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chem exam 3 terms
| Question | Answer |
|---|---|
| electromagnetic spectrum | range of all types of electromagnetic radiation visible light makes up a small portion of it trend: most energy at smallest wavelengths (10^-12), and least energy at larger wavelengths (10^3) |
| how to characterize a wave | by its wavelength and frequency |
| wavelength | lambda distance between any 2 identical points on adjacent waves unit: usually nanometers (x10^-9), but any length can be used |
| wave | oscillation or periodic movement that can transport energy from one point in space to another |
| frequency | measured in hertz (cycles per second) number of successive wavelengths that pass a given point in a unit time |
| amplitude | 1/2 the distance between peaks and troughs |
| constructive interference | when two waves meet and they combine (to make bigger wave) |
| destructive interference | when 2 waves meet and they cancel each other out think noise cancelling headphone function |
| wave-particle duality | particles and waves are connected on a fundamental level, called this (this is the fundamental level waves and particles are connected on) |
| speed of light | product of frequency and wavelength c=2.998 x 10^8 m/s |
| wavelength and frequency relationship | theyre inversely proportional we cannot have a very long wave that moves very fast (related back to wave spectrum and energy trends) theyre related by the wave speed, which for light is c; c=(lambda)(freq) |
| when do standing waves form | when a wave is oscillating at a certain frequency where it looks like its standing still |
| a system with fixed endpoints | restricts the number and type of possible waveforms so we refer to the number of half wavelnegths before the endpoints as n |
| planck's constant | light is transmitted in packets called photons (particle like properties) h=6.626 x 10^-34 Js (doesnt need to be memorized) |
| photoelectric effect explanation | light is transmitted in packets called photons these photons interact with matter in quantifiable ways, and the energy of a photon is proportional to its frequency E=hv, v=c/lambda, so E=hc/ lambda how remotes work! |
| line spectra | contrasts continuous light can also occur as discrete line spectra having very narrow line widths interspersed throughout the spectral regions each emission line consists of a single wavelength of light |
| 2 types of light spectrums (spectra) | continuous (the one with x ray, visible light, etc) line spectra |
| how to produce line spectra | exciting a gas at low partial pressure using an electrical current, or heating it requires matter to be spread out |
| what determines line spectra | each element has its own distinct line spectra |
| wave/particle duality | both light and very small bits of matter (like electrons) can act like BOTH particles AND waves this "wave/particle duality" results in phenomena that we dont observe for large objects that we see in our everyday life ONLY FOR SUPER SMALL PARTICLES |
| tying quant mechanics to why we care | quantum mechanics uses the math of waves to describe matter in terms of probability distributions to figure out where the wave like matter will be most rich you can also think of this as where you are most likely to find a particle |
| more on tying quant mechan to why we care | these wave equations result in only certain energy levels being "allowed" for any system with a small amount of mass resting in a potential well (i.e. attracted or repelled from something else) the |
| why do we care about quantum mechanics and this light wave stuff | those idea combine to explain electronic structure in atoms which then explains why atoms only emit certain wavelengths of light |
| question | if electromagnetic radiation can have particle-like chaaracter, can electrons and other submicroscopic particles exhibit wavelike character? Louis de broglie explored this in 1925 |
| Louis de Broglie | extended the wave-particle duality of light that einstein used to resolve the photoelectric effect paradox to material particles extended to lambda=h/mv=h/p |
| de Broglie wavelength | a characteristic of particles, not electromagnetic radiation |
| how do you know if youre looking at a de Broglie wavelength or electromagnetic radiation wavelength | are you looking at a particle or at light? particle--de broglie light--electromagnetic |
| C.J. Davisson and L.H. Germer | demonstrated experimentally that electrons can exhibit wavelength behavior (video in class) they showed that electrons travelling through a regular atomic pattern in a crystal produced a interference pattern, which is strikingly similar to that of light |
| more on davisson and germers findings | the interference pattern for electrons passing through very closely spaced slits demonstrates that quantum particles such as electrons can exhibit wavelike behavior the experimental results illustrated here demonstrate the wave-particle duality in e- |
| einstein | proposed that light consists of quanta or particles of electromagnetic energy called photons the energy of each photon is proportional to its frequancy (e=h(lambda)) |
| photoelectric effect | the ejection of an electron from the surface of a metal or other material when light shines on it |
| particle wave duality of light | light has properties of both waves and matter neither understanding is sufficient alone |
| continuous spectrum | contains all wavelengths of light |
| line spectrum | shows only certain colors or specific wavelengths of light when atoms are heated they emit light this process produced a line spectrum that is specific to that atom "emission spectra" |
| energy level postulate | an electron can have only certain energy values, called energy levels energy levels are quantized (aka not existing in a spectrum; specific values) |
| transitions between energy levels | an electron can change energy levels by absorbing energy to move to a higher energy level or by emitting energy to move to a lower energy level (when they talk about how much energy it takes to ____) |
| when delta E is positive | light is absorbed by an atom (when electron transition is from lower n to higher n (nf>ni) |
| delta E is negative | light is emitted from an atom (its lost; negative) when electron transition is from higher n to lower n (nf<ni) |
| when nf=infinity | electron is ejected |
| heisenburg uncertainty principle | by Werner Heisenberg he determined there is a fundamental limit to how accuratley one can measure both a particle's position and its momentum simultaneously the more accuratley we can determine its position at that time, and vise versa equation for it |
| erwin schrodinger | extended de broglies work by incorperating the de broglie relation into a wave equation this is today known as schrodinger equation thought of the electron in terms of a 3d stationary wave, or wavefunction represented by psi |
| max born | electrons are still particles and so the waves represented by psi are not physicial waves but instead are complex probability amplitudes |
| waveforms | can be used to determine the distribution of the electrons density with respect to the nucleus in an atom |
| as n (principle quantum number) increases | energy of atomic orbitals increases and size of orbital increases and the electrons spend more time farther from the nucleus thus, attraction to nucleus is weaker and energy associated with orbital is higher (less stable) |
| when the orbital is higher | attraction to nucleus is weaker when orbital is higher it is also less stabilized |
| in any atom with two or more electrons | the repulsion between the electrons makes energies of subshels with different values of l differ |
| which orbitals do electrons fill first | low energy |
| electron configuration of an atom | the arrangement of electrons in the orbitals of an atom is a particular distribution of electrons among subshells |
| in 2p6, for example, what does each number/letter represent | 2; the principle quantum shell, n p; the letter that designates the orbital TYPE 6; a superscript number that designates the number of electrons in that particular subshell |
| orbital diagram of an atom | shows how the ORBITALS of a subshell are occupied by electrons orbitals represented by a circle electrons represented by arrows; up for +1/2 spin and down for -1/2 spin |
| aufbau principle | "building up principle" scheme to reproduce the ground-state electron configurations by successively filling subshells with electrons in a specific order (building up order) this order is filling orbitals LOWEST TO HIGHEST ENERGY ex: 1s2 first--lowest |
| valence electrons | the electrons occupying the orbitals in the outermost shell (highest value of n) the highest in energy; what usually reacts in experiments |
| core electrons | the electrons occupying the inner shell orbitals |
| the core electrons represent | noble gas configuration |
| complete electron configuration | shows every subshell explicitly all of the numbers and letters across the period; the whole thing |
| noble gas configuration | substitutes the core configuration for the preceding noble gas and explicitly shows subshells beyond that |
| electron configurations for transition metals | additional electrons added to 3d subshell d subshells can fill to capacity with 10 electrons 5 d orbitals have a combined capacity of 10 electrons after 3d subshell is filled, electrons are next added to the 4p subshell |
| electron configuration for inner transition metals (La through Lu and Ac through Lr) | electrons are added to an f subshell for l=3 (f orbital) there are 2l+1=7 values of ml seven f orbitals have a combined capacity of 14 electrons |
| exceptions to building up order (aufbau) | Chromium (24) and copper (29) have been found to have different ground electron configurations Cu: expect [Ar] 4s2, 3d9///actual: [Ar] 4s1, 3d10 Cr; expect [Ar] 4s,3d4 actual; [Ar] 4s1, 3d5 in each, difference is in 3d and 4s subshells |
| electron configuration for main group elements | in the form ns^Anp^B sum of A and B is equal to group number so for an element in group VA of 3rd period, config would be [Ne]3s2,3p3 |
| how was hunds rule formed | in 1927 Friedrich Hund discovered, by experiment, a rule for determining the lowest-energy configuration of electrons in orbitals of a subshell |
| Hunds rule | the lowest energy arrangement of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin before pairing electrons (fill with single electrons before pairing up) |
| ground state | lowest energy configuration of an atom any other configuration is an excited state |
| excited state | any configuration other than ground state electrons have jumped to higher subshells (keep in mind subshells cannot hold more electrons than its allowed (ex: no p^7)) |
| can subshells hold more electrons than assumed | no no p&7 or d^13 |
| Pauli exclusion principle | no two ELECTRONS can have exactly the same set of 4 numbers (cant be in same orbital and spinning same way) |
| group 3A and 5A metals often exhibit | 2 different ionic charges; one that is equal to the group number and one that is 2 less than the group number the higher charge is due to loss of both s subshell electrons and the p subshell electrons lower charge due to only loss of p subshell electr |
| transition metals | form several ions the atoms generally lose the ns electrons before losing the (n-1) d electrons as a result, one of the ions transition metals generally form is the +2 ion |
| cation | forms when one or more electrons are removed from an atom |
| cation formation for main group elements | electrons that were added last are the first removed ex: Be: 1s2, 2s2 //Be2+; 1s^2 |
| cation formation for transition metals and inner transition metals | the highest ns electrons are lost first, and then the (n-1)d or (n-2)f electrons are removed |
| anion formation | electrons are added in the order predicted by the aufbau principle |
| anion | negatively charged ion formed when one or more electrons are added to a parent atom |
| inner transition metals | that block at the bottom f block 58 through 71 and 90 to 103 ** |
| why do elements take the charges they do | to get a full valence shell aka the nearest noble gas's electron config ex: P is 3 away from having Ar config, so it wants to gain 3 electrons to get all 8; this gives it its typical 3- charge |
| magnetic properties of atoms | although an electron behaves like a tiny magnet, two electrons that are opposite in spin cancel each other so only atoms with unpaired electrons exhibit magnetic susceptiblity this allows us 2 classify atoms based on their behavior in a magnetic field |
| when do we see magnetism in an atom | only when it has unpaired electrons it will behave a certain way in a magnetic field |
| paramagnetic substance | one that IS weakly attracted by a magnetic field, usually as result of unpaired electrons |
| dimagnetic susbtance | NOT attracted by a magnetic field generally because it only has paired electrons di=die magnetism |
| periodic law | states that when the elements are arranged by atomic number, their physical and chemical properties vary periodically (by period) aka organization of periodic table means something properties of elements change as their electronic structure changes |
| 3 periodic properties | atomic radius, ionization energy, electron affinity these govern the chemical behavior of elements |
| size (radius) of atoms and ions | changes for covalent and ionic radii find them in the stack |
| covalent radii | 1/2 the distance between the nuclei of 2 identical atoms when they are joined by a covalent bond |
| ionic radii | the radius of a monatomic ion in an ionic crystal structure the measure used to describe the size of an ion |
| ionization energies | the amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state how much energy do i need to put in or take out to take away an electron ex: Na to Na+ opposite of electron affinity |
| electron affinities | the energy change for the process of adding an electron to a gaseous atom to form an anion ex: Fe to Fe- opposite of ionization energy |
| atomic radius | while an atom does not have a definite size, we can define it in terms of covalent radii (the radius in covalent compounds) |
| atomic radius trend | within each group (vertical) the atomic radius increases with period number explained by the fact that each successive shell is larger than the previous shell as you go down a group it increases but across a period it decreases |
| effective nuclear charge | the positive charge that an electron experiences from the nucleus equal to the nuclear charge, but is reduced by shielding or screening from any intervening electron distribution (inner shell electrons) outside e pull exerted on an electron by nucleus |
| nuclear charge (Z) vs effective nuclear charge(Zeff) of hydrogen | theyre the same because there is only 1 electron nothing to block |
| why is effective nuclear charge different from nuclear charge | electrons are partially shielded from the pull of nucleus by other electrons present |
| trend of effective nuclear charge | increases as we move L to R across period so, atomic size decreases across a period because as that center pull get stronger, radius gets smaller because theyre physically pulled tighter to nucleus |
| a cation is always _______ than the atom from which it is derived | smaller -as electrons are removed, remaining electrons experience a larger effective charge and are drawn closer to nucleus -less road blocks Al larger than Al 3+ |
| an anion is always __________ than the atom from which it is derived | larger -as electron are added, a greater repulsion results and a decrease in the effective nuclear charge, which causes valence e to be farther from nucleus more road blocks S smaller than S2- |
| isoelectric | atoms and ions that have the same electron configuration ex: N3-, O2-, F-, Ne, Na+, Mg2+, Al3+ (1s2, 2s2, 2p6) for isoelectric atoms or ions, number of protons determines the size ex: N3- and F- have same electrons, but N3- has less protons so is larg |
| in a series of isoelectronic ions and atoms, the greater the nuclear charge_______________ | the smaller the radius |
| ionization energy | the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state |
| ionization trend | down a group, first ionization energy decreases explained by understanding that the smaller an atom, the harder it is to remove an electron, so the larger the ionization energy |
| what accounts for small deviations in ionization energy between groups IIA and IIIA and groups VA and VIA | examining the valence configurations for these groups helps us understand IIA ns2 IIIA ns2np1 it takes less energy to remove the np1 electron than ns2 electron VA ns2np3 VIA ns2np4 it takes less energy to remove the np4 |
| what are the excepptions for ionization energy trend | groups IIA and IIIA and then groups VA and VIA |
| when do we see a dramatic increase in ionization energy | when the first electron from the noble gas core in a configuration is removed |
| ionization energy trend explained | electrons can be successively removed from an atom. each successive ionization energy increases, because the electron is removed from a positive ion of increasing charge (therefore greater pull--requires more enrgy |
| electron affinity | the energy change for the process of ADDING an electron to a neutral atom in the gaseous state to form an anion |
| in electron affinity, a negative energy change indicates | exothermic indicates a stable anion is formed the larger the negative number, the more stable the anion small negative energies indivate a less stable anion |
| in electron affinity, the larger the negative number | the more stable the anion |
| in electron affinity, a positive energy change indicates | endothermic the anion is unstable (more negative more stable) |
| what participate in chemical bonding | valence electrons |
| a cation is always ________ than the atom it came from | smaller |
| an anion is always _________ than the atom it came from | larger |
| ionic radius for isoelectric series as atomic number increases | it decreases |
| covalent bond | is a chemical bond in which two or more electrons are shared by two atoms |
| why should two atoms share electrons | if its not favorable for one to take the other then one will be unhappy if they share they both win |
| ionic bond | favorable for one to give the other one an electron then theyre oppositely charged and that electromagnetic energy is keeping them together |
| double bond | two atoms share two pairs of electrons 2 lines 4 electrons |
| triple bond | two atoms share 3 pairs of electrons 3 lines 6 electrons |
| bond length trend | triple bond < double bond < single bond the shorter the length the stronger/tighter the bond |
| polar covalent bond/polar bond | a covalent bond with greater electron density around one of the two atoms 1 atom has greater electron density than the other |
| electronegativity | the ability of an atom to attract towards itself the electrons in a chemical bond hold onto electrons stronger nonequal custody |
| linus pauling | made many important contributions to the field of chemisty he was also a prominent activist, publicizing issues related to health and nuclea weapons |
| no difference in electronegativity bond type | covalent bond |
| difference in electronegativity >/= 2 (greater than 2) | ionic bond |
| difference in electronegativity 0< and <2 | polar covalent |
| polar covalent | partial transfer of electrons |
| ionic | transfer of electrons |
| covalent | share electrons |
| formal charge | of an atom the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to the atom in a Lewis structure |
| resonance structure | one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure the best one is going to have a formal charge of 0, or lowest and will satisfy octet rule |
| how do we measure the strength of covalent bonds | by the energy required to break it (energy necessary to seperate the bonded atoms) |
| bond energy | aka bond dissociation energy the energy required to break a specific covalent bond in one mole of GASEOUSmolecules |
| bond energy rankings | single bond < double bond < triple bond triple bond requires hella energy to break because theyre real tightly bound |
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