Chemistry 1 Final Exam Review - Definition Notes PDF

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three states of matter solid, liquid, gas

solid definite shape and volume

liquid matter that has a definite volume but no definite shape

gas a state of matter with no definite shape or volume

conservation of energy a fundamental principle stating energy cannot be created nor destroyed but only changed from one form to another

conservation of matter/mass matter is neither created nor destroyed

dalton's atomic theory all matter is composed of extremely small particles called atoms

accuracy how close a measurement is to the true value

precision how close a group of measurements are to each other

physical change

a change from one state (solid or liquid or gas) to another without a change in chemical composition

chemical change A change that occurs when one or more substances change into entirely new substances with different properties

intensive properties a property that does not depend on the amount of matter present, such as pressure, temperature, or density

extensive properties a physical property, such as mass, length, and volume, that is dependent upon the amount of substance present

pure substances a sample of matter, either a single element or a single compound, that has a definite chemical and physical properties

mixtures two or more substances that are not chemically combined

homogeneous mixture a mixture that is uniform throughout. A solution.

heterogeneous mixture A mixture in which the composition is not uniform throughout

Celsius to Fahrenheit F= 1.8 (C) +32

metric prefixes tera, giga, mega, kilo, hecto, deka, base unit, deci, centi, milli, micro, nano, pico, femto, atto

tera 10¹²

giga 10^9

mega 10^6

kilo 10^3

hecto 10^2

deka 10^1

deci 10^-1

centi 10^-2

milli 10^-3, or 1/1000

micro 10^-6

nano 10^-9

pico 10^-12

density mass/volume

law of definite proportions the law that states that a chemical compound always contains the same elements in exactly the same proportions by weight or mass

thomson discovered the electron, Plum Pudding Model

milikan oil-drop experiment; found mass and charge of an electron

Rutherford gold foil experiment - atoms are made up if mostly empty space with a small dense nucleus in the center, discovered protons

common polyatomic ions No3(-) / Nitrate, SO4(2-) / Sulfate, OH(-) / Hydroxide, CN(-)/ Cyanide, PO4(3-)/ Phosphate, CO3(2-)/ Carbonate, HCO3(1-)/ Hydrogen Carbonate

ionic compounds compounds composed of cations and anions

binary molecular compounds combination of two nonmetals

common acids HCl = hydrochloric acid HNO3 = nitric acid H2CO3 = carbonic acid H2SO4 = sulfuric acid HC2H3O2 = acetic acid H3PO4 = phosphoric acid a molecule that can release protons(H+) into a solution(proton donor)

nomenclature a system of scientific names

limiting reagent the reactant used up first in a reaction

actual yield the amount of product actually produced by a chemical reaction

theoretical yield in a chemical reaction, the maximum amount of product that can be produced from a given amount of reactant

empirical formula

a formula showing the lowest whole number ratio of atoms in a compound. ex CH3

molecular formula ex. C2H6

molarity (moles of solute) / (liters of solution)

solution volume M1V1 = M2V2

SOLUTION X IS 10 PERCENT ALCOHOL BY VOLUME, AND SOLUTION Y IS 30 PERCENT ALCOHOL BY VOLUME. HOW MANY MILLILITERS OF SOLUTION Y MUST BE ADDED TO 200 MILLILITERS OF SOLUTION X TO CREATE A SOLUTION THAT IS 25 PERCENT ALCOHOL BY VOLUME? draw it out. 10% of 200 ml is 20 ml. So 20 ml alchohol and 180 ml water. then solution y is .3(30%) k alcohol and .7k ml water. total solution y is k ml. New solution is 180 + .7k ml and 20 + .3k ml and total volume is 200 + k ml.

so you have an equation: 1/4 (25%) = 20 + .3k / 200 + k. Solve for K to get 600.

strong electrolytes Substance that dissociates completely into its ions when added to water and conducts electricity well.

weak electrolytes produce relatively few ions when dissolved in water, weak acids, weak bases

nonelectrolytes substances that form no ions in water and cannot conduct electricity

qualitative

observations that describe, sight, sounds, smells, and textures., an observation made without measurement, descriptive data

quantitative Data that is measurable; observations based on numbers

equilibrium when the concentration of a solute is the same throughout a solution

dynamic equilibrium result of diffusion where there is continuous movement of particles but no overall change in concentration

Arrhenius concept a concept postulating that acids produce hydrogen ions in aqueous solution, while bases produce hydroxide ions

Bronsted- Lowry concept 1. an acid is a proton (H+) donor 2. a base is a proton (H+) acceptor

strong acids HCl, HBr, HI, HNO3, HClO4, H2SO4

weak acids don't completely dissociate in water HF, HNO2, HClO2. CH3COOH

oxidation-reduction reaction any chemical change in which one species is oxidized (loses electrons) and another species is reduced (gains electrons); also called redox reaction

precipitation reaction A reaction in which two solutions react to form an insoluble solid, which separates from the solution.

electromagnetic radiation form of energy that exhibits wavelike behavior as it travels through space

atomic spectra Spectrum of certain absorbed wavelengths of light corresponding to an atom's spectrum of emitted frequencies of light. Unique to each element. AAS can be used to indentify an element.

the bohr theory Electrons can only occupy certain orbits or shells in an atom and each orbit represents a definite energy for the electrons in it

quantum atomic structure - the electrons in an atom are grouped around the nucleus into shells - the farther a shell is from the nucleus, the larger it is, the more electrons it can hold, and the higher the energies of those electrons (as you move further away from the nucleus, the attraction to the nucleus decreases)

atomic orbitals the regions around the nucleus within which the electrons have the highest probability of being found

diamagnetism all electrons paired (ends of the blocks)

paramagnetism the presence of unpaired electrons

frequency (nu (ν)) The number of cycles a wave undergoes per second, expressed in units of 1/second

wavelength (lamda(λ) The number of cycles a wave undergoes per second, expressed in units of 1/second

speed of light (c) 3.00 x 10^8 m/s; c= (lamda)(nu)

plank's constant (h) h=6.626 x 10^-34

energy of radiation ΔE = hc/λ = hν, E=nhν

rydberg equation 1/λ=R (1/(n1)^2-1/(n2)^2)

effective nuclear charge stays the same down a group, but increases right

ionization energy periodicity increases up and to the right of periodic table

electron affinity periodicity increases up and to the right

atomic radii periodicity increases down and to the lef

electronegativity periodicity increase to top right, tendency of an atom to attract an electron away from another

ionic bond Formed when one or more electrons are transferred from one atom to another

lattice energy the enthalpy change that accompanies the reverse of this equation—1 mol of ionic solid separating into gaseous ions.

covalent bond bond formed by the sharing of electrons between atoms

resonance refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure.

bond polarity determined using difference in electronegativity, a measure of how equally or unequally the electrons in any covalent bond are shared

molecular polarity uneven distribution of molecular charge

VSEPR theory a theory that predicts some molecular shapes based on the idea that pairs of valence electrons surrounding an atom repel each other.

valence bond theory bonds are formed by sharing of electrons from overlapping atomic orbitals

hybrid orbitals Orbitals of equal energy produced by the combination of two or more orbitals on the same atom.

pi bonding results from side-to-side or parallel overlap of two atomic orbitals

isomers compounds with the same formula but different structure

structural isomer Compounds that have the same molecular formula but differ in the covalent arrangements of their atoms.

optical isomer molecules that are mirror images of each other

thermodynamics the study of energy transformations that occur in a collection of matter

laws of thermodynamics 1) energy can not be created or destroyed, only converted from one form to another, 2) each time you convert one form of energy to another, some energy is converted to a non-usable form (more energy efficient to consume plants because they exist very close to the initial source of energy)

enthalpy total energy in a system; the amount of energy in a system capable of doing mechanical work

entropy A measure of disorder, or randomness

gibbs free energy ∆G=∆H-T∆S; , Changes in it, ΔG, are useful in indicating the conditions under which a chemical reaction will occur. The equation is ΔG = ΔH - TΔS, where ΔH = change in enthalpy and ΔS = change in entropy. If ΔG is negative, the reaction will proceed spontaneously to equilibrium.

ΔG = ΔH - TΔS where H is enthalpy(heat content, bond energy),T is temperature,and S is entropy,G, the Gibbs free energy

Calorimetry precise measurement of the heat flow into or out of a system for chemical and physical processes

average bond energies ΔH= Σreactants - Σproducts

heats of reaction to measure, we must look at the heat required to raise the temperature of a substance

hess's law - Can find ΔH of any reaction

- Arrange series of rxns that give ΔH rxn - Reverse sign if you flip - Multiply ΔH rxn if you change number of moles to get something to cancel

Kinetic-Molecular Theory describes the behavior of matter in terms of particles in motion

Avogadros law

the volume is directly proportional to the number of moles of gas, at constant temp and pressure. V=n

daltons law Total pressure is equal to the sum off all partial pressures.

ideal gas law PV = nRT;

grahams law RateA/RateB=√Mb/√Ma; , the gas diffusion rate is inversely proportional to the square root of its molar mass

law of combining gas volumes At the same temperature and pressure, volumes of gases combine or decompose in small, whole number ratios; reactions must be balanced

Interparticle forces forces between particles that make up a substance

intermolecular forces attractive forces between molecules

ion-ion an atom or group of bonded atoms that has a positive or negative charge

london forces the intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles

dipole-dipole

Created when atoms are joined by a polar covalent bond. The positive end of a dipole in one compound will be attracted to the negative dipole in another compound creating weak attraction between the two compounds.

hydrogen bonding Very strong intermolecular force where a H covalently bound to an N, O, or F is attracted to another N, O, or F

ion-dipole type of van der Waals force between an ion and the partial charge on the end of a polar molecule

ion-induced dipole Exists between ions and non-polar compounds. Stronger than Dipole-Induced Dipole because Ions are truly positive, not partially.

vaporization the change from a liquid to a gas

boiling point the temperature at which a liquid becomes a gas

freezing point The temperature at which a liquid changes into a solid

sublimation point solid to gas

phase changes substances change phases by adding or subtracting energy; Phase Changes that require energy includes melting, sublimation and vaporization; phase changes that release energy include condensation, deposition and freezing.

solution a mixture in which one substance is dissolved in another

solute the substance that is dissolved

solvent The dissolver. ex: water.

saturated solution a solution in which no more solute can dissolve

unsaturated solution a solution in which more solute can be dissolved

supersaturated solution a solution that holds more dissolved solute than is required to reach equilibrium at a given temperature

solvation the process of dissolving

lattice energy the energy required to separate one mole of the ions of an ionic compound

Le Chatelier's principle States that if a stress is applied to a system at equilibrium, the system shifs in the direction that relieves the stress.

endothermic dissolving Dissolving that absorbs heat energy and ammonium nitrate added to water (cold pack)

exothermic dissolving dissolving that releases heat energy and sulfuric acid added to water

effect of pressure on solubility when a gas is part of a solution, the higher the pressure, the more soluble the gas

concentration calculation Percent Solute (g solute/g solute+g soln) or Molarity (moles/Liter) exothermic gives off heat to the surrounings

endothermic absorbs heat from the surroundigns

List 4 State functions 1 Volume 2 Temperature 3 Pressure 4 Composition

change of state function depends on: initial and final states of the system

change of state function is independent of: path which energy is accomplished

Is energy a state function? yes

Is work and heat a state function? no

Is enthalpy a state function? yes

Constant-volume and constant-pressure calorimeters are used to measure what? heat changes that occur in physical/chemical processes

What is Hess's law? the overall enthalpy change in a reaction is equal to the sum of enthalpy changes for individual steps of the overall reaction

How can the standard enthalpy of a reaction be calculated? From the number of moles and standard enthalpies of formation for reactants and products

equation for standard enthalpy of reaction ∆H˚rxn=∑n∆H˚f(products)-∑m∆H˚f(reactants)

equation for calculating heat change in terms of specific heat q=ms∆t

equation for calculating heat change in terms of heat capacity q=C∆t

equation for calculating work done in gas expansion or gas compression w=-P∆V

equation for density Mass/Volume mass of an obj dived by its volume Most used for (solids and liquids): (g/cm^3) or (g/mL) Most used for (gases): (g/L)

define & give exs. of physical properties measured and observed without change in composition or identity of a substance ex: color melting point boiling point

define & give exs. of chemical properties must cary out chemical change ex. burning

define & give exs. extensive property measured value depends on how much matter is being considered. Additive! ex. mass volume length

define & give exs. intensive property does not depend on how much matter is considered. Not additive! ex. density

temperature

SI base unit for amount of substance mole (mol)

SI base unit for temperature kelvin (K)

SI base unit for mass kilogram (kg)

SI base unit for time second (s)

SI base unit for length meter (m)

convert 1 liter to mL, cm^3, and dm^3 volume occupied by ONE CUBIC DECIMETER 1 L= 1 dm^3 1,000 cm^3 1,000 mL

convert mL to volume 1mL=1cm^3

How many sig figs in x or division? +/-? * or / =least number of sig figs

+ or - = smallest number of decimal places

convert 5.2 L to m^3 1L = 1dm^3

Precipitation Reaction: A reaction in which an insoluble substance forms and separates from the solution. The solid is known as a precipitate and the solution above the precipitate is called the supernatant.

Know the different gas laws PV = nRT V1/T1 = V2/T2 P1*V1 = P2*T2 V₁/n₁ = V₂/n₂

Know how to find percent error: {[Absolute Value of (Experimental Value - Theoretical Value)] / (Theoretical Value)} * 100 %

Know what is precision and accuracy o Precision is a measure of how close several trials of an experiment are to each other. o Accuracy is how close an experimental value is to the theoretical value.

Know about electrolytes An electrolyte is a substance that ionizes when dissolved in suitable ionizing solvents such as water. This includes most soluble salts, acids, and bases. Some gases, such as hydrogen chloride, under conditions of high temperature or low pressure can also function as electrolytes.

Know how to read a buret/graduated cylinder: For a buret and a graduated cylinder read at the bottom of the meniscus at eye level, to four significant figures.

Know key information about the SPARK system ...

Know key information, and 2 purposes from the Percentage of Oxygen experiment ...

Reagent A reagent is a solution that has the potential to react with one or more ions in a solution ofen producing a chemical reaction.

List the 4 indications that a chemical reaction has occurred? 1. Formation of Precipitate? 2. Change in color/texture? 3. Evolution of a gas. 4. Re-dissolving a precipitate through the addition of heat.

How do you differentiate between a color change in a solution and the formation of a precipitate. By centrifuging a sample, the precipitate will separate from supernatant, giving the indication of whether there was an actual color change or simply a precipitate formed or both?

Avogadro's Law ** V₁/n₁ = V₂/n₂** The equation shows that, as the number of moles of gas increases, the volume of the gas also increases in proportion.

Boyle's Law ***P₁ x V₁ = P₂ x V₂*** The equation states that product of pressure and volume is a constant for a given mass of confined gas as long as the temperature is constant.

Gay-Lussac's Law

***P₁/T₁=P₂/T₂*** & *** P₁T₂ = P₂T₁. If a gas's temperature increases, then so does its pressure if the mass and volume of the gas are held constant.

Charle's Law ***V₁/T₁=V₂/T₂*** & *** V₁T₂ = V₂T₁. This law describes how a gas expands as the temperature increases; conversely, a decrease in temperature will lead to a decrease in volume.

Know key information and purposes about Conductivity (including the graphs) Purpose: 1. Determine whether substances are strong, weak, or non-electrolytes based on their conductivity. 2. Perform a conductometric titration to observe changes in conductivity during a chemical reaction.

Titration Is a common lab technique used to determine the concentration of a particular reactant in a chemical reaction.

How were conductometric titrations used in the conductivity lab? Conductometric Titrations in the conductivity lab will be done using two different aqueous solutions, and the conductiv...