\n Properties of a gas (6) - atoms in constant random motion, fills the container it occupies, low density, compressible, mixtures are homogenous, fluid

Ideal gas assumption - often intermolecular forces are essentially negligible, size of the molecules can often be ignored

pressure is a measure of - the total amount of this push (force) exerted by gas molecules hitting the entire surface at one instant

Boyle’s Law - gases are compressible P1V1 = P2V2

At constant T and n (Boyle’s Law) - P and V are inversely proportional so as V decreases, pressure increase

Charles’s Law - Volume of a gas extrapolates to zero at absolution zero Kelvin V1/T1 = V2/T2

At constant P and n (Charles’s Law) - V and T are proportional so as T increases, V increases

Avogadro’s Law - Equal volumes of gases contain same number of moles V1/n1 = V2/n2

At constant P and T (Avogadro’s law) - V and n are proportional so as volume increases as moles increases

Amontons’s Law - Pressure of a gas increases as the temperature of the gas increases P1/T1 = P2/T2

At constant V and n (Amontons’s Law) - P and T are proportional so as T increase, P increases

Combined Gas Law - PiVi/Ti = PfVf/Tf

STP (standard temperature and pressure) - 0 degrees C, 273.15 K, 1atm, 22.4L/mol

Ideal gas constant or molar gas constant or universal gas constant (R=) - 0.08206Latm/molK

Ideal Gas Equation - PV = nRT : P in atm, V in liters, n in mols, T in K

Molecular Weight Determination - Mm = mRT/RV

Density Determination - D = MmP/RT

Dalton’s Law - sum of pressures of all different gases in a mixture equals the total pressure Ptot = Pa + Pb + Pc +…

Moles fraction X - a fraction of moles of “A” in the total moles of the mixture Xa = na/ntot = Pa/Ptot

Von der Waals - corrects for the nonideal nature of real gases (P + n^2a/v^2)(V-nb) = nRT

Kinetic-Molecular Theory - the volume of particles is negligible, particles are in constant motion, no inherent attractive or repulsive forces, and the average kinetic energy of particles is proportional to the temperature

Molecular speed (u) equation - u=√(3RT/Mm)

Diffusion - transfer of a gas through space over time

Effusion - transfer of a gas through a membrane or orifice rate of effusion is proportional to 1/√(Mm)

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The energy of a system is - its capacity to do work

Work - is done to achieve motion against an opposing force

Work= - (Force)(Distance)

When a system changes energy as a result of a difference in temperature between the system and surroundings ___ is exchanged - heat

Thermal motion - disorderly motion

Exothermic Process - a chemical reaction or physical change in which heat is evolved (q<0), heat transfers from a system to the surroundings

Endothermic Process - a chemical reaction or physical change in which heat is absorbed (q>0), heat transfers from the surroundings to a system

System - the part we are interested in

Boundary - can be physical or imaginary, permeable or impenetrable to the flow of matter, conducting or insulating to the heat, rigid or flexible

Surroundings - the rest of the universe

Kinetic Energy - energy of motion or energy that is being transferred

Potential Energy - energy that is stored in an object, or energy associated with the composition and position of the object

Internal Energy (U) - the sum of the kinetic and potential energy of the particles making up a substance

Kinetic or potential? Energy stored in the structure of a compound is - potential

Kinetic or potential? Thermal energy is - kinetic (motion of molecules)

Law of Conservation of Energy - Energy can neither be created or destroyed

Joule (J) - is the amount of energy needed to move an object by 1m with 1N force

1J = : 1M\*n=1kgm^2/s2

Calorie (Cal) - is the amount of energy needed to raise one gram of water by 1°C

1cal : = 4.184J

The internal energy of an isolated system is _ (1st law of thermodynamics) - The internal energy of an isolated system is CONSTANT

State Function - only depends on the initial and final conditions not on the process used (E and U are state functions)

Pressure Volume Work - work based on a change in volume

q=heat (thermal ) energy - the energy that flows into or out of a system because of a difference in temperature between the system and its surroundings

Heat - exchange of thermal energy, occurs when system and surroundings have a difference in temperature, heat flows from matter with high temperature to matter with low temperature until both objects reach the same temperature

Calorimetry - calculation of the amount of heat (q) from temperature change, means by which energy is transferred from a hot body to a colder body when the two are placed in thermal contact

Heat Capacity (C) - the quantity of heat required to raise the temperature by 1°C

Specific heat capacity (s) - heat required to raise the temperature of 1g of a substance by 1°C

Thermal Equilibrium - heat flows from matter with high temperature to matter with low temperature until both objects reach the same temperature

Enthalpy - the heat absorbed or evolved in a chemical reactions

Change in enthalpy for a reaction at a given temperature and pressure: H(products) minus H(reactants)

Change in enthalpy - the heat of reaction at constant pressure

△Hcond - molar enthalpy of condensation <0

△Hfreez - molar enthalpy of freezing <0

△Hvap - molar enthalpy of vaporization >0

△Hfus - molar enthalpy of fusion >0

Calorimetry at constant pressure - reactions done in aqueous solution are at constant pressure, the calorimeter is often nested foam cups containing the solution

Law of summation of heats of formation - the enthalpy of a reaction is equal to the total formation energy of the products minus that of the reactants

Standard enthalpy of formation - the enthalpy change for the formation of one mole of a substance in its standard state from its component elements in their standard state

Enthalpy of formation for a pure element in its standard state - is always 0

Heat of summation - for a chemical equation that can be written as the sum of two or more steps, the enthalpy change for this equation is then the sum of the enthalpy changes for the individual steps

Relationships Involving Hess’s Law - if a reaction can be expressed as a series of steps, then the △Hxn for the overall reaction is the sum of the heats of reaction for each step

Four “rules” of Hess’s Law - 1 Phases of reactants and products are important, 2 Temperature of the reaction is important, 3 If you reverse the reaction the sign of △rH° changes, 4 If the reaction is multiplied then △rH° must be multiplied as well

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Electromagnetic Radiation - A type of energy embodied in oscillating electric and magnetic fields traveling through space

Wavelength - The distance between any two adjacent identical points

Frequency - the number of wavelengths that pass a fixed point per second

Is light a wave? - Yes

Can light be diffracted? - Yes

Light has _ and _ interference - constructive and destructive

Electromagnetic spectrum - the range of wavelength or frequency of electromagnetic radiation

Blackbody radiation (the ultraviolet catastrophe) - when you heat a solid it can emit radiation (red-hot = 2000K, white-hot = 5000K, blue-hot > 8000K

Planck’s Quantum Theory - Energy of atoms and molecules take only discrete quantities

Planck’s Quantization of Energy - the atoms of a solid oscillate with a certain frequency

Photoelectric Effect - the ejection of electrons from the surface of metals by radiation, no electrons are ejected unless the frequency exceeds some threshold (depends on metal), the kinetic energy increases linearly with frequency but independent of intensity

Is light a particle? - Yes

Photoelectric Effect - An electron is ejected when it is struck by a single photon, it behaves like a particle, its energy is taken up by the electron

Energy level (Bohr’s Postulates) - An electron can have only specific energy levels in an atom

Transitions between energy levels (Bohr’s Postulates) - An electron can change energy levels by a “transition” from one energy level to another

When an electron undergoes a transition from a higher to a lower energy level, the energy is __ - When an electron undergoes a transition from a higher to a lower energy level, the energy is EMITTED AS A PHOTON

Is light a particle or a wave - Both (according to De Broglie)

De Broglie predicted that the wavelength of a particle was - inversely proportional to its momentum

Standing waves - only certain wavelengths can occur on a string

Circular standing wave according to De Broglie - if an electron behaves like a standing wave in a hydrogen atom the length of the wave must EXACTLY fit the circumference of the orbit

Heisenberg uncertainty principle - we cannot measure position and momentum with greaat precision simultaneously

The energy of an electron depends on its “residence” in certain orbital - a shell in a certain distance from the center, a subshell in a certain shape, the subshell with an orientation, e’s spin direction

Principal q.n. (n) - the “shell number” in which an electron “resides”, the smaller n is the smaller the orbital the lower the energy of the electron

Angular momentum g.n. (l) - distinguishes “sub shells” in a given shell that have different shapes, each main 0:s, 1:p, 2:d, 3:f, 4:g

Magnetic q.n. (ml) - distinguishes orbitals in a subshell that have different shapes and orientations in space, each subshell is subdivided into “orbitals” capable of holding a pair of electrons, and each orbital within a given subshell has the same energy

Spin q.n. (ms) - two possible spic orientations of electrons residing in a given orbital, each orbital can hold only two electrons whose spins must oppose one another

Electron Spin (ms) - electron pairs residing in the same orbital are required to have opposing spins, electrons behave like tiny bar magnets

Pauli exclusion principle - no two electrons can have the smae four quantum number

Electron Configuration - a particular distribution of electrons among available sub shells

Orbital diagram - each orbital is represented by a square or line, each group of orbitals is labeled by its sub shell notation

Electronic States - every atom has many possible electron configurations (ground state, excited states)

Aufbau Principle - A scheme to build the ground state electron configs of atoms by the “building up” order

Valance Electrons - Electrons that reside in the outermost shell, primarily involved in chemical reactions, same valence shell: similar chemical properties, elements within a given family have similar configurations

Paramagnetic Substance - weakly attracted by a magnetic field (unpaired)

Diamagnetic Substance - not attracted by a magnetic field (paired)

Periodic Law - when elements are arranged by Z, their physical and chemical properties vary periodically

Atomic Radius Trend - in a period (row) r decrease with Z increasing, in a group (column) r increases with period number

Ionization Energy - minimum energy needed to remove an electron from an atom (gas state, valence electron easiest to remove)

Periodic trend of IE - IE increase with Z increases within a given row, IE decreases as we go down a column

Successive Ionization - less electrons → stronger attraction → more IE is needed

Electron attachment - the amount of energy required to remove an electron from an anion

Periodic Trend of Electron Affinity - the more negative the EA, the more stable the negative ion; from lower left to upper right, EA becomes more negative

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Bonding - a chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of the separate atoms

Lewis structures - that allow us to predict many properties of molecules (Electron Dot Structure)

Isoelectronic group of ions - ions that have the same number and configuration of electrons

Ionic Bonds - bond formed by electrostatic attraction between (+) and (-) ion

Covalent Bonds - Bond that share valence electrons, share sufficient number of electrons in order to achieve a noble gas electron configuration

Octet rule - tendency to have 8ei in outer shells (It’s 2e- for H2)

Coordinate Covalent bond - bonds between atoms where both electrons are donated by one of the atoms

Double bond - two pairs of electrons are shared between atoms

Triple Bonds - three pairs of electrons are shared between atoms

Polar covalent bond - the bonding electrons spend more time near one of the two atoms

Nonpolar covalent bond - for alike atoms (as in H2) the bonding electrons are shared equally

Polar bond - shared not equally (HCl)

Electronegativity - ability of an atom to draw bonding electrons to itself, increases from lower-left to upper-right corner

Polarity of Bond - the absolute difference in electronegativity of two bonded atoms

Nonpolar Bond - when the difference in electronegativity is small <0.5

Polar Bond - when the difference in electronegativity is larger >0.5

Ionic Bond - when the difference in electronegativity exceeds about 1.8

Steps to drawing Lewis Structures - 1. find total of all valence e- 2. arrange atoms radially with the least electronegative in the center placing one paif of e- per bond 3. distribute remaining electrons to electronegative atoms to satisfy octet rule 4. distribute any remaining electrons to center atom

Delocalized Bonding: Resonance - can represented by two different Lewis e-dot formulas or pair is shared (O3)

Expanded Octet Exception - the central atom has more tha n8 electrons (3rd period or greater)

Fewer then 8 Electrons Exception - the central atom does not need a full octet (very small atoms smaller than C)

Formal Charge - used to determine which structure is the most likely

Bond Theory VSEPR Model - Predicts shapes of molecules by assuming that the valence e pairs are arranged as far from one another as possible (to minimize repulsion)

electron pair arrangement steps - 1. Draw Lewis structure 2. Determine how many electron pairs are around the central atom 3. Arrange electron pairs according to VSEPR 4. Obtain geometry from directions of bonding pairs

Arrangement of 2 electron pairs - linear (180 degrees)

Arrangement of 3 electron pairs - trigonal planar (120 degrees)

Arrangement of 4 electron pairs - tetrahedral (109.5 degrees)

Arrangement of 5 electron pairs - trigonal bipyramidal (90, 120 degrees)

Arrangement of 6 electron pairs - octahedral (90 degrees)

Molecular Geometry of an atom with ***2*** bonding pairs and ***0*** lone pairs - linear

Molecular Geometry of an atom with 3 bonding pairs and ***0*** lone pairs - trigonal planar

Molecular Geometry of an atom with ***2*** bonding pairs and ***1*** lone pair - bent or angular

Molecular Geometry of an atom with ***4*** bonding pairs and ***0*** lone pairs - tetrahedral

Molecular Geometry of an atom with ***3*** bonding pairs and ***1*** lone pair - trigonal pyramidal

Molecular Geometry of an atom with ***2*** bonding pairs and ***2*** lone pairs - bent or angular

Molecular Geometry of an atom with 5 bonding pairs and 0 lone pairs -Trigonal bipyramindal

Molecular Geometry of an atom with ***4*** bonding pairs and ***1*** lone pairs - Seesaw

Molecular Geometry of an atom with ***3*** bonding pairs and ***2*** lone pairs - T-Shaped

Molecular Geometry of an atom with ***2*** bonding pairs and ***3*** lone pairs - linear

Molecular Geometry of an atom with ***6*** bonding pairs and ***0*** lone pairs - Octahedral

Molecular Geometry of an atom with ***5*** bonding pairs and ***1*** lone pairs - Square Pyramidal

Molecular Geometry of an atom with ***4*** bonding pairs and ***2*** lone pairs - Square Planar

Dipole Moment - a measure of the degree of charge separation in a molecule

Nonpolar Molecule (Dipole) - perfectly symmetric, having a 0 dipole moment

Polar Molecule (Dipole) - exhibiting any asymmetry, having a nonzero dipole moment