General Chemistry Flashcards

Molecular Weight

• Also molecular mass, is the mass of a single molecule of that compound, generally given in atomic mass units.

• Determined with periodic table

Molecular Formula

• Symbols for the elements of the compound, with a subscript indicating the number of atoms of that element if that number is greater than one.

Percent Mass

• The empirical formula can be determined by dividing the percent mass of each element by the atomic mass of that element and finding the ratio of the results.

• Mass of solute/ mass of solution x 100

Avogadro’s Number

• Basis definition of the mole.

• Converting between micro and macroscopic units, moles and molecules, or grams and atomic mass units.

• 6.022 × 10^-23

Definition of Density

• Amount of mass in a substance or object per unit volume.

• Rho, p.

Oxidation Number

• The measurement of its effective charge within the molecule.

• Number of elections that have been removed from or added to the atom in question.

• The more electronegative atom in a compound usually has a lower oxidation state.

Common Oxidizing and Reducing Agents

• Oxidizing agents: oxygen, chlorine, and halogens, hydrogen peroxides, and sulfuric acid.

• Reducing agents: hydrogen gas, sodium, alkali metals, iron, carbon monoxide.

Limiting Reactants

• The reactant the runs out first when the reaction takes place, the others are called excess reactants.

• Balance equation, convert to moles, divide by coefficient, and compare values. Lowest is the limiting reactant.

Theoretical Yields

• The amount that will be produced if the maximum amount of the reactants reacted.

• Determined by identifying the limiting reactant and determining how much of the product it would produce based on the ratios of the coefficients.

Absolute Temperature and the Advantages of the Kelvin Scale

• 0K is absolute zero of temperature.

• C + 273.15 = K

  (32°F − 32) × 5/9 + 273.15 = 273.15K

Pressure, Simple Mercury Barometer

• Due to Pascal’s Law the pressure from the surrounding air will press on the mercury in the open container throughout the mercury.

• At equilibrium the hydrostatic pressure due to the mercury column will be equal to the atmospheric pressure of the surroundings, so the atmospheric pressure can be calculated by the height, h, of the column of mercury.

• Patm= rhoHg(gh)

Ideal Gas

• Made of particles with the following characteristics

  • Volume is negligible compared to the volume of the container

  • Do not interact except during collisions

  • All collisions are perfectly elastic (no KE is lost)

  • Between collisions, the particles travel at constant speed in straight lines.

Ideal Gas Law

• An equation that relates the pressure, volume, temperature, and number of moles or molecules of an ideal gas.

• PV=nRT

• Combines Boyle’s law, Charle’s law, and Avogadro’s law.

• Boyle’s law: P1V1= P2V2, pressure and volume of a gas have an inverse relationship when temperature is held constant.

• Charles’ law: V1/T1 = V2/T2, volume of a gas is directly proportional to its absolute temperature when pressure is held constant.

• Avogadro’s law: volume of a gas is directly proportional to the number of moles of gas present, provided the temperature and pressure stay constant.

Heat Capacity

• Measurement of the amount of heat that must be transferred to an object to raise its temperature by a specific amount.

• C= Q/deltaT

• Q= heat

• DeltaT= change in temperature

Boltzmann’s Constant

• k or kB, is a constant that is equal to 1.38 × 10^-23 J/K

Dalton’s Law

• The sum of the partial pressure of all the constituents of a particular mixture of ideal gases is equal to the total pressure of the gas.

Intermolecular Forces: Hydrogen Bonding

• Between two polar molecules arising between a hydrogen atom with a partial positive charge in one molecule is attracted to an atom in another molecule with a partial negative charge.

• Cause molecules to stick together, why they have a higher boiling point and viscosities.

Intermolecular Forces: Van der Waal’s Forces

• London dispersion forces are weak forces that arise between molecules because of temporary polarization- even if a molecule is nonpolar, temporary dipole moment at a given time bc the electrons happen to be at that time, distributed more toward one side of the nucleus, and two molecules with such temporary dipole moments may be attracted to each other similarly to a dipole-dipole interaction.

• London dispersion forces and dipole-dipole interactions are collectively called Van der Waals forces.

Phase Diagram

• Shows the phase a particular compound or element is in at a different pressure and temperature.

Heat of Fusion and Heat of Vaporization

• Heat of fusion, Lf, is the amount of heat released when a quantity of compound changes from liquid to solid state. Or solid to liquid.

• Heat of vaporization, Lv, is the amount of heat released when a quantity of compound changes from gas to liquid state, or liquid to gas.

Polar and Nonpolar Covalent Bonds

• Polar covalent bonds result when electrons are shared unequally between atoms.

• Nonpolar covalent bonds are shared equally.

• Unequal sharing is due to the differences in electronegativities of two atoms.

• Partial charges develop because of this.

Intensive Properties

• Intensive do not depend on the amount of matter or quantity of the sample, meaning they do not change if the sample size is increased or decreased.

• Color, hardness, melting point, boiling point, density, ductility, malleability, specific heat, temperature, concentration, and magnetization.

Extensive Properties

• Do depend on the amount of matter or quantity of a sample.

• If sample size is increased, property increases.

• Volume, mass, weight, energy, entropy, number of moles, and electrical charge.

Mixtures

• Two or more substances that are combined in various proportions.

• Homogenous: uniform composition

• Heterogenous: not uniform

Solution, Solute, and Solvent

• Solution is a homogenous mixture of substances that cannot be separated by filtration or centrifugation.

• Solutions are made by dissolving one or more solutes into a solvent.

Colloid

• Heterogenous mixture in which small particles are suspended, but not dissolved in a liquid. Can be separated by centrifugation.

• Milk.

Freezing Point Depression

• Colligative property of solutions that depends only on the number of particles in solution, not on the identity of those particles.

• Adding a nonvolatile solute to a solution will lower the freezing point of that solution.

• This decrease in temperature is known as freezing point depression.

• Vapor pressure lowers, causing an decrease in the freezing point.

Solubility Product Constant

• Ksp, is a constant that relates to the equilibrium conditions for the concentrations of ions in a slightly soluble ionic compound.

• Ksp= [A^n+]^m[C^m-]^n

Common-Ion Effect

• Slightly soluble ionic compounds become less soluble in a solution that contains one of the same ions as the compound.

• Product of the concentrations of the ions is constant.

Osmosis

• Diffusion through a semipermeable membrane.

• Net flow of solvent from a solution with a lower concentration to a solution with a higher concentration until equilibrium is reached. (high water to low water)

Ionization of Water and pH

• pH is the measurement of the conc of H ions in an aqueous solution and therefore of the solution’s acidity.

• -log10 of the ion concentration.

Strong and Weak Acids and Bases

• Strength of an acid or base refers to how readily it donates or accepts a proton.

• A strong acid dissociates completely in water, practically every molecule loses a proton (increase negativity).

• A weak acid partially dissociates.

• Strong base, practically every molecule accepts a proton- because the base dissociates into a cation and a proton-accepting hydroxide ion.

• Weak base accepts a proton but less readily.

Buffers

• A solution that resists change in pH.

• Contains either a weak acid and its conjugate base or a weak base and its conjugate acid.

• When a strong base is added to a solution containing a weak acid, the weak acid gives up a proton to the base, changing into its conjugate acid.

Indications (Acids and Bases)

• A substance that when added to a solution, changes color depending on the pH of the solution.

• Litmus red in acidic and blue in basic.

Neutralization (Acids and Bases)

• Refers to the reaction of an acid and a base to form a salt.

• Arrhenius bases - those that include a hydroxide (OH-) ion, the reaction produces water.

• Acid and base are neutralized when no excess acid or base remains, but resulting solution may be acidic, if it results from a strong acid and a weak base, OR basic if it results from a weak acid and strong base.

Redox Titration

• Small amounts of a substance of known concentration are added to a fixed volume of a substance of unknown concentration (analyte).

• In redox titration: titrant is the reducing agent, and analyte is an oxidizing agent.

Reaction Rate

• Measurement of what quantity of the reactants reacts per unit time.

• Measure in mol/ L(s) so the reaction rate, r, means that in each liter of the substance, r moles of the reactants reach each second.

Rate Constant

• r=k[R1]^x1[R2]^x2[R3]^x3…

• 298K

Activation Energy

• Minimum energy required for the reaction to proceed.

Arrhenius Equation

• Relationship between reaction rate and temperature.

• k= Ae Ea/RT

• A is a constant, Ea is activation energy, R is ideal gas constant (8.314 J/mol x K), T is temperature in K.

Catalysts

• A compound that lowers the activation energy of a reaction and thus facilitates that reaction taking place without being expended in the reaction.

• Increases rate of reaction.

Equilibrium Constant

• Large equilibrium constant implies that the products will predominate at equilibrium, small says reactants will.

Equilibrium in Reversible Chemical Reactions: Application of Le Chatelier’s Principle

• A change in some property of a system in equilibrium will produce a shift in the equilibrium that counteracts the change.

• If concentration of one of the reactants increased, the reaction will tend to proceed toward the right, using up some of the reactants and forming more of the product.

Zeroth Law of Thermodynamics

• Two bodies both in thermal equilibrium with a third body are in thermal equilibrium with each other.

• Heat may flow through them but no net head flow between the objects.

First Law of Thermodynamics

• Total energy of an isolated thermodynamic system is constant.

• Non-isolated, energy may be added or taken away from the system in the form of work or heat, but this means the system’s internal energy will change by the same amount.

Second Law of Thermodynamics

• Total entropy of an isolated system will never decrease.

• Must increase or stay the same.

Hess’s Law

• Hess’s Law of Heat Summation states that if a chemical reaction can be broken down into smaller steps, then the total heat of the reaction is equal to the sum of the heats of reaction of each step.

Measurement of Heat Changes (Calorimetry)

• Process of measuring heat transferred to or from a system.

Heat Transfer: Conduction

• Transfer of heat between two objects in contact.

• Pot on stove.

Heat Transfer: Convection

• Transfer of heat within a fluid by the movement of the fluid.

• Air or ocean currents, hot and cold.

Heat Transfer: Radiation

• Transfer of heat by the emission of electromagnetic radiation.

• Very slow method.

• Sun and earth.

Isotopes

• An isotope of an element is an atom of that element with a particular number of neutrons.

• A specific isotope is named in terms of the atomic number and mass number.

Half-Life of Radioactive Material

• The time it takes for half of the atoms in a sample to undergo radioactive decay.

Bohr Atom

• An atomic model developed by Niels Henrik David Bohr.

• Electrons orbited the nucleus in circular paths, but only in specific, discrete stationary orbits, and were only able to absorb or emit quantities of energy that would transfer them from one stationary orbit to another.

Ground State of An Atom

• The state with the lowest energy.

• All of the electrons are in the lowest-energy states possible, subject to the Pauli Exclusion Principle and the maximum number of electrons in an orbital.

Absorption Line Spectra

• Produced by the absorption by a gas or other transparent medium of photons of particular frequencies from a continuous spectrum.

• When an atom or molecule absorbs a photon, the absorbed energy puts the particle into an excited state.

• This places a constraint on the frequencies of photons that can be absorbed; only the photons that have the exact amount of energy to raise the particle to an excited state can be absorbed.

Fluorescence

• Phenomenon where certain materials absorb light at one wavelength and emit light at a wavelength of a lower frequency.

Pauli Exclusion Principle

• For a certain class of particles called fermions, two such particles in the same system cannot be the same state at the same time.

• Photons, for example, are not fermions.

Paramagnetism and Diamagnetism

• Ferromagnetism is the type of magnetism present in familiar permanent magnets

• Only a few materials are capable of this, including cobalt, nickel, and iron.

• Adjacent atoms tend to align so that their electrons spins are in the same direction, causing their individual magnetic moments to build into a significant magnetic field.

Heisenberg Uncertainty Principle

• States that the position and momentum of an object cannot simultaneously be known with arbitrarily high precision.

• (change in x)(change in p) > or equal to h/4pi

• change in x = uncertainty in position

• change in p= uncertainty in momentum

• h= Planck’s constant, 6.626 × 10^-34 J/s

Effective Nuclear Charge

• Zeff, is the effective net nuclear charge experienced on electrons given the presence of other electrons in the atom.

• Since each electron is both attracted to the nucleus and repelled by the other electrons present, the electrical force from the other electrons effectively cancels much of the force from a like number of protons.

Photoelectric Effect

• Light shined on certain materials causes them to emit electrons.

• Assuming that the energy of light knocks electrons from the atom, several observed aspects of the phenomenon isnt explained by this.

• The photoelectric effect does not occur when the light’s frequency is below a certain threshold, regardless of the intensity of the light.

Albert Einstein and the Photoelectric Effect

• Explained that the effect can be explained if the light came in distinct quanta, now called photons, with an energy dependent on the light’s frequency.

• The number of photons determined the intensity, but only if individual photons had sufficient energy would they be able to remove electrons from atoms.

Lewis Electron Dot Formula

• Each atom is represented by its chemical symbol; a line drawn between atoms represents a bond and a dot represents an unbonded electron.

Dipole Moment

• Dipole moment of a molecule is a measurement of the polarity of the molecule- the difference in charge between the ends.

• The magnitude of the dipole moment of an atomic bond in a molecule is equal to the charge difference between the bonded atoms times the length of the bond.

• Its a vector quantity, pointing from the positive toward the negative charge.

Sigma and Pi bonds

• Sigma bond is a bond formed by two atomic orbitals of similar orientation overlapping end to end.

• Overlapping s-orbitals always form sigma bonds because they are spherically symmetrical and their orientation is irrelevant, but p-orbitals and higher orbitals can also form sigma bonds if they overlap along their axes of symmetry.

• Pi bond forms when two atomic orbitals overlap along their sides, such as two p-orbitals that are displaced along a line perpendicular to their axes of symmetry.

Delocalized Electron

• An electron that may pertain to a particular molecule or ion but is not associated with any one atom or bond.

• The presence of delocalized electrons tends to lead to greater electrical and thermal conductivity, because of their mobility.

• Delocalized electrons are common in metals, which form metallic bonds in which valence electrons move freely among the atoms in a sort of “electron sea” rather than remaining near one or two atoms.

Mass Spectrometer

• A device used to separate molecules by mass.

• Relies on particles being charged, uncharged molecules can be separated by a mass spec if they are first ionized.

Alkali Metals

• Occupy the leftmost column of the periodic table, Group I.

• Lithium, sodium, potassium, and so on. Hydrogen, is not among one of these.

Halogens

• Occupy the second column from the right.

• Fluorine, chlorine, bromine, end in -ine.

Noble Gases

• Rightmost column, end in -on. Helium is a noble gas.

• Boron, carbon, and silicon end in -on, but are not noble gases.

Transition Metals

• Groups III through XII, starting at the column with Sc on the left and ending with Zn on the right.

• Sometimes the lanthanides and actinides, usually shown in separate rows below the rest of the table, are included and called inner transition metals.

• Elements that are not transition metals are called main-group elements or representative elements.

Metals and Nonmetals

• Metals are characterized in their solid state by malleability (reshaped by pressure) and ductility (stretched into thin wires).

• Nonmetals are brittle, cracking under stress, rather than reshaping.

• Metals are good conductors of both heat and electricity, from the fact that their outer electrons are relatively free to move between atoms.

Oxygen Group

• Third column from the right. Sometimes called chalcogens.

• Sulfur, selenium, tellurium, and polonium.

Metalloids

• Boron, silicon, germanium, arsenic, antimony, tellurium, polonium, astatine.

• All of the elements above and to the right are nonmetals; all of the metals below and to the left are metals, except for hydrogen.

Valence Electrons

• The electrons on the outermost shell of an atom.

• The number of valence electrons in an atom determines its chemical properties.

• Elements in the same group have the same amount of valence electrons.

Ionization Energy

• The amount of energy needed to ionize an atom of that element- to remove an electron from another atom.

• The energy required to remove one electron from a neutral atom is the first ionization energy, the energy required to remove another electron from an atom that has already lost an electron (a singly charged positive ion) is the second ionization energy, and so on.

Electron Affinity

• Measurement of the charge in energy when an atom or molecule in the gaseous state gains an electron.

• Equivalently, the electron affinity can be defined as the amount of energy required to remove an electron from a singly charged negative ion.

• The larger the electron affinity, the more stable the negative ion.

Electronegativity

• Tendency of an atom to attract electrons.

• Need to define a reference point, by convention a hydrogen has electronegativity of 2.20 and this is the point for all other atoms.

Radioactive Decay and Half-Life

• Set of processes that allow unstable atomic nuclei, or nuclides, to emit subatomic particles or radiation.

• Radioactive decay occurs at an exponential rate.

• A= Aoe^kt

• Ao is the amount of material present at the time t=0, and k is the activity of the material.

Gamma Rays

• Gamma ray (y) is electromagnetic radiation that is produced from radioactive decay or other nuclear processes, such as electron annihilation in which an electron and positron collide.

• No physical difference exists between X-rays and gamma rays that have the same energy, they are set apart by their origin.

• Gamma rays are high-energy electromagnetic radiation that result from nuclear transitions, while X-rays are high-energy radiation caused by energy transitions from electrons that are accelerating.

Nuclear Binding Energy

• Energy that would be required to disassemble the nucleus into its constituent nucleons.

• Calculated by determining the difference in mass between the nucleus and the sum of masses of its constituent particles.

• Mass is converted to energy by E= mc²

• Where c= the speed of light in a vacuum.

Fusion

• Requires a large input of energy to overcome the repulsive force between nuclei, but after transformation it produces an even greater amount of energy.

• Fission is chosen over fusion, since it is easier to control.

Fission

• Initiated by bombarding atoms with a stream of neutrons, destabilizing the atoms and causing a fission into smaller nuclei and emission of numerous radiation particles, including more neutrons, thus perpetuating the reaction.

• Ideal for generating power.