Chemistry Flashcards
Principal Quantum Number
n, describes the radial distance of an electron’s orbit from the nucleus
- Electrons orbit at fixed distances from the nucleus but the distance between orbits decreases with distance from the nucleus
- Energy increases with distance from the nucleus
Radioactive decay
Unstable nuclei are said to be radioactive and they undergo a transformation to make thm more stable altering the number and ratio of protons and neutrons or just lowering the energy
Alpha particle
- When a large nucleus wants to become more stable by reducing the number of protons and neutronss, it emits an alpha particle
- Particle consists of 2 protons and 2 neutrons (4/2 alpha) which is equivalent to a helium 4 nucleus so it can be denoted by 4/2 He and the decay reduces the parent’s atomic number by 2 and the mass number by 4
- Alpha particles are emitted with high energy from the parent nucleus but this energy is quickly lost as the particle travels through matter or air so the particles do not travel far and can be stopped by the outer layers of human skin or a piece of paper
Beta decay
- Three types of beta decay (b-, b+, and electron capture) Each type of beta decay involves the conversion of a neuron into a proton through the action of the weak nuclear force
- Beta particles are more dangerous than alpha particles since they are less massive and have more energy and penetrating ability but can be stopped by aluminum foil or centimeter of plastic or gas
B- Decay
When an unstable nucleus contains too many neutrons it may convert a neutron into a proton and an electron which is ejected. The atomic number is 1 greater than parent nucleus but the mass number remains the same. Most common type of beta decay
B+ Decay (positron emission)
When an unstable nucleus contains too few neutrons, it converts a proton into a neutron and a positron which is ejected. The positron is the electron’s antiparticle (identical to an electron except charge is positive).
Atomic number is 1 less than parent but mass number remains same
Electron Capture
Another way for unstable nucleus to increase its neutrons is to capture an electron from the closest electron shell and to use it in conversion of a proton into a neutron
Causes atomic number to be reduced by 1 while mass number remains same
Gamma decay
A nucleus in an excited state can relax to its ground by emitting energy in the form of one or more photons of electromagnetic radiation (photons = gamma photons).
Gamma photons
High frequency and energy and can penetrate most matter effectively
Ejection changes neither mass number nor atomic number
Nuclear Binding Energy
Energy that was released when the individual nucleons (proton and neutrons) were bound together by the strong force to form the nucleus. It’s also equal to the energy that would be required to break up the intact nucleus into tits individual nucleons. The greater the binding energy per nucleon, the more stable the nucleus
Mass defect
When the nucleons bind together to form a nucleus, some mass is converted to energy, so the mass of the combined nucleus is less than the sum of the masses of all its nucleons individually. The difference (delta m) is called the mass defect and its energy equivalent is the nuclear binding effect
delta m = (total mass of separate nucleons) - (mass of nucleus)
Emission Spectrum
Gives an energetic “fingerprint” of that element because it consists of a unique sequence of bright lines that correspond to specific wavelengths and energies. The energies of the photons, or particles of light that are emitted, are related to their frequencies and wavelengths by
Ephoton = hf = h (c/wavelength)
Energy subshell
Comprised of one or more orbitals and is denoted by a letter (s,p,d,f) that describes the shape and energy of the orbital
Aufbau Principle
Electrons occupy the lowest energy orbitals available. Electrons subshells are filled in order or increasing energy
Hund’s Rule
Electrons in the same subshell occupy available orbitals singly, before pairing up
Pauli Exclusion Principle
There can be no more than two electrons in any given orbital
Diamagnetic
An atom that has all of its electrons spin paired; contain an even number of electrons and have all of its occupied sub shells filled. Since all the electrons in a dia are spin-paired, the individual magnetic fields created cancel leaving no net magnetic field. Such an atom will be repelled by an externally produced magnetic field
Paramagnetic
Atom’s electrons are spin-paired so the atoms are attracted into externally produced magnetic fields
Which way do you move in periodic table if atom becomes an anion? cation?
Anion: right
Cation: left
Ionization Energy
The amount of energy necessary to remove the least tightly bound electron from an isolated atom’s ionization energy. As we move from left to right across a period or up a group the ionization energy increases since the valence electrons are more tightly bound. IE2 always greater than IE1
Electron affinity
The energy associated with the addition of an electron to an isolated atom. If energy is released when the electron is added, EA is neg. The halogens have large negative EA values since the addition of an electron would give them the desired octet configuration. So they readily accept an electron to become an anion; the increase in stability causes energy to be released. The noble gases and alkaline earth metals have positive electron affinities because the added electrons begin to fill a new level or sub level and destabilizes the electron configuration. Tend to become more neg as we move to the right or up a group
Electronegativity
Is a measure of an atom’s ability to pull electrons to itself when it forms a covalent bond; the greater the tendency to attract electrons, the greater the value is
Acidity
measure of how well a compound donates protons, accepts electrons, or lowers pH in a chemical system. The ease with which an acid (HX) donates its H+ is related to the stability of the conjugate base. The more electronegative the element bearing the negative charge is the more stable the anion will be. Therefore acidity increases from left to right across a period. The larger the anion, the more the negative charge can be delocalized and stabilized. Acidity increases down a group in the periodic table
Bond dissociation energy
The energy required to break a bond homolytically
Homolytic bond cleavage
One electron of the bond being broken goes to each fragment of the molecule so two radicals form
Heterolytic Bond Cleavage
Both electrons of the electrons of the electron pair that make up the bond end up on the same atom; form a cation and an anion
Relationship between bond order bond strength
The higher the bond order the shorter and stronger the bond- should only be compared for similar bonds (carbon-carbon with c-c)
Bond order
Number of bonds between adjacent atoms so a single bond order of 1
Bond length and bond strength
longer the bond weaker it is; the shorter the bond, the stronger it is
Covalent bond
Formed between atoms when each contributes one or more of its unpaired valence electrons. The electrons are shared by both atoms to help complete both octets.
Coordinate covalent bond
One atom will donate both of the shared electrons in a bond (NH3)
Lewis base
-Donates a par of electrons; Can act as a ligand or a nucleophile
Hydrogen Bond
- Strongest type of intermolecular force between two neutral molecules
- A molecule must have a covalent bond between H and either N, O, F and another molecuele must have a lone pair of e on N, O, F
Vapor Pressure
Pressure exerted by the gaseous phase of a liquid that evaporated from the exposed surface of the liquid. Weaker IMF, the higher vapor pressure and the more easily it evaporates
ALso temperature dependent. Vapor pressure increases with the temp of the substance. INc the average kinetic energy of the particles allows them to overcome the intermolecular forces holding them together and increases the proportion of particles that can move into the gas phase
Ionic Solids
Held together by the electrostatic attractions between cations and anions in a lattice structure. Solid at room temp; The greater the charge, the stronger the force of attraction between the ions
Network Solid
Atoms are connected in a lattice of covalent bonds meaning that all interactions between atoms are covalent bonds. The intermolecular forces are identical to the tintramolecular forces. Very strong and tend to e very hard solids at room temperatures.
Network Solid
Atoms are connected in a lattice of covalent bonds meaning that all interactions between atoms are covalent bonds. The intermolecular forces are identical to the intramolecular forces. Very strong and tend to e very hard solids at room temperatures.
Metallic solids
-Covalently bound lattice of nuclei and their inner shell elections surrounded by a sea or cloud of electrons. At least one valence electron is not bounded to any one particular atom and is free to move throughout the lattice (conduction electrons). Metals excellent conductors and malleable and ductile. Solids at room temp
Molecular solids
Molecules held together by one of three types of intermolecular interactions: hydrogen bonds, dipole-dipole, or LD forces.
- these forces are significantly weaker than ionic, network, or metallic bonds, molecular compounds have much lower melting and boiling points than other types of solids
- Often liquids or gases at room temperature and are more likely to be solids as the strength of IMF increase
Zeroth law of Thermodynamics
If two systems are both in thermal equilibrium with a third system, then the two initial systems are in thermal equilibrium with one another. If systems are in thermal equilibrium with one another, their temperatures must be the same.
First Law of Thermodynamics
The total energy of the universe is constant. Energy must be transformed from one form to another, but it cannot be created or destroyed
Enthalpy
Measure of the heat energy that is released or absorbed when bonds are broken and formed during a reaction that’s run at constant pressure
Exothermic
If the products of a chemical reaction have stronger bonds than the reactants then more energy is released in the making of product bonds than was to put in to break the reactant bonds. Energy is released. Products are in lower energy state than the reactants
Endothermic
The products of a chemical reaction have weaker bonds than the reactants then more energy is put in during the breaking of reactant bonds than is released in the making of product bonds. The products are in a higher energy state than the reactants and enthalpy is positive because heat had to be added to the system from surroudings
Second Law of Thermodynamics
The disorder of the universe increases in a spontaneous process. Nature has a tendency to become increasingly disorganized and another way to state the second law is that all processes tend to run in a direction that leads to maximum disorder
Positive Delta S
If randomness increases or order decreases
Entropy Cases
- LIquids have more entropy than solids
- Gases have more entropy than solids or liquids
- Particle in solution have more entropy than undissolved solids
- Two moles of a substance have more entropy than one mole
- The value of delta S for a reverse reaction has the same magnitude as that of the forward reaction, but with opposite signs
Third Law of Thermodynamics
Defines absolute zero to be a state of zero-entropy. Thermal energy is absent and only the least energetic thermodynamic state is available to system. If only one state is possible, then there is no randomness to the system and S=0. The least thermally energetic state and has the lowest achievable temperature
Gibbs Free Energy
The energy that is free to do useful work from a chemical reaction.
Determining Spontaneity
Delta H Delta S Delta G Rxn is ?
- + - Spontan.
+ + - @high T S
+ @ low T NS
- - + @ high T NS
- @ low T S
+ - + NS
Kinetics v Thermodynamics
Just because a reaction is thermodynamically favorable does not mean that it will be taking place rapidly. Thermo predicts spontaneity and equilibrium not rates
Key property of a physical change
No intramolecular bonds are made or broken; a physical change affects only the intermolecular forces between molecules or atoms
Heat of fusion
The amount of heat that must be absorbed to change a solid into a liquid
Heat of Vaporization
Energy absorbed when a liquid changes to a gas
Formula for the amount of heat accompanying a phase transition
q = n x delta H phase change
Temperature Change vs Phase Change
In between phase changes, matter can absorb or release energy without undergoing transition as an increase or decrease in temperature. When a sample is undergoing a phase change, it absorbs or releases heat without a change in temperature. When a substance absorbs or releases heat, one of two things can happen: either its temperature changes or it will undergo a phase change but not both at the same time
Specific heat
Intrinsic property that tells us how resistant it is to changing its temperature
Temperature Change and Heat Capacity
Temperature change inversely proportional to the substance’s heat capacity. A substance like water with a high specific heat will undergo small change in termperature
Phase Diagrams
Phase of a substance doesn’t depend on the temperature it also depends on the pressure. even at high temperatures, a pressure can be squeezed into the liquid phase if the pressure is high enough
Phase Diagram for Water
An increase in pressure at constant temperatures can favor the liquid phase not the solid phase222222
Volume
1 cm^3 = 1 cc = 1 mL and 1 m^3 = 1000 L
Pressure conversion
1 atm = 760 torr = 760 mm Hg = 101.3 kPa
Deviations from Ideal Gas Behavior
Attractive forces between a particle causes a decrease in pressure if the volume of the container is fixed and accounting for particle volume causes a decrease in free space if the pressure is fixed
Gases behave more ideally at higher temperatures.
Why do higher pressures and lower temperatures cause larger deviations from ideal behavior?
As pressure increases, gas particles become closer to one another. This accentuates the effects of attractive intermolecular forces, causing a decrease in observed pressure (preal < pideal). At low temperatures, IMF become more important. Those gases that behave most ideally have the weakest IMF and the smallest molecular weights and volumes. By maintaining conditions of high temperature and low pressure, the interactions between particles are minimized and particle volume remains insignificant compared to container size helping to favor more ideal behavior for all gases
A container holds methane and sulfur dioxide at a temperature of 227. Let KEm denote the average kinetic energy of the methane and the KEs the average KE of the SO2 molecules. Which describes the relationship between these energies?
KEm= KEs
Both gases have samp temp so avg kinetic energies will be the same
The temperature of neon gas in a glass tube is increased from 10 degrees Celsius to 160. As a result, the avg Kinetic energy of the neon atoms will increase by a factor of?
less than 2
convert C to Kelvins!!!
What happens to carbohydrates in oxidation
broken down to CO2
BOnd between two sugar molecules
GLycosidic linkage; covalent bond formed in a dehydration reaction that requires enzyme catalysis
Fatty acid structure
long unsubstituted alkanes that end in a carboxylic acid. The chain is typically 14 to 18 Carbons long and because they are synthesize two carbons at a time from acetate, only even numbered made in human cells
Triacylglycerols
Storage form of fatty acid is fat (triacylglycerol). Composed of three fatty acids esterified to a gycerol molecule.
lipases
enzymes that hydrolyze fats; triacylglycerols are stored in fat cells as an energy source
Fats as Energy Storage Molecules
Efficient because of packing and energy content
- Packing: hydrophobicity allows fats to pack together more closely than carbs. Carbs carry a great amount of water-of-solvation (water molecules hydrogen bonded to hydroxyl group)
- Energy Conten: Fat molecules store much more energy than carbs
Bond between two phosphates
Anhydride bond
Why are phosphate anhydride bonds able to store so much energy?
- When phosphates are linked together, their negative charges repel each other strongly
- Phosphates have more resonance forms and thus a lower free energy than linked phosphates
- Phosphates has a more favorable interaction with biological solvent than linked phosphates
Intermediate
Substance that is produced in one elementary step and then consumed in a subsequent step. Intermediates will not pile up like dirty dishes they will shuttle back and forth between reactants and products until the slow step takes it forward.
Rate determining step
Slowest step in a process
What is reaction rate determined by?
- How frequently the reactant molecules collide
- The orientation of the colliding molecules
- Their energy
What makes the reaction rate faster
- Lower the activation energy, faster the reaction
- The greater the concentration of the reactants, the faster the reaction rate (favorable collisions are more likely as the concentrations of reactant molecules increase)
- The higher the temperature of the reaction mixture, the faster the reaction rate (more molecules have sufficient energy to overcome the activation energy barrier and molecules collide at a higher frequency so the reaction can proceed at a faster rate)
Catalyst
Provides reactants with a different route, usually a shortcut to get to products. Will make the reaction go faster by either speeding up the rate determining step or providing an optimized route to products b lowering activation energy and energy of the highest energy transition state
Difference between reactants and catalysts
Catalyst remains unchanged at the end of a reaction
Phase Solubility Rules
- The solubility of solids in liquids tends to increase with increasing temperature
- Solubility of gases in liquids tends to decrease with increasing temperature
- Solubility of gases in liquids tends to increase with increasing pressure
Salt solubility Rules
- All Group 1 (Li+, Na+, K+ Rb+, Cs+) and ammonium salts are soluble
- All nitrate (NO3-) Perchlorate (ClO4-), and acetate (C2H3)2-) salts are soluble
- All silver (Ag+) lead (Pb2+/Pb4+) and mercury (Hg2 2+/Hg2+) salts are insoluble except for their nutrates, perchlorates, and acetates
Ka and strength of acid
The larger the Ka value, the stronger the acid. The smaller the Ka value the weaker the acid
Conjugate acid and base tendencies
- The conjugate acid of a strong base has no acidic properties in water. Ex: the conjugate acid of LiOH is Li+ which does not act as an acid in water
- The conjugate acid of a weak base is a weak acid (the weaker the base, the stronger the conjugate acid) Ex: the conjugate acid of NH3 is NH4+ which is a weak acid
Neutralization Reactions
When an acid and a base are combined they will react in a neutralization reaction. Oftentimes this reaction will produce a salt and water.
Henderson Hasselbalch Equation for acid
pH = pKa + log (conjugate base/ weak acid)
Henderson-Hasselbalch Equation (for base)
pOH= pKb + log (conjugate acid/weak base)
Indicator
Weak acid that undergoes a color change when it’s converted to its conjugate base
Acid-base equivalence points
- For a weak acid (titrated with a strong base) the equivalence point will occur at a pH > 7
- For a weak base (titrated with a strong acid), the equivalence point will occur at a pH < 7
- For a strong acid (titrated with a strong base) or for a strong base (titrated with a strong acid), the equivalence point will occur at pH=7
