Chemistry 101 Final Unit 1 Flashcards
What are elements, compounds and mixtures?
- element: only one kind of atom, cannot be broken down physically or chemically
- compound: 2 or more elements chemically bonded into molecules or extended structures + only broken down into component elements through chemical changes
- mixture: heterogeneous or homogeneous, 2 or more elements physically bonded, can vary in their parts by pass + will retain many of the properties of its components + only broken down through physical changes
What is Dalton’s Atomic Theory + the mass laws it accounted for?
- Dalton’s Atomic Theory: matter consists of indivisible atoms,
atoms cannot be converted to atoms of another element but
can combine with other types of atoms to form compounds- incomplete, doesn’t explain how bonds form + new
experiments suggest an atom could be divisible into
subatomic particles (electrons + nucleus)
- incomplete, doesn’t explain how bonds form + new
- 3 mass laws
- conservation of mass: total mass is same in chemical in a
chemical reaction - constant (definite) composition: all samples of a given
compound consist of same fraction (ex. calcium carbonate) - multiple proportions: in different compounds of same
elements, masses of combining elements can be
expressed in simple ratios (ex. CO + CO2)
- conservation of mass: total mass is same in chemical in a
What did Thomson discover?
- existence of negatively charged particles + that they’re smaller and lighter than atoms
- proposed plum pudding model of atom: maintains neutral charge of atom (area of positive charge + negatively charged particles)
Atoms can lose or gain electrons to form charged species called?
Ions (atomic number: # of protons = # of electrons)
What are ionic compounds?
- name cation then anion
- add -ide to name of anion
- for metals that can form more than one ion, specify the charge by Roman numerals
What are polyatomic ions (ionic compound)?
- consists of 2 or >2 atoms joined by covalent bonds + exists as charged unit (ex. NaCN is not Na + C + N, but Na + CN ions)
- # of O increases: hypo-ite, ite, ate, per-ate
What are binary compounds (covalent compound)?
- written in order of increasing EN
- add -ide to name of 2nd element
- add numerical prefixes where necessary (mono prefix is not used for 1st element)
- ex. N2O: dinitrogen monoxide
What are acids (covalent compound)?
- produce H+(aq) when dissolved in water (formula has at least one H atom usually written first)
- binary acids have no O: hydro-ic acid (ex. HF(aq): hydrofluoric acid)
- oxyacids contain H+O+non-metal, name depends on oxoanion
- -ate: -ic acid, -ite: -ous acid
What is spectroscopy?
Study of interaction of matter with radiated energy.
What is wavelength (λ), frequency (v), amplitude + speed?
- wavelength: distance between 2 consecutive crests (ex. distance the wave travels during one cycle)
- units: m, cm, nm
- frequency: # of cycles the wave undergoes per second
- units: s^-1: H2
- amplitude: height of cresti related to intensity of radiation
- speed of a wave: distance it moves per unit time
- formula: C: λv
Each atom changes it’s energy by an amount, △Eatom…
When it absorbs or emits one proton (one particle of light) whose energy is related to its frequency (not its amplitude).
What is threshold frequency?
Electrons break free when it absorbs a photon of enough energy. Electrons cannot be ejected by saving up energy, whose energy isn’t sufficient.
What is absence of a time lag?
Some current flows as soon as light of sufficient energy (frequency) strikes the metal surface.
What is atomic spectra?
- an element that’s vaporized + excited will emit a series of fire lines (not a continuous spectrum)
- spectral lines are produced when electrons move from one energy to another energy level
- absorption: excitation from lower allowed energy level to higher energy level (absorbs a proton)
- emission: relaxation from higher allowed energy level to lower energy level, often ground level (emits a proton)
What are the 3 parts of Bohr’s Model + its limitation?
- electron occupies stationary states
- ground state: electrons in first orbit (n:1), closest to nucleus + H atom is in its first energy level
- excited state: electrons is in any orbit farther from nucleus (n>1)
- ex. if n:2, the atom is in the first excited state
- n:∞: indicates an electron has been completely removed from nucleus
- electron doesn’t lose energy while in a stationary state
- electron undergoes a transition from one stationary state to another by absorbing or emitting photons
- absorption process: nf > ni, △E always +
- emission process: nf < ni, △E always -
- limitation: fails to predict line spectrum of any other atom + completely for atoms with more than one electron since e-e repulsions, nucleus-e are unaccounted for.
- H-like species: one electron species
- He+: 1 e, 2 p, z:2
- Li2+: 1 e, 3 p, z:3
- H-like species: one electron species
What are infrared, visible + ultraviolet series?
- infrared: nf:3, when electron drops from outer orbits (ni:4,5,6…) to n:3 orbit
- visible (Balmer) series: nf:2, when electron drops from outer orbits (ni:3,4,5…) to n:2 orbit
- ultraviolet (Lyman) series: nf:1, when electron drops from outer orbits (ni:2,3,4…) to n:1 orbit
What does the smallest energy transition correspond to?
Shortest distance on energy plot since smallest △E corresponds to longest λ.
What is electron probability density?
A measure of the probability of finding the electron in some tiny volume of the atom. It decreases with distance from nucleus along line.
What are the 4 quantum numbers for the state of an electron?
- principle quantum number (n)
- allowed values: any + integer
- size of orbital
- angular momentum quantum number (l)
- allowed values: 0 to n-1
- shape of orbital
- 0:s, 1:p, 2:d, 3:f
- magnetic quantum number (ml)
- allowed values: -l-+l
- spatial orientation of orbital (e cloud)
- 1 s-orbital, 3 p-orbitals, 5 d-orbitals, 7 f-orbitals
- spin quantum number (ms)
- describes spin of the electron (not orbital)
- half of electrons are attracted by large external magnetic
field while other half are repelled by it - electrons with same spin value: parallel spin
- electrons with different spin value: opposing spin
Orbital shapes
- s-orbitals are spherical, p-orbitals nucleus lies at nodal plane, d-orbitals orientation of nodal planes always lie between orbital lobes
What is Pauli’s Exclusion Principle?
No two electrons can have all four quantum numbers alive, each orbital can be occupied by maximum of 2 electrons of opposing spin.
- ex. specify all quantum #s of the 2 electrons in a He atom
He atom: 2 p, 2 e
1. n=1, l=0, ml=0, ms=+1/2
2. n=1, l=0, ml=0, ms=-1/2
How do orbital energies differ in one vs many electron atoms?
The outermost electron feels…
- attraction from positive nucleus
- repulsion by other e-e is somewhat screened/shielded from
the nucleus by other electrons
What is the effect of nuclear charge (Zeff)?
The higher the nuclear charge, the greater the pull on the electron. Energy is lowered and that the system is more stable.
What happens during shielding?
Inner electrons prevent the outer electron from experiencing the full attraction of the nucleus, making its removal somewhat easier than expected.
What does orbital energy depend which factors?
- greater nuclear charge lowers orbital energy
- e-e repulsions raises orbital energy
- electrons in outer orbitals (higher n) are shielded from the full nuclear charge, so they have higher energy
- orbitals with good penetration (having e-density close to nucleus) have lower energy
- penetration: how well the outer electrons are shielded by
the inner electrons
- penetration: how well the outer electrons are shielded by
What is Hund’s rule?
When orbitals of equal energy are available, maximize unpaired spins.
What’s special about Cr + Cu during electron configurations?
They have “anomalous” configurations attributed to “special stability” of half filled or completely filled 3d subshells.
- ex. Cr: [Ar]4s^2-3d^4 to [Ar]4s^1-3d^5
For electron configuration of ions…
- metals tend to lose valence electrons to form cations
- transition-metal atoms can form several cations, with the ns electrons removed before the (n-1)d electrons
For electron configuration of ions…
- metals tend to lose valence electrons to form cations
- transition-metal atoms can form several cations, with the ns electrons removed before the (n-1)d electrons
- nonmetals tend to gain valence electrons to form anions
Difference between paramagnetic and diamagnetic?
- paramagnetic: unpaired electrons
- diamagnetic: all electrons are paired
For Zeff…
- increase Zeff as there’s a net + charge: decrease radius (half the size of neutral atom)
- decrease Zeff as there’s a net - charge: increase radius (2x the size of neutral atom)
- decreases down the group, increases across period
- ex. rank in order of increasing size: K+, S2-, Cl-
K < Cl < S (more negative ion: more e: bigger size)
For atomic radius…
- increases down the group (less Zeff cause electrons aren’t held as tight, so the probability that outer electrons spend most of their time from nucleus increases), decreases across period (electrons are held tighter)
What is ionization energy and its trend?
- energy that must be added to remove an electron from isolated gaseous atom or ion
- 1st IE: outermost electron (x = x+ + e)
- 2nd IE: next electron (x = x2+ + e)
- each successive ionization needs more energy (IE4 > IE3
>IE2>IE1)
- decreases down the group, increases across period
- anomaly: Be VS B: B has 1 e in 2p orbital (more energy than
2s) hence removing this e is easier than removing a 2s e
from Be which has a full subshell)
- anomaly: Be VS B: B has 1 e in 2p orbital (more energy than
What is electron affinity and its trend?
- the energy change that occurs when a gaseous atom or ion gains an electron
- EA1 is mostly negative (when an electron is added to a neutral atom, energy is released in an exothermic process) where EA2 is always positive (the energy required to add a second electron to an anion)
- decreases down the group, increases across period
What is the difference between reducing agent and oxidizing agent?
- reducing: X loses e, X is oxidized by Y, oxidation # of X increases
- oxidizing: X gains e, Y is reduced by X, oxidation # of Y decreases)
Are alkali metals good oxidizing or reducing agents?
Reducing agents, since they want to lose an electron so they want to be oxidized.
Are halogens good oxidizing or reducing agents?
Oxidizing agents, since they want to gain an electron so they want to be reduced.
How can we tell if a substance is likely a good oxidizing or reducing agent?
- good reducing agents: low IE + EN, small e-affinity, low oxidation state
- good oxidizing agents: high IE + EN, large e-affinity, high oxidationg state
