that e- can only have specific (quantized) energy values.  Light is emitted as e moves from one energy level to a lower energy level: E n = -R H1/n ) where n (principal quantum number) = 1,2,3, etc. and R H (Rydberg constant) = 2.18 x 10 J -18  The interaction between waves is called interference.  Constructive interference: Waves that interact so that they add to make a larger wave are said to be in phase.  Destructive interference: Waves that interact so that they cancel each other are said to be out of phase.  When traveling waves encounter an obstacle or opening in a barrier that is about the same size as the wavelength, they bend around it; this is called diffraction.  The diffraction of light through two slits separated by a distance comparable to the wavelength results in an interference pattern of the diffracted waves.  De Broglie relation: λ=h/mv  Heisenberg stated that the product of the uncertainties in both the position and speed of a particle were inversely proportional to its mass.  – x = position, Δx = uncertainty in position, v= velocity, Δv = uncertainty in velocity, m = mass  Δ- × mΔv > h/4 π  e is both a particle and wave. 2  Schrodinger’s wave equation describes both the particle and wave nature of the e- HΨ = EΨ  H is called an "operator": in this case taking the second derivative with respect to x, y and z. E is the energy. Solution is a Wave function: Ψ  A plot of distance versus ψ represents an orbital, a probability distribution map of a region where the electron is likely to be found.  These integers (solutions) are called quantum numbers and there are four quantum numbers: o Principal quantum number, n is the energy level.  Values of n can be any whole number integer ≥ 1.  It determines the size (overall) and energy of an orbital. The larger the value of n, the more energy the orbital has, and the larger the orbital. –18 2  En=–2.18 × 10 J (1/n ) where n = 1, 2, 3, etc. o Angular momentum quantum number, l is the orbital type.  l can have integer values from 0 to (n – 1).  Each value of l is designated by a particular letter that designates the shape of the orbital.  s orbitals are spherical.  p orbitals are like two balloons tied at the knots (dumbbell shape).  d orbitals are mainly like four balloons tied at the knots.  f orbitals are mainly like eight balloons tied at the knots. o Magnetic quantum number, m is thelposition of orbital in an X-Y-Z plot.  Values are integers from −l to +l.  When l = 2, the values of m arl −2, −1, 0, +1, +2, which means there are five orbitals with l = 2. o Spin quantum number, m is ths orientation of the spin of the electron.  Orbitals with the same value of n are in the same principal energy level and are also called the principal shell.  Orbitals with the same values of n and l are said to be in the same sublevel and are also called a subshell.  In general, the following is true: o The number of sublevels within a level = n. o The number of orbitals within a sublevel = 2l + 1. o The number of orbitals in a level = n . 2  When an electron is excited, it transitions from an orbital in a lower energy level to an orbital in a higher energy level.  When an electron relaxes, it transitions from an orbital in a higher energy level to an orbital in a lower energy level.  For an electron in energy state n, there are (n – 1) energy states it can transition to. Therefore, it can generate (n – 1) lines.  The energy of a photon released is equal to the difference in energy between the two levels the electron is jumping between. 3  ΔE electron E final state initial statere ΔE=-2.18*10 J(1/n -1/n ) 2f 2i  E photon= ΔE = E - f i  E = -R (1/n ), E = -R (1/n ) nf = 1 ni = 2 nf = 1 ni = 3 nf = 2 ni = 3 f i ΔE = f H f i H i -R H1/n -1fn ) 2i  Schrodinger Wave Equation: Ψ = fn(n, l, m , m ) l s o n = 1, 2, 3, etc. is the distance of e from the nucleus. o l for a given value of n, l = 0, 1, 2,3…n-1 Shape of the “volume” of space that the e- occupies.  n = 1, l = 0  l = 0 s orbital  n = 2, l = 0 or 1  l = 1 p orbital  n = 3, l = 0, 1, or 2  l = 2 d orbital o m il orientation of the orbital in space. For a given value of l m = -l, l …., 0, …. +l o Spin quantum number m : m = +1/2sor -s/2 4 Chapter 4 Cliff Notes  Existence (and energy) of electron in atom is described by its unique wave function Ψ.  Pauli exclusion principle - no two electrons in an atom can have the same four quantum numbers.  Coulomb’s law describes the attractions and repulsions between charged particles. o For like charges, the potential energy (E) is positive and decreases as the particles get farther apart as r increases. o For opposite charges, the potential energy is negative and becomes more negative as the particles get closer together. o The strength of the interaction increases as the size of the charges increases. Electrons are more strongly attracted to a nucleus with a 2+ charge than to a nucleus with a 1+ charge.  From radial distribution function: 2s orbital penetrates more deeply into 1s orbital than does 2p.  The weaker penetration  electrons in the 2p sublevel experience more repulsive force & are more shielded from nucleus (less attractive force).  In the fourth and fifth principal levels, the effects of penetration become so important that the s orbital lies lower in energy than the d orbitals of the previous principal level.  Penetration  sublevels are not degenerate, 4s<3d, 5s<4d, etc., ΔE less as n increases.  Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom.  Aufbau Principle: Energy levels and sublevels fill from lowest energy to highest: – s → p → d →f  Pauli exclusion principle states that there can be no more than two electrons per orbital.  The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (Hund’s rule).  Orbitals fill in the following order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s  Orbitals can hold no more than two electrons each. When two electrons occupy the same orbital, their spins are opposite.  When orbitals of identical energy are available, electrons first occupy these orbitals singly with parallel spins rather than in pairs (Hund’s rule).  1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s 2 2 6 2 6 2 10 2  Electron configuration Short Cut: Ge: 1s 2s 2p 3s 3p 4s 3d 4p easier: [Ar] 4s 3d 10 4p i.e. write [last noble gas] remainder of orbitals.  Periodic law: – When the elements are arranged in order of increasing mass, certain sets of properties recur periodically.  Rows of the table are referred to as periods (or just plain rows).  Columns in the table are referred to as groups or a family. 5  Elements in the periodic table are classified as the following: Metals, Nonmetals, and Metalloids.  Main-group elements properties tend to be largely predictable based on their position in the periodic table.  Transition elements (or transition metals) and inner transition metals properties tend to be less predictable based simply on their position in the periodic table.  The group number corresponds to the number of electrons in last (unfilled valence shell) called valence electrons. Core electrons are in inner shell.  Transition metals (d block) and inner transition metals (f block) exhibit trends differing from those of main-group elements (s block and p block). Because of sublevel splitting, the 4s sublevel is lower in energy than the 3d sublevel; therefore, the 4s orbital fills before the 3d orbital.  Elements in the same column have similar properties.  The elements in a period show a pattern that repeats.  The noble gases have eight valence electrons. Except for He, which has only two electrons.  Alkali Metals have one more electron than the previous noble gas and occupy the first column. In their reactions, the alkali metals lose one electron, and the resulting electron configuration is the same as that of a noble gas. Forming a cation with a 1+ charge.  Alkaline Earth Metals have two more electrons than the previous noble gas and occupy the second column. In their reactions, the alkaline earth metals lose two electrons, and the resulting electron configuration is the same as that of a noble gas. Forming a cation with a 2+ charge.  Transition and Inner Transition Metals are located in the d-block area of the periodic table. In chemical reactions, they will lose electron(s) from s and then d orbitals to form cations.  p-block metals are located in the p-block area (left-hand side of the metalloids) of the periodic table. In chemical reactions, they will lose electrons from the s and p orbitals to form cations.  Metalloids are located in the d-block area of the periodic table between the metal and nonmetal elements. Sitting on the “steps” of the zigzag diagonal line indicated on the periodic table. Metalloids in chemical reactions can exhibit metallic or nonmetallic behaviors. Metalloids can either lose electron(s) from s and then p orbitals to form cations or gain electrons into their p orbitals to form anions.  Nonmetals are located in the upper right-hand side of the periodic table. p- block area. In chemical reactions nonmetal elements will gain electrons into the p orbitals, resulting in their ions having the same electron configuration as a noble gas at the end of their period (row). Nonmetals form anions.  Halogens are nonmetals. They have one fewer electron than the next noble gas. In their reactions with metals, the halogens tend to gain an electron and attain the electron configuration of the next noble gas, forming an anion with charge 1−.  Metals form cations (positively charged atoms). Alkali metals (group 1A) form only +1 cations. Alkaline earth metals (group 2A) form only +2 cations.  Nonmetals form anions (negatively charged atoms). Halogens (group 7A) usually gain one electron to form –1 anions. 6  The effective nuclear charge is a net positive charge that is attracting a particular electron.  Z is the nuclear charge, and S is the number of electrons in lower energy levels. Z effective – S  Atomic radius decreases across a period (left to right).  Atomic radius increases down a group.  The larger the effective nuclear charge an electron experiences, the stronger the attraction it will have for the nucleus.  Atoms in the same group increase in size down the column.  Atomic radii of transition metals are roughly the same size across the d block.  Electron configurations that result in unpaired electrons mean that the atom or ion will have a net magnetic field; this is called paramagnetism.  Electron configurations that result in all paired electrons mean that the atom or ion will have no magnetic field; this is called diamagnetism.  Isoelectronic= same number of electrons.  Ions in the same group have the same charge.  Ion size increases down the column.  Cations are smaller than neutral atoms; anions are larger than neutral atoms.  Ionization Energy (IE) is the minimum energy needed to remove an electron from an atom or ion in the gas phase. It is an endothermic process (requires the input of energy to remove the electron).  First IE decreases down the group.  First IE generally increases across the period.  Electron Affinities (EA) increases across a period.  Metals are malleable and ductile, shiny, lustrous, reflect light, conduct heat and electricity, most oxides basic and ionic, form cations in solution, lose electrons in reactions, and are oxidized.  Nonmetals are brittle in solid state, dull, nonreflective, solid surface, electrical and thermal insulators, most oxides acidic and molecular, form anions and polyatomic anions, and gain electrons in reactions making them reduced.  Metallic character is how closely an element’s properties match the ideal properties of a metal.  Metallic character decreases left to right across a period.  Metallic character increases down the column.  Metals generally have smaller first ionization energies and nonmetals generally have larger electron affinities. 7 Chapter 5 Cliff Notes  When two or more elements combine, a molecule can be formed.  When two or more elements combine to form a compound, an entirely new substance results.  Compounds are made of atoms held together by chemical bonds. The attraction is electrostatic: between protons and electrons.  Two general types of bonding between atoms found in compounds, ionic and covalent. o Ionic bonds result when electrons have been transferred between atoms, resulting in oppositely charged ions that attract each other. They are generally metal atoms bonded to nonmetal atoms. o Covalent bonds result when two atoms share some of their electrons. They are generally found when nonmetal atoms bonded together.  A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance.  An empirical formula shows the simplest whole-number ratio of the atoms in a substance.  Ionic compounds are metals and nonmetals. There are no individual molecule units, instead they have a 3-dimensional array of cations and anions made of formula units. Many contain polyatomic ions.  Octet Rule: Atoms prefer to have a filled valence shell meaning 8 valence electrons.  Lewis theory predicts the number of electrons a metal atom should lose or a nonmetal atom should gain in order to attain a stable electron arrangement.  Exceptions to the octet rule: H, Li, Be, and B attain an electron configuration like that of He. o He can have ONLY two valence electrons, a duet. o Li loses its one valence electron. o H shares or gains one +lectron. – Though it commonly loses its one electron to become H . o Be loses two electrons to become Be 2+ o B loses three electrons to become B 3+  The extra stability that accompanies the formation of the crystal lattice is measured as the lattice energy. It is the energy released when the solid crystal forms from separate ions in the gas state. It always exothermic.  Ionic solids are relatively hard.  Ionic solids are brittle. When struck, they shatter.  Ionic solids do not conduct electricity.  Ionic compounds conduct electricity in the liquid state or when dissolved in water.  Ionic compounds always contain positive and negative ions. In a chemical formula, the sum of the charges of the positive ions (cations) must equal the sum of the charges of the negative ions (anions). The formula of an ionic compound reflects the smallest whole-number ratio of ions.  Organic compounds: predominantly carbon Usually plus H, O, N, S. 8  Inorganic compounds: everything else.  Ionic Compounds: Cation followed by anion +-ide.  Binary compounds contain only two different elements. The names of binary ionic compounds take the following form: name of cation (metal) + base name of anion (nonmetal) +-ide. o Ex. the name for KCl consists of the name of the cation, potassium, followed by the base name of the anion, chlor, with the ending -ide. – KCl is potassium chloride.  The Lewis Model: Valence electrons are represented as dots.  Lewis electron-dot structures (Lewis structures) depict the structural formula with its valence electrons.  When atoms bond, they tend to gain, lose, or share electrons to give a noble gas–like configuration.  Nonmetals: period 2 elements must obey the octet rule.  Most polyatomic ions are oxyanions, anions containing oxygen and another element.  If there are two ions in the series, the one with more oxygen atoms has the ending -ate; and the one with fewer has the ending -ite. o Ex. NO i3 nitrate, NO i2 nitrite.  If there are more than two ions in the series, then the prefixes hypo-, meaning less than, and per-, meaning more than, are used. o Ex. ClO is hypochlorite, ClO 2s chlorite, ClO 3s chlorate, ClO i4 perchlorate.  Hydrates are ionic compounds containing a specific number of water molecules associated with each formula unit.  Electrons that are shared by atoms are called bonding pairs.  Electrons that are not shared by atoms but belong to a particular atom are called lone pairs.  When two atoms share one pair of electrons, the result is called a single covalent bond (two electrons). H only duet.  When two atoms share two pairs of electrons, the result is called a double covalent bond (four electrons between the two atoms). Elements that can double-bond with each other and themselves are C, N, O, S, and P.  When two atoms share three pairs of electrons, the result is called a triple covalent bond (six electrons between the two atoms). Elements that can triple-bond with each other and themselves are C, N, O, and S.  Lewis theory of covalent bonding implies that the attractions between atoms are directional.  Molecular compounds are composed of two or more nonmetals. st  Binary molecular compounds are structured as follows: prefix + name of 1 element + prefix + base name of 2 ndelement + -ide.  Molecular mass (or molecular weight) is the sum of the atomic masses (in amu) in a molecule.  For any molecule molecular mass (amu) = molar mass (grams).  Use molar mass to convert to the amount in moles. Then use Avogadro’s number to convert to number of molecules: g  mol  molecules. 9  Mass percent composition of an element in a compound is (n * molar mass of element)/(molar mass of compound) * 100% where n is the number of moles of the element in 1 mole of the compound.  The mass percent tells you the mass of a constituent element in 100g of the compound.  The molecular formula is a multiple of the empirical formula. Molecular Mass is the same multiple of the mass of the empirical formula.  Empirical Formula is the simplest whole-number ratio of the atoms of elements in a compound.  A common technique for analyzing compounds is to burn a known mass of compound and weigh the amounts of product made. This is generally used for organic compounds containing C, H, or O. o All the original C forms CO , the original H forms H O, and the original 2 2 mass of O is found by subtraction.  Organic compounds are mainly made of C and H, sometimes with O, N, P, S, and trace amounts of other elements. They are mainly made up of C.  Carbon atoms bond almost exclusively covalently.  There are two main categories of organic compounds: hydrocarbons and functionalized hydrocarbons.  Hydrocarbons contain only C and H.  Functional groups are non-carbon groups that are on the molecule. 10 Chapter 6 Cliff Notes (Beginning)  Molecular compounds have low melting points and boiling points.  Melting and boiling involve breaking the attractions between the molecules, but not the bonds between the atoms.  Molecular compounds do not conduct electricity in the liquid state.  Molecular acids conduct electricity when dissolved in water, but not in the solid state.  When dissolved in water, molecular acids are ionized, and have the ability to move through the structure and therefore conduct electricity.  Electronegativity (EN) is a measure of an elements ability to pull electrons toward it.  The ability of an atom to attract bonding electrons to itself is called electronegativity.  EN Increases across a period (left to right) and decreases down a group (top to bottom).  Dipole moment, µ, is a measure of bond polarity. A dipole is a material with a + and − end. It is directly proportional to the size of the partial charges (q) and directly proportional to the distance (r) between them. o µ (dipole moment) = (q)(r) Where it is measured in Debyes, D.  If the difference in electronegativity between bonded atoms is 0, the bond is pure covalent.  If the difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent.  If the difference in electronegativity between bonded atoms is 0.5 to 1.9, the bond is polar covalent.  If the difference in electronegativity between bonded atoms is larger than or equal to 2.0, the bond is ionic.  The percent ionic character is the percentage of a bond’s measured dipole moment compared to what it would be if the electrons were completely transferred.  Resonance, used when two or more valid Lewis structures can be drawn for the same compound. o Extensions of Lewis theory suggest that there is some degree of delocalization of the electrons; the concept is called resonance.  Formal charge, an electron bookkeeping system that allows us to discriminate between alternative Lewis structures. o {(# of valence electrons) – (# of nonbonding electrons)- (1/2 × # of bonding electrons)} 11 Examples of Worked Through Questions and Answers 12 13 14


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