Chemistry Fundamentals

Emission

(-) energy change - Electrons in an excited state can return to a lower energy orbit, emitting a photon equal in energy to the energy difference between the energy levels - electrons can return to the ground state in a single transition, or in multiple transitions E3 (excited state) --> (E2 excited state) = when electron drops to lower energy level the electron emits photon

electrons

(e-) charge = -1 mass = 0 amu The charge (C) gives: - the number of (p+) - e- ex: 12p+ and 10 e- = 2+ (2 positive charges) ex: C = p+ - e- e- = -C + p+ e- = -(2) + 12 = 10e-

neutrons

(n^0) charge = 0 mass = 1 amu = 1 g/mol The mass number (A) = (p+) + (n^0) : - determines the isotope

protons

(p+) charge = +1 mass = 1 amu = 1 g/mol atomic number (Z) gives: - number of p+ - nuclear charge - determines the element * not possible for the same element to have the same # of protons

Absorption

+ energy change - electrons absorb only specific, allowed quantities of energy - allowed energies match the energy difference between an electron's - n=4 is the highest energy level, furthest distance - n=1 is the lowest energy level - n=2 and n=3 are the excited states = when an electron jumps to a higher energy level: E1 (ground state) --> E2 (excited state) = incoming photon absorbed by electron Ephoton = Ef - Ei

Paramagnetic

- Any atom or ion that aligns itself parallel to a magnetic field is defined as paramagnetic, and all such species have unpaired electrons in their orbitals. - paramagnetic atoms (which have electrons that are not spin-paired) will be ATTRACTED to an externally produced magnetic field. Q: Which of the following ions will align itself parallel to the poles of a magnetic field? A. Ca2+ = diamagnetic B. V3+ Correct Answer; paramagnetic C. Cr6+ = diamagnetic D. Se2- =diamagnetic

Rank the C-O, C-F, Si-F, and Si-O bonds in order of decreasing polarization. A. Si-O, Si-F, C-O, C-F B. Si-F, Si-O, C-F, C-O Correct Answer C. C-O, C-F, Si-F, Si-O D. C-F, C-O, Si-O, Si-F

- Based upon the trend in electronegativity (it increases as one moves to the right and up within the periodic table), the order of electronegativity is Si < C < O < F. - Therefore the most polar bond exists between atoms with the greatest difference in electronegativity. So, the bond polarizations listed in decreasing order are Si-F, Si-O, C-F, C-O (choice B). Note that based upon the trend, another possible answer might be Si-F, C-F, Si-O, C-O, but this is not given as a choice. strongest polarization will have the greatest difference in EN

Diamagnetic

- Diamagnetic atoms will be repelled by an externally produced magnetic field because all the electrons in a diamagnetic atom are spin-paired, so the individual magnetic fields that they create cancel - leaving no net magnetic field - Atoms with all electrons spin-paired are diamagnetic - must contain an even number of electrons and have all of its subshells filled

Hydrogen Absorption Spectrum

- H takes color & detector cannot catch them (why you see dark bands, this represents missing wavelength) - Hydrogen sample is absorbing the energy levels of the red, blue, and purple colors; electrons are absorbing a photon - each band represents ONE absorbed energy

Compared to carbon, nitrogen has a greater ionization energy and more negative electron affinity

- Ionization energy increases going left to right in a period due to greater effective nuclear charge (gaining protons because atomic number is increasing), resulting in valence electrons being held more tightly. Thus, nitrogen will have a greater ionization energy - electron affinity becomes more negative (i.e., more favorable) going left to right in a period (once again due to the greater effective nuclear charge), making nitrogen more negative compared to carbon (more negative because accepting electrons gives you a negative charge)

Hydrogen Emission Spectrum

- OPP of absorption spectrum - dark background with bright bands = energies released by Hydrogen sample; electrons are emitting a photon - detector was able to catch the energies released by H

The principal quantum number is a measure of which of the following? A. Approximate radial size of an electron cloud B. Approximate shape of an electron cloud C. Number of valence electrons that orbit a nucleus D. Number of protons and neutrons found in the nucleus of an atom

- The answer to this question is A because the principal quantum number n is most closely associated with the potential energy of the electron. Since potential energy is proportional to the square of the distance of two oppositely charged particles by Coulomb's Law, it is also true that n is associated with the radial "size" of the electron cloud. - n ranges from 1 -7=horizontal rows = energy shells

Oxidation Potential (OIL)

- The greater an element's oxidation potential, the more readily the atom will lose an electron and the more reactive the element is. This can be approximated by the relative ionization energies of the elements. - ionization energies increase across a row and decrease down a column of the periodic table - Elements farthest to the left are most reactive, and reactivity often increases down a metal family, especially for groups I and II. high oxidation potential = strong reducing agent - will reduce something else, and become oxidized itself highest oxidation potential: Cs Sr Ca Fe lowest oxidation potential

how to find electrons

- The number of electrons in a neutral atom is equal to the number of protons - 2 e- in innermost shell, next 8e-, next 18e-; all group 1 elements have one valence e- because they have one e- in their outer shell) ex: Oxygen has 8 electron but O2- has 10 e- GROUP = VERTICAL

Reactivity and Shielding

- The reactivity of sodium compared to potassium is that potassium has a higher reactivity because the valence electron on potassium is farther from the nucleus (has a higher principle quantum number, n). (K and Na are in same vertical group so have same # of valence e-) - Reactivity of metallic elements increases down a column of the periodic table. This increase in reactivity corresponds with increased shielding from inner electron shells, increased atomic radius, and decreased ionization energy. (= doesn't require very much energy to lose outer electron) - Thus, electrons are easier to remove in potassium versus sodium. increase shielding, increase reactivity

Does a small atomic radius correspond to a high ionization energy?

- Yes, because the shorter distance between the positive nucleus and the negative electron enhances electrostatic attraction, and thus makes it difficult for these electrons to be removed (increased effective nuclear charge as more protons are added) - Ionization energy decreases with period (n) and the second ionization energy is greater than the first. - As period increases, ionization energy decreases due in part to the greater degree of shielding present. (=less effective nuclear charge, larger radius)

Limiting Reagent

- always a reactant - the reactant you run out of first, so it limits ho much product you can produce - take what question stem gives you / balance reaction - see notes

Atomic radius

- an atom's electron cloud represents almost all of its volume - as Fe increases valance electrons are pulled more strongly toward the nucleus - as electrons are pulled toward the nucleus because they feel a greater effective nuclear charge, atomic radius decreases L to R = decrease Top to Bottom = increase, due to increased shielding experienced by the valence e- as you travel down more shells are added with each period. The valence e- are less tightly bound since they feel a smaller effective nuclear charge - as we go down a group, atomic radius increases due to the increased shielding and decreased effective nuclear charge

Moving top to bottom down a group...

- core electrons are added at the same rate as protons (Zeff remains constant) - the number of valence electrons remains the same (C remains zero) - the size of the valence shell increases (r increases; aka n increases) Fe (attraction) DECREASES going down a group (n increases) in the periodic table

Aufbau Principle

- describes how electrons are added to or removed from orbitals of DIFFERENT ENERGY - electrons are added to orbitals from lowest to highest energy - valence electrons are in the highest energy shell (highest 'n' value) - electrons are first removed from valence orbitals from highest to lowest energy - electrons are then removed from the remaining orbitals from highest to lowest energy LOWEST: 1s < 2s < 2p <3s < 3p < 4s < 3d <4p< 5s < 4d< 5p< 6s <7s< 5f HIGHEST - even though 3d is higher energy than 4s, the valence 4s electrons aree removed BEFORE the non-valence 3d electrons

Hund's Rule

- describes how electrons are added to or removed from orbitals of the SAME ENERGY (degenerate orbitals) - electrons fill degenerate orbitals (same energy) one per orbital before pairing Paramagnetic = at least one unpaired electron (ex: d5, d6) Diamagnetic = all electrons are paired (s2, p6, d10, f14; d6 can go on any of the orbitals, doesn't have to be first)

Pauli Principle

- describes the carrying capacity of an orbital - there can be no more than 2 electrons in any given orbital - no 2 electrons may be identical (i.e if they have same n-value then they must have opposite spin multiplicity) - this limits he occupancy of an orbital to a maximum of 2 electrons - first orbital can spin up or down

mole fraction

- expresses the fraction of moles in a given substance

Acidity trend

- how well a compound donates protons (Bronsted) or accepts electrons (Lewis), or lowers the pH - horizontal periodic trend = the more EN the element bearing the negative charge is, the more stable the anion will be (acidity increases L to R) - vertical trend = depends on the size of the anion. The larger the anion, the more the negative charge can be delocalized and stabilized (acidity increase going down a group/family)

Ionic Radius

- if an ion forms, the radius will decrease as e- are removed (bc the ones that are left are drawn in more closely to the nucleus), and the radius will increase as e- are added X+ < X < X- (largest because gaining electrons) cation radius < neutral-atom radius < anion radius ***ionic radius increases with increasing negative charge*** **for isoelectronic ions the species with the more protons will have the smaller radius** - an ion's electron cloud represents almost all of its volume - as Fe increases valence electrons are pulled more strongly toward the nucleus - valence electron repulsion is slightly increased in anions versus a neutral atom - valence electron repulsion is slightly decreased in cations versus a neutral atom Which is the largest? C- (6p, 7e-) = LARGEST size, less attraction due to shielding C (6p, 6e-) C+ (6p, 5e-) = strongest attraction, but SMALLEST size

Electron Affinity

- is the energy change when adding an electron to the valence shell of an atom in its gaseous state = chemical reaction necessary to determine - most elements tend to release energy upon the addition of an electron (exothermic or negative values, -EA) (ex: halogens - adding an e- gives them a desired octet configuration) - if energy is required in order to add the electron, the electron affinity is positive (ex: noble gases and alkaline earth metals because adding an e- begins to fill a new level or sublevel and destabilizes the electron configuration) - as Fe (electrostatic interactions) increases, additional electrons releases more energy (-EA) - elements with closed-shell and closed-subshell configurations tend to require energy upon the addition of an electron (endothermic or positive values, +EA) L (m) to R (nm) = increase electron affinity = more exothermic; negative electron affinities (except noble gases) Top to Bottom = decrease electron affinity and is less exothermic

The difference between thee actual mass of a nucleus and the sum of the masses of the protons and neutrons that form is called the:

- mass defect = the difference between the mass of an isotope and its mass number. (the mass number is the total number of protons and neutrons in a nucleus, and atomic weight is the relative mass of an atom based on a scale where 12C is assigned a mass of 12. The critical mass is the minimum mass of a fissionable material necessary to sustain a clear nuclear chain reaction)

Energy subshell (l)

- s-block (n) -d-block (n-1) p-block (n) f-block (n-2) ex: highest energy level for Vanadium? - n=4 - d block so 4-1, n=3 so 3d (belongs to d-block) ex: Uranium? - n=7 - f block so 7-2 = 5 so 5f (f block)

-ide

- stands for something with a negative charge ex: Fluoride (F-) has how many e-? C = p+ - e- -1 = 9p+ - (e-) e- = 10e-

standard state conditions

- temperature is 298K (25C) and pressure is 1 atm - all solids and liquids are said to be pure - solutions are at a concentration of 1M - will have circle subscript

Moving left to right across a period...

- the number of core electrons remain constant while protons are added (Zeff increases) (CORE ELECTRONS = HORIZONTAL ROW = N=1-7) - valence electrons are added at the same rate as protons (C remains zero) (VALENCE E- = VERTICAL GROUP) - the size of the valence shell remains constant (r remains constant) Fe INCREASES going across a period in the periodic table

angstrom

10^-10 m - atomic radii bond lengths are usually 1-3 A

standard temperature and pressure

= STP - OC = 273K

nuclear charge

= atomic number, number of protons an atom has - increases from left to right as the number of protons increases (atomic number) - As the positive nuclear charge increases, effective nuclear charge increases and electrons are pulled closer to the nucleus, resulting in a decreased atomic radius. - decreased atomic radius = increased nuclear charge causes electrons to be held more tightly - ex: Compared to the atomic radius of S, the atomic radius of Al is larger, due to decreased nuclear charge.

ampere

A electric current

In organic solution, F- deprotonates dissolved HCl. Which of the following explains this observation? A. F- has a smaller radius than Cl-. B. F- is a stronger acid than Cl-. C. F- has a greater electronegativity than Cl-. D. F- has a lower ionization energy than Cl-.

A and C are true but since F- has a smaller radius than does Cl-, it is less stable with excess charge, and therefore a better Bronsted base. It out competes Cl- for the proton.

Which of the following electron transitions (between energy levels, labeled by n) could account for the emission of photons of red, yellow, and blue light from a pure noble gas? A. n = 4 to n = 3, n = 3 to n = 2, and n = 2 to n = 1, respectively or B. n = 1 to n = 0, n = 2 to n = 1, and n = 3 to n = 2, respectively

A. The energy of the photons satisfy red < yellow < blue. Therefore, the energy difference between the energy levels that produce these photons must follow the same order. - Since the spacing between energy levels decreases with increasing n, then n = 4 to n = 3, n = 3 to n = 2, and n = 2 to n = 1 is best. (n=2 to n=1 is the greatest energy difference because the spacing between these orbitals is the largest, which should be paired up with blue light because it has the most energy)

Excited State Electron Configurations vs (ground state configurations = lowest energy configurations)

An excited state is any of an infinite number of configurations that have higher energy than the lowest energy electron configuration - make sure the configuration has the correct total number of electrons - the electrons can be in ANY orbital as long as that orbital exists Is 1s2,2s2,2p6,3s2 an excited state configuration for Na (11e-)? = NO, too many e-, this is Na-(12e-) Is 1s2,2s2,2d6,3s1? = NO, because 2d doesn't exist. This is the wrong excitation Is 1s2,2s2,2p6,3p1? = YES has 11e-, excitation send e- to 3p1

Deuterium

An isotope of hydrogen with one proton, one neutron, and one electron in the nucleus having an atomic weight of 2.014 2^H (1p+ + 1n^0)

Which of the following is true about the amount of shielding the highest energy electrons of calcium and arsenic experience? A. The electrons of calcium have a greater amount of shielding than the electrons of arsenic. B. The electrons of calcium have a lesser amount of shielding than the electrons of arsenic. Correct Answer C. The electrons of calcium have the same amount of shielding as the electrons of arsenic. D. The amount of shielding between the electrons of calcium and of arsenic can only be determined experimentally.

B. The amount of shielding that the highest energy electrons of an atom feels is determined by the number of filled shells in that atom (inner core electrons), as well as subshells within an energy level (eliminate choice D). Because calcium and arsenic are in the same period (row), they have the same number of inner core electrons. However, the highest energy electrons in arsenic are in the p subshell and are shielded by the electrons in the s subshell. Since the highest energy electrons in calcium are in the s subshell, there is more shielding for arsenic, eliminating choices A and C.

Charge (C)

C>0 = cation (+) C = 0 = atom (0) C<0 = anion (-)

Electromagnetic Spectrum (what does h-constant =?)

E = hf = (hc) / wavelength h = 6.6 x 10^-34 Js c = 3 x 10^8 m/s left side (gamma first) = high E high f, hi refractive index (n), low speed (v) low wavelength right side (radio first) = low E low f high wavelength, low refractive index (n), high speed (v) Roman Men Invented Very Unusual X-ray Guns Visible light: red = 700 nm (lower energy, lower frequency), violet = 300 nm (higher energy, higher frequency)

Energy shell (n) and subshell

Each period corresponds to a different energy shell - each shell is higher energy and larger than the last - n=1 to n=7 (highest), n=1 behaves the most like an ideal gas Each block corresponds to a different subshell (l) - each subshell is more complex and higher energy than the last - s-block -d-block -p-block -f-block

iron (III) ion (Fe3+)

Fe = [Ar] 4s23d6 Fe3+= [Ar] 3d5 (paramagnetic)

Which is more stable?

Fe2+ or Fe 3+ (both are paramagnetic) Fe: [Ar]4s2,3d6 Fe2+: [Ar]3d6 Fe3+: [Ar]3d5 = half filled subshells tend to be more stable half filled subshells: p3 d5 f7

Diatomic elements

Have - H2(g) No - N2(g) Fear - F2(g) Of - O2 (g) Ice - I2 (s) Cold - Cl2 (g) Beer - B2 (l) = molecules composed of two atoms. ... Covalent bonds are used to link two atoms together in a diatomic element through the action of sharing electrons. This type of bonding can be observed in diatomic elements by viewing the electron configuration of the molecule 1 mole of O2 = 32 g 1 mole oh H2 = 2 g

period on periodic table

Horizontal rows on the periodic table n=1 (H, He) n=2 (Li, Be, B, C...etc) aka principle quantum numbers n = energy shells

kelvin

K temperature 298K = 25C 273K = 1C

liter

L = unit of volume equal to 1/1000 of a cubic meter 1000 L = 1 m^3 1L = 1000 cm^3

3 Rules for Electron Filling in an Atom (PAH)

Pauli Principle Aufbau Principle Hund's Rule

Periodic Trends

TOP RIGHT: small size high IE, EN, EA BOTTOM LEFT: large size low IE, EN, EA

Ground State Electron Configurations

The ground state is the lowest energy electron configuration - make sure the configuration has the correct total number of electrons - follow the structure of the periodic table to determine which electrons are present - use the previous noble gas to the element to denote the core electrons that don't contribute to the chemistry of the atom - for CATIONS or ANIONS write the configuration of the atom and THEN add/remove electrons appropriately - closed-shell configurations (octet and duet) tend to be stable

Shielding

Valence electrons have an electrostatic attraction due to the nucleus given by Coulomb's law Fe = (Zeff + C) / r^2 (aka n^2) - Each filled shell between the nucleus and the valence e- shields/protects the valence e- from the full effect of the positively charged protons in the nucleus = nuclear shielding or the shielding effect. The electrical pull by the proton in the nucleus is reduced by the negative charges of the electrons in the filled shells in between; the result is an effective reduction in the positive elementary charge from Z (protons, atomic number) to a smaller amount = Zeff = effective nuclear charge - core electrons shield the valence electrons from the full nuclear charge - the shielding effect between valence electrons is relatively small - the nuclear charge experienced by a valence electron = the effective nuclear charge: Zeff = Z - core e-

Multiple Ionization

Which is greater, the energy required to remove an electron from an atom X (IE1), or from its cation, X+ (IE2)? - IE2 > IE1 - as the positive charge on a given ion increases, so does its ionization energy (=oxidation potential) - elements with closed-shell and closed-subshell configurations tend to require more energy to remove an additional electron ex: Ca IE1 = 600 kj/mol (valence e-) IE2 = 1200 kj/mol (another valence e- bc only double the energy) IE3 = 5000 kj/mol (this is a core electron/breaking the inner shell costs more E because it would have a full octet at this point) * do electron configurations and valence electrons to determine these questions

Copper metal (Cu) (follows same rules as Cr)

[Ar] 4s1, 3d10 - anomalous electron configurations = the exceptions - Since completely filled d subshells are more stable for transition metals, we observe that one s electron is promoted into the d subshell in the copper family, of which silver is a member. Ag ground state = [Kr] 5s1 4d10

chromium metal (Cr) (ODD one out think ODD numbers)

[Ar] 4s1,3d5 - the 1e- jumps from s to d - d5 is more stable that d4 - anomalous electron configurations = the exceptions

iron metal (Fe) (normal)

[Ar] 4s2,3d6

sulfide ion (S2-)

[Ar] = [Ne]3s2,3p6 stable, full octet

lithium ion (Li+)

[He] more stable because has full valence and acts like noble gas

lithium metal (Li)

[He]2s1

sulfur (s)

[Ne]3s2,3p4

Bohr model

a Bohr atom contains only one electron = 1 e- system so only H has one e- so must take e- away from other elements: He (0 charge) but has 2p+ so: e- = 0 +2, e- = 2 (this means one electron must be taken away) He -> He+ Li (has 3 e-) --> Li2+ (has 1 e- now)

anhydrous

describes a substance that does not contain water (could be a crystalline compound)

Coulomb's Law

electric force between charged objects depends on the distance between the objects and the magnitude of the charges.

Electronegativity

intrinsic property of an atom - EN is the ability of an atom to attract electrons to itself when it forms a covalent bond; the greater this tendency to attract electrons, the greater the atom's EN - as EN increases, shielding decreases (decreased nuclear shielding allows for a stronger pull on the valence electrons = increased effective nuclear charge) - as Fe increases, the ability to attract electrons increases MOST EN F O N Cl Br I S C = H

Ionization Energy

is the minimum amount of energy required to remove the outermost electron from an atom in its gaseous state - since an atom's positively charged nucleus is attracted to the electrons in the atom, it takes energy to remove an electron. The amount of energy necessary to remove the least tightly bound electron from an isolated atom = the atom's first ionization energy (IE or IE1) - the second ionization energy (IE2) of an atom, X, is the energy required to remove the least tightly bound electrons from the cation X+ - IE2 will always > IE1 - requires E = endothermic process - as Fe (electrostatic interactions) increases, removal of electrons requires more energy OPP to size: - moving from L to R or up a group, the IE increases since the valence e- are more tightly boundd - the IE for any atom with a noble-gas configuration will always be large L to R = more endothermic & harder to remove e- top to bottom = less endothermic and easier to remove *increase size (atomic radius) = decrease ionization energy* * can remove more than one e- but it gets harder (second ionization energy)

kilogram

kg mass

meter

m length

mole

mol amount of substance - one mole of anything contains 6x10^23 atoms (number that links atomic mass units and grams)

molarity (M)

mol of solute/L of solution - tells you concentration of a solution - [Na+] = 1.0 M = the concentration is equivalent to 1 mole of sodium ions per liter of solution

second

s time

isoelectric

same number of electrons (think isoelectric point being chargeless)

isotomes

same number of neutrons

isotopes

same number of protons