Test 3 Flashcards
how can you tell valence electrons from the periodic table?
group
how can you tell the number of shells from the periodic table?
period
Orbitals and orientations (there can be two electrons in each orientation)
Spherical - 1
Peripheral - 3
Diffuse - 5
Fundamental - 7
Orbital
- also known as an orientation
- a “pocket” within the subshell that can hold up to 2 electrons of different spin
Aufbau Principle
At ground state, electrons must have the lowest possible energies
Pauli exclusion principle
If n, l, and ml are the same for two electrons, they must have different ms (or spin)
-If a box already has an arrow, the second arrow must point in the opposite direction.
Hund’s rule
- orbitals within a subshell have equivalent energy, and since electrons repel, they spread out through out the orbitals as much as possible
- Because electrons repel each other, they spread out until they no longer can
how many orbitals and electrons can fit in the s subshell?
1-2
how many orbitals and electrons can fit in the p subshell?
3-6
how many orbitals and electrons can fit in the d subshell?
5-10
how many orbitals and electrons can fit in the f subshell?
7-14
principle quantum number
- n
- the number of shells
- n = 1,2,3 etc..
angular momentum quantum number
- l
- The subshell
- l = 0,1,2,3 for s, p, d, and f respectively
magnetic quantum number
- ml
- tells the orbital
- 1,0 +1
slit screen experiment
-showed that electrons must be spinning because sometimes they are attracted to the magnet. Whether an electron has an up or down spin is completely random. IF one electron in an orientation changes, the other must change as well.
spin quantum number
- ms
- ms = +1/2 up
- ms = -1/2down
- when something is spinning and it hits something, its spin will determine whether it deflects to the right or to the left
paramagnetism
a weak attraction to a magnetic field
-occurs when there are one or more unpaired electrons
diamagnetism
- weak repulsion from a magnetic field
- occurs when all electrons are paired
two forms of stability
- having the lowest possible energy
- having exactly filled or half-filled subshells
however, having lowest possible energy in order to be stable breaks down in the transition metals because….
sometimes, an electron will move up a level (breaking the principle) to further stabilize an atom by having either a filled or half-filled subshell for the valence electrons
levels of stability (even charge distribution leads to the most stability
- most stable: full s and p
- stable: full subshells in general
- medium stable: exactly half-filled subshells
- not particularly stable: anything else
transition metals
- have incompletely filled d subshells, and will readily lose electrons to try to stabilize
- Form cations (although they will not be quite isoelectronic with a noble has, they will have full s and p subshells like a noble gas
ionic reaction
any time something loses/gains electrons
-happens because every participating atom either loses or gains electrons ignored to become isoelectronic with a noble gas
atomic radius
- increases diagonally from the top right to the bottom left (protons decrease to the left and shells increase toward the bottom)
- Increases in the same pattern as shielding because shielding lessens the nucleus’s ability to hold on to its electrons
- distance from the nucleus to the valence shell (measured as 1/2 the distance between two nuclei that are bonded together
Electronegativity (EN)
- increases to the top right (the less shielding there is, the more electronegativity there will be)
- measures how much an element will own an electron in a bond
- An electron spends more time with the nucleus that is more electronegative
electron affinity (EA)
- Increases diagonally from the bottom left to the top right (represents how attractive a nucleus is to an electron)
- Highest in the halogens. This is because the noble gases do not want to receive electrons
- The energy that is released when an electron is gained (to produce an anion)
- If the attraction is strong, lots of energy is released
- EA and IE are basically the same thing, just with different experimental methods
Ionization energy (IE)
- increase diagonally from the bottom left to the top right
- the energy it takes to remove an electron
- we can electrocute atoms, and if there’s enough energy, an electron will fly off and form a cation
- it takes more energy to remove the second electron than the first and so on. The energy before the big jump is the final valence electron.
shielding
- electrons are attracted to the nucleus and repelled by each other
- Because of repelling core electrons, outer electrons do not feel as attracted to the nucleus
- proportional to the number of shells. Adding more electrons to a row does not really make a difference. In fact, it decreases shielding because across a row, number of protons increases as number of electrons increases.
- Inversely proportional to the number of protons (fewer protons = more shielding towards the bottom left)
effective nuclear charge
Zeff
- The force of attraction is inversely proportional to the distance o
- Fbetween charges [fish proportional symbol] 1/ distance
- Fbetween charges is proportional to the magnitude of the charge
force
the interaction between two charges
when you cross the staircase, do elements want to gain or lose electrons?
they want to gain electrons
covalent bond
- when the electrons in a bond are shared equally
- happens among nonmetals
- delta EN among non polar covalent bonds ranges from 0-0.5
- Strong bonds (not broken by melting/boiling)
- Atoms have no charge
Polar covalent bond
- electrons are shared unequally
- transfer of an electron that results in charges that attract
- occurs among nonmetals that are far enough away from each other
- 0.5-1.7
- Atoms have slight positive and slight negative charges
ionic bond
- when electrons are transferred
- always happens between a metal (cation) and a nonmetal (anion)
- A big difference in EN
- 1.7-3.3
- Atoms have full charges
- Disconnect in water
- Satisfies the octet rule by transferring electrons
metallic bonds
- valence electrons are shared collectively, delocalized across all the nuclei
- the source of a metal’s strength, conductivity, and malleability
dative bond
if something shares a pair of electrons, and both electrons originally came from
octet rule
generally, all atoms will have 8 valence electrons (full s and p subshells) in the final structure
bonding pairs
a pair of shared electrons
lone pair
a pair of electrons belonging to only one atom
formal charge
represents how many electrons are gained or lost by an atom (valence electrons - associated electrons)
- associated electrons = unshared + shared/2
- the more stable structure has more evenly distributed formal charge
ionic compound binary naming
- name of the metal + name of the nonmetal with -ide suffix
- metal (cation) always comes before the nonmetal (anion)
- If a transition metal, the charge is a roman numeral after it
- if a polyatomic atom, use the collective name
molecular compounds
- name first nonmetal and then second nonmetal with -ide suffix (in the order of the formula)
- If more than one of each type of atom, use a prefix to denote the number
1
mono
2
di
3
tri
4
tetra
5
penta
6
hexa
7
hepta
8
octa
9
nona
10
deca
