Chemistry 161 General Chemistry I PDF

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D Brayton...

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Chemistry 10.13 Chemistry Notes (10.13) -electrons are incredibly small -election behavior determines much of the behavior of atoms

Quantum-Mechanical Model: explains the manner electrons exist and behave in atoms helps us understand and predict the properties of atoms that are directly related to the behavior of the electrons

electromagnetic radiation: light is formed from this ** speed of light: 3.00 x 10^(8) m/s in a vacuum

Characterizing Waves:amplitude-  the height of the wave wavelength- a measure of the distance covered by the wave frequency- (v) the number of wave that pass a point in a given time

The Relationship between Wavelength and Frequency: for waves traveling at the same speed, the shorter the wavelength, the more frequently they pass, meaning that the wavelength and frequency of electromagnetic wave are inversely proportional:

v(s^(-1) = c(m/s)/…

Color: The color of light is determined by its wavelength; you need to know the difference in energy between the types of light.

Electromagnetic Spectrum: visible light comprises only a small fraction of all the wavelengths of light (short wavelength (high frequency) light has high energy) < these can potentially damage biological molecules

Interference: the interaction between waves Constructive Interference: when waves interact so that they add to make a larger wave Destructive Interference: when waves interact so they cancel each other out

Diffraction: the process by which a beam of light or other system of waves is spread out as a result of passing through a narrow aperture or across an edge, typically accompanied by interference between the waveforms produced

Ejected Electrons: Kinetic Energy = E(photon) - E(binding) KE = hv - change

** Threshold Frequency: observed that there was a maximum wavelength for electrons to be emitted

Einstein’s Explanation:  The light energy was delivered to the atoms in packets called quanta or photons: energy of a photon of light was directly proportional to its frequency

-emission spectra: a spectrum of the electromagnetic radiation emitted by a source

Bohr’s Model and Emission Spectra

Chemistry 10.15 Chemistry Notes (10.15)

Not on TEST: Slides 55, 56, 45, 44, 43, 41, 53, 54, 69, 71 On TEST: 60-68, 73

Bohr’s Model: Neils Bohr proposed that the electrons could only have very specific amounts of energy - quantized:   fixed amounts - stationary states: the electrons traveled in orbits that were a fixed distance from the nucleus Electron Energy: electron energy and position are complimentary - orbital: for an electron with a given energy, the best we can do is describe a region in the atom of high probability of finding it Wave Function:Calculations show that the size, shape and orientation in space of an orbital are determined be three integer terms in the wave function. Electron Transitions: In order to transition to a higher energy state, the electron must gain the correct amount of energy corresponding to the difference in energy between the final and initial states. Electrons in high energy states are unstable and tend to lose energy and transition to lower energy states. ** Hydrogen Energy Transitions Probability & Radial Distribution Functions: represents the total probability at a certain distance from the nucleus  nodes where the probability drops to 0 in the function The shapes of Atomic Orbitals: the l quantum number primarily determines the shape of the orbital; l can have integer values from 0 to (n-1); each value of l is called 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 d orbitals are like four balloons l orbitals f orbitals Energy shells and subshells **

Orbitals & Subshells On the TEST: Slide 20, 22, 23, 25-26, 28, 39, 41 Electrons Configurations: the ground  state of the electron is the lowest energy orbital it can occupy -electron configuration: the  distribution of electrons into the various orbitals in an atom in its ground state - the # designates the principal energy level - the letter designates the sub level an type of orbital - the superscript designates the number of electrons in that sub level Definitions Paramagnetic: odd energy Diamagnetic: even energy Degenerate: same energy

Orbital Diagram: unoccupied orbitals have no arrows in the box; orbital w/ 1 electron has ^; orbital w/ 2 electrons has ^v Understand penetration and shielding: Penetration: because of this, sub levels within an energy level are not degenerate - Penetration of the 4th and higher energy levels is so strong that their s sub level is lower in energy that the d sublevel of the previous energy level. Order of Subshell Filling in Ground State Electron Configurations(Slide 18) energy shells fill from lowest energy to high subtle fill from lowest energy to high s > p >d >f orbitals that are in the same subtle have the same energy no more than 2 electrons per orbital when filling orbitals that have the same energy, place one electron in each before completing pairs Valence Electrons: electron configuration Valence Electrons: The electrons in all the subtle with the highest principal energy shell. Transition Metals: For the block metals, the principal energy level is one less than valence shell one less than the period number sometimes s electron “promoted” to d sublevel

Periodic Trends Chapter 8: Periodic Trends Electron Configuration of Cations in their Ground State Trends in Ionic Radius: Ions in same group have same charge (51) Ionization Energy: minimum energy needed to remove an electron from an atom gas state endothermic process valence electron easiest to remove General Trends in 1st Ionization Energy: larger the effective nuclear charge on the electron, the more energy it takes to remove it the farther the most probable distance the electron is from the nucleus, the less energy it takes to remove it 1st 1E decreases   down the group : valence electron farther from nucleus 1st generally increases across the period : effective nuclear charge ++ Metal Characters:  ● malleable & ductile ● shiny, lustrous, reflect light ● conduct heat and electricity ● most oxides basic and ionic ● from cations in solution ● lose electrons in reactions Nonmetal Characters: ● brittle in solid state ● dull ● electrical and thermal insulators ● most oxides are acidic and molecular ● form anions and polyatomic anions ● gain electrons in reactions - reduced Trends in the Noble Gas: ● atomic radius increases down the column ● ionization energy decreases down the column CHEM II Review:

Understand real world questions: global warming, hydrogen fuel cells

Specific heat x speed of light & wavelength equations

Endothermic vs exothermic

PV = nRT P (pressure) = atm V (volume) = L n ( moles) = g x mol/g R (rate) = 0.0821 T (temperature) = K

Paramagnetic x diamagnetic

Chapter 9-10 Bonding Theories: ● Lewis Theory ● Lewis structures: allows us to predict many properties of molecules ○ aka Electron Dot Structures ○ such as molecular shape, size, polarity Why do Atoms Bond? The atoms have lower energy ** ● processes are spontaneous if they result in a system with lower potential energy ● chemical bonds form because they have lower potential energy between charged particles

Potential Energy Between Charged Particles: ● 8.85 x 10 ^(-12) C^(2)/Jm ● charges w/ the same, potential energy is + (magnitude gets less positive as the particles get farther apart) ● charges w/ opposite signs, potential energy is - (magnitude gets more negative as the particles get closer together)

Bonding: 1. a chemical bond forms when the potential energy of the bonded atoms is less that the potential energy of the separate atoms Types of Bonds: ● metals to nonmetals > ionic > electrons transferred

Ionic Bonds: when metals bonds to nonmetals, some electrons from the metal atoms are transferred to the nonmetal atoms ■ metal have low ionization energy, relatively easy to remove an electron from ■ nonmetals have high electron affinities, relatively goof to add electrons to nonmetals to nonmetals > covalent > electrons shared metal to metal > metallic > electrons pooled

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Lew Symbols of Ions (13): ● Cations have Lewis symbols without valence electrons ○ Lost in the cation formations ● Anios have Lewis symbols with 8 valence electrons ○ electrons gained in the formation of the anion ● Octet Rule >> when atoms bond, they tend to gain, lose or share electrons to result in 8 valence electrons EXCEPT H, Li, Be, B

Properties of Ionic Compounds

Lew Theory and Ionic Bonding ● Lewis symbols can be used to represent the transfer of electrons from metal atoms to nonmetal atom, resulting ….

Lattice Energy: the energy released when the solid crystal forms from separate ions in the gas state ● always exothermic ● hard to measure directly, but can be calculated from knowledge of other processes. ● ● ** (ON TEST) Born-Haber Cycle (25) ● method for determining the lattice energy of an ionic substance by using other reactions ● use Hess’s Law to add up heats of other processes 11.5 Covalent Bonding: Bonding and Lone Pair Electrons ● Covalent bonding results when atoms share pairs of electrons to achieve an “octet”

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 (non-binding pairs) Lewis Theory: ● Allows us to predict the formula of molecules

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Model vs. Reality: ● 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 ○ the covalent bonds are strong ○ the attractions….

Electronegativity and Bond Polarity: ● If difference in electronegativity between bonded atoms is 0, the bond is pure covalent (sharing) ● If difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent ● If difference is 0.4 to 0.9...

Bond Dipole Moments: ● the dipole moment is quantitative way of describing the polarity of a bond ○ a dipole is a material with positively and negatively charged ends ○ measured ● Dipole moment: a measure of bond polarity

Water - a polar molecule

Slide 57

11.10

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Bond energies ○ Shorter = stronger ○ More electrons = stronger Bonding Theory (Ch.10) ○ Using Bond Energies to Estimate ‘Change in heat' ○ average bond energies

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○ bond breaking is endothermic VH(breaking) = + ○ bond making is exothermic VH(making) = Estimating the Enthalpy Change of a Reaction from Bond Energies ○ Estimate the enthalpy of the following reaction (83) Bond Lengths: ○ the distance between the nuclei of bonded atoms is called bond length ○ because the actual bond length depends on the other atoms around the bond we often use the average bond length ○ Trends in Bond Lengths: ■ the more electrons two atoms share, the shorter the covalent bond ■ decreases from left to right across period ■ increases down the column ■ as bonds get longer, they get weaker Chemical Bonding II VSEPRE Theory ○ valence shell electron pair repulsion Tetrahedral Geometry

11.12.15 Molecule Polarity: ● The O-C bond is pool. The bonding electrons are pulled equally toward both O ends of the molecule. Net Result = non polar molecule ● The N-H bond is polar.

Molecular Polarity Affects Solubility in Water: ● polar molecules attracted to other polar molecules ● water is polar therefore other polar molecules dissolve well in water ○ as well as ionic compounds ● molecules can have both polar and nonpolar parts

** Polar LIKES water ** Non polar has a water phobia

[practice equations on slide 69-70, 90]

Valence Bond Theory: ● atoms would rise when orbitals on those atoms interact to make a bond ● interaction depends on whether the orbitals align along the axis …

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orbital interaction: as two atoms approached, the partially filled or empty valence atomic orbitals on the atoms would interact to form molecular orbitals Hybridization

Valence Bond Theory: ● Valence electrons in an atom reside in the quantum mechanical atomic orbitals or hybrid orbitals ● hybridization: mixing different types of orbitals to make a new set of degenerate orbitals ● some atoms hybridize their orbitals to maximize bonding ● better explain observed shapes of molecules

Hybrid Orbitals: ● Hydrogen can’t hybridize ● # of standard atomic orbitals combined = 3 of hybrid orbitals formed ● 3 and type of standard atomic orbitals combined determines the shape of the hybrid orbitals

Types of Bonds: ● Sigma Bond: results when the bonding atomic orbitals point along the axis connecting the two bonding nuclei ○ either standard atomic orbitals or hybrids ● Pi Bond: results when the bonding atomic orbitals are parallel to each other and perpendicular to the axis connecting the two bonding nuclei ○ between un-hybridized parallel p orbitals ● the interaction between parallel orbitals is not as strong as between orbitals that point at each other; therefore, sigma bonds are stronger than pi bonds...