General Honors Chem Notes

Front 1

percent error

Back 1

= ((accepted value - experimental value) / (accepted value)) x 100

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density

Back 2

grams / meter cubed

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atomic number

Back 3

(symbol Z) = number of protons

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isotopes

Back 4

'atoms of a given element that differ in the number of neutrons, and consequently in mass

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mass number

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(symbol A) = number of protons + number of neutrons

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all common naturally existing diatomic molecules

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H2, O2, N2, F2, Cl2, Br2, I2

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oxidation

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loss of electrons in ion

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reduction

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gain of electrons in ion

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molecular formula

Back 9

chemical formula that indicates the actual number and types of atoms in a molecule

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empirical formula

Back 10

chemical formula that gives only the relative number of atoms of each type in a molecule

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electrolyte

Back 11

substances which produce ions in solution- will conduct an electric current

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stoichiometry

Back 12

the study of quantitative relationships in chemical reactions

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reagent

Back 13

the limiting reactant in a chemical reaction- the chemical that is completely consumed in the reaction

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theoretical yield

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the quantity of product that is calculated to form when all of the limiting reactant reacts

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actual yield

Back 15

the amount of product actually obtained in a reaction

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percent yield

Back 16

((actual yield) / (theoretical yield)) x 100

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mole

Back 17

6.02 x 10^23 atoms or molecules per substance being measured

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exothermic reaction

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when a chemical reaction releases energy to the surroundings (written as a product in a chemical equation i.e. + energy)

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endothermic reaction

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when a chemical reaction requires more energy to break the bonds in the reactants than it gives off by forming products

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synthesis reaction

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the combination of two or more substances to form a compound

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decomposition reaction

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(the opposite of a synthesis reaction) one substance breaks down to form two or more simpler substances i.e. AB --> A + B

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single- replacement reaction

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occur when one element is replaced by another in a compound ( A + BC --> AC + B )

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double- replacement reaction

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occur when ions of two compounds exchange places in an aqueous solution to form two new compounds ( AX + BY --> BX + AY )

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precipitation reaction

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(often double- replacement reactions) reaction between two aqueous solutions where the formation of a solid occurs

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acid- base reactions

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acid + base --> salt + water ( HB + XOH --> XB + HOH )

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combustion reaction

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the reaction of a substance with oxygen, releasing a large amount of energy in the form of light and heat

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formula mass

Back 27

atomic mass of an ionic compound

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molar mass

Back 28

the mass, in grams, of 1 mole of a substance- or the mass of 6.02x10^23 formula units of a substance

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hydrate

Back 29

compounds in which molecules of water are attached to the lattice structure of an ionic compound (waters of hydration)q

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molarity (M)

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concentration of a solution- moles of solute per volume of solution in liters: (moles of solute) / (liters of solution, or solvent)
use brackets [ ] to indicate molarity

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solute

Back 31

substance of less quantity in a homogeneous mixture or solution

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solvent

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substance of greater quantity in a solution

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Pressure

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force exerted per unit area
P= (force) / (area)
1 Pascal (Pa) = 1 Newton (N) / m^2

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force

Back 34

= (mass x acceleration)

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Newton

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(N) = 1 kg x m / s^2

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Charles Law

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when the pressure and number of moles of a gas are held constant, the volume of the gas in directly proportional to it's kelvin temperature:
(V1) / (T1) = K = (V2) / (T2)

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Boyle's Law

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when the temperature and number of moles are held constant, the volume of a gas is inversely proportional to the pressure applied on a gas:
(P1)(V1) = k = (P2)(V2)

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Avogadro's Law

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the volume of a gas is directly proportional to the number of moles if the temperature and pressure remain constant:
(V1) / (n1) = constant = (V2) / (n2)

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Gay- Lussac Relationship

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states that pressure is directly proportional to temperature:
(P1) / (T1) = k = (P2) / (T2)

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Combined Gas Law

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where # of moles are constant but pressure, volume, temperature conditions are all changing:
(P1) (V1) / (T1) = k = (P2) (V2) / (T2)

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The Ideal- Gas Equation

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PV = nRT where R = ideal gas constant, .082 (atm)(L)/(mol)(K)

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effusion

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the escape of gas molecules through a tiny hole into an evacuated space

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diffusion

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the spread of one substance throughout a space or throughout a second substance

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Graham's Law

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the relative rates at which two gases under identical conditions of temperature and pressure will diffuse vary inversely as the square roots of the molecular masses of the gases:
(rate a) / (rate b) = square root ((molar mass b) / (molar mass a))

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joule

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1 joule = 1 kg x m^2 / s^2

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delta E

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= E final - E initial

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enthalpy

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(H) heat absorbed or released under constant pressure

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delta H

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= H final - H initial = qp
qp equals the heat gained or lost by the system when the process occurs under constant pressure.

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state functions

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properties (like energy) that depend only on an object's current state

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heat capacity

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the amount of heat required to raise its temperature by 1K (or 1*C)

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molar heat capacity

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the heat capacity of 1 mole of a substance

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specific heat capacity

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the heat capacity of 1g of a substance

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specific heat

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(quantity of heat transferred) / ((grams of a substance) x (temperature change))
Cp = q / (m x delta T)

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Hess's Law

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if a reaction is carried out in a series of steps, delta H for the reaction will be equal to the sum of the enthalpy changes for the individual steps

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ΔH° rxn

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∑nΔH°f (products) - ∑mΔH°f (reactants)

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ΔH°f (Standard Molar Enthalpy of Formation of a compound)

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equal to the enthalpy change when one mole of a the compound is formed at a constant pressure of 1 atm and a fixed temperature (usually 25°C) from the elements in their stable states at that pressure and temp.

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ΔHfus (Heat of fusion)

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heat absorbed when a solid melts in units of J/kg or J/mole:
ΔH = (mass) (ΔH fusion)

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ΔHvap (Heat of Vaporization)

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the heat absorbed when a liquid vaporized

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ΔHcomb (Heat of Combustion)

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Energy released when a substance reacts with oxygen to form CO2 and H20

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spontaneous processes

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a change that proceeds on it's own without any outside intervention

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Entropy (symbol S)

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the measure of the degree of disorder of a system

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ΔS

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= S products - S reactants

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Gibbs Free Energy

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a function of the state of a reaction system has been defined to relate both the enthalpy and entropy factors at a given temperature. This combined enthalpy- entropy function is called the free- energy of the system

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ΔG

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change in free energy

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Gibbs-Helmholtz Equation

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ΔG = ΔH - TΔS

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frequency (v)

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the # of complete wavelengths or cycles that pass a given point in 1 second:
= # waves / second = Hertz

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amplitude

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the height of the wave from the origin to its crest or peak
(measures brightness or intensity)

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wavelength (λ)

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the distance between similar points in a wave

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wavespeed

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the speed of light: 3.0 x 10^8 m/sec. symbol C:
C = λv

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Max Planck

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proposed the idea that energy could be released or absorbed by atoms only in "chunks" or "quanta"

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quantum

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(fixed amount) the smallest quantity of energy that can be emitted or absorbed as electromagnetic radiation

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E of a quantum

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E = hv (energy of a single quantum equals a constant times its frequency)

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Planck's Constant (h)

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6.63 x 10^ -34 J x sec

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equation to find E of a quantum

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E = (h)(C) / (λ)

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photoelectric effect

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light or radiant energy shining on a clean, shiny metal surface can emit electrons

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photon

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a unit "packet" of light energy-- behaves like a tiny "particle" of light

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dual nature of light

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possesses both wavelike and particle like properties

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Niels Bohr

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made contributions to understanding line spectra

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monochromatic radiation

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radiation composed of a single wavelength

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continuous spectrum

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continuous merging of colors from certain radiation passed through a prism, refracting its different wavelengths

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line spectrum or discontinuous spectra or bright line spectra

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spectra showing lines of color present of a few wavelengths from specific sources- specific gasses under pressure

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Bohr's Model

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an atomic model proposing only electron orbits of certain radii, corresponding to certain definite energies permitted:
proposed the idea that electrons could absorb/release specific "quanta" of energy to make changes in energy states

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Ground State

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the lowest possible energy state f an atom

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Excited State

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when the electron is in a higher energy state- has excess energy

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Louis De Broglie

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proposed the idea of "matter waves" to describe the wave characteristics of material particles (electrons)

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Werner Heisenberg

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uncertainty principle- it is not possible to accurately determine the position and momentum for a particle as small as an electron simultaneously.

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Erwin Schrödinger

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proposed the idea of wave functions- symbol Ψ

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probability density

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Ψ^2: represents the probability that the electron will be found at that location

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atomic orbital

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a region around the nucleus where an electron with a given energy is likely to be found

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quantum numbers

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(n, l, ml, ms) designations that describe the position and spin of an electron- give a probable location

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1st quantum number- principle quantum number (n)

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assigned values in order of increasing energy: n= 1, 2, 3, 4...
n^2 gives the total # of orbitals in any energy level
2n^2 = maximum number of electrons for level

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2nd quantum number- angular momentum (l)

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designates the shape of the orbital- describes which sublevel the electron is in (s, p, d, or f)
l can be 0 up to (n-1)

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3rd quantum number- the magnetic quantum number (ml)

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assigning ml values helps determine the # of orbitals in a given sublevel- distinguish one orbital from another in possible type/shape/orientation
assigned: +l,... 0... -l

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4th quantum number- spin quantum number (ms)

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indicates the e- spin (clockwise or counter clockwise)
assigned +1/2 or -1/2

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Pauli Exclusion Principle

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No two electrons can have an identical set of 4 Q#'s

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effective nuclear charge

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the net positive charge attracting the electron: Z eff. It is the atomic # - the average # of electrons between the nucleus and electron in question:
Zeff= Z - S

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Aufbau Principle

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electrons enter orbitals of the lowest energy first

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Hund's Rule

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the lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli Exclusion Principle in a particular set of degenerate orbitals.

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Valence Electrons

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electrons in outermost (highest) principle energy level of the atom