AP Chem Unit 1

Binding Energy/Ionization Energy

• the energy required to remove an electron from an atom, molecule, or ion

IE = hv - KE

IE: binding/ionization energy

hv = photon energy

KE = kinetic energy

across a period: increases because valence electrons are bound more tightly with higher Zeff

down a group: decreases because valence electrons are farther away and more shielded

Electron Affinity

• an atom’s tendency to gain an electron

• most elements have a negative(exothermic) first electron affinity

• a positive EA means energy must be put in to force the atom to accept the electron(already stable elements like noble gases)

• opposite of ionization energy

across a period: becomes more negative(greater tendency to gain an electron) because effective nuclear charge increases and atomic size decreases, leading to stronger pull on incoming electron

down a group: becomes less negative as less energy is released(due to increased atomic radius and the shielding effect) EXCEPT second period often have a less negative EA than the element directly below them

electronegativity

an atom’s tendency to attract shared electrons towards itself in a chemical bond

across a period: increases because nucleus’ stronger pull on shared bonding electrons

down a group: decreases because increased distance weakens pull

s subshell(type of apartment)

spherical, one orbital, max 2e-

p subshell(type of apartment)

peanut, three orbitals, max 6e-

d subshell(type of apartment)

cloverleaf, five orbitals, max 10e-

f subshell(type of apartment)

flower, seven orbitals, max 14e-

organization of electron “levels”

principal energy levels—>subshell—>orbital(max 2e-)

quantum “number” n

principal - distance from nucleus, any positive integer

quantum “number” l

angular - shape of orbital, any value from 0 to (n-1)

quantum “number” m

magnetic - orientation in space, any integer from -1 to 1

quantum “number” s

spin - electron spin, -.5 or +.5

Aufbau’s Principle

electrons fill atomic orbitals in order of increasing energy, so electrons occupy lowest available energy levels first

• seek most stable(lowest energy configuration)

Pauli Exclusion Principle

no two electrons in an atom can have the exact same set of four quantum numbers

• an individual orbital holds a maximum of two electrons(which must have opposite spins)

Hund’s Rule

when filling a set of degenerate orbitals(with the same energy, like three p orbitals or five d orbitals), electrons will occupy the orbitals singly with parallel spins before pairing up

• electrons do this to maximize spin multiplicity to achieve a lower energy, more stable configuration by minimizing electron-electron repulsion

electron-electron repulsion

when two negatively charged electrons meet, they repel

• occurs in shielding, where core electrons repel valence electrons

Coulomb’s Law Formula

F = k((q1)(q2)/r²)

Coulomb’s Law

• the more protons in an atom’s nucleus, the higher the charge of the atom is, therefore it will have a higher effective nuclear charge and a stronger force(makes it harder to remove an electron)

• a greater distance between the nucleus and the valence electrons equals a lower effective nuclear charge and a weaker force(makes it easier to remove an electron)

Shielding Effect

core electrons shield the outer valence electrons from the pull of the nucleus, meaning the farther an electron is from the nucleus, the weaker the proton pull of the nucleus is

As effective nuclear charge increases, how is the nucleus’ attraction for an incoming electron affected?

it is stronger

Effective Nuclear Charge

net positive charge experienced by valence electrons

Across a Period: atomic radius decreases, therefore positive nuclear charge increases, pulling electrons in more tightly→Zeff increases

Down a Group: adds more principal energy levels, so shielding and distance occur because of increased atomic radius→Zeff kind of decreases

Why does atomic radius increase as you move down a group, even though the number of protons increases?

addition of new principal energy levels overrides proton pull

The Shielding Formula

Zeff = Z - S

Z = atomic number

S = shielding constant(the # of inner electrons)

Atomic Radius

• radius of an atom is ½ of the internuclear distance

• when an atom forms a positive ion by losing one or more electrons, the ionic radius is smaller than the atomic radius of the atom

• when an atom forms a negative ion by gaining one or more electrons, the ionic radius is larger than the atomic radius of the atom

across a period: decreases, Zeff increases as protons are added, pulling electrons towards the nucleus

down a group: increases, add new principal energy levels, causing more shielding, but also Zeff decreases

When an atom loses an electron it becomes

postively charged(cation), often metals

When an atom gains an electron it becomes

negatively charged(anion), often nonmetals

Photoelectron Spectroscopy

• takes stream of a certain atom and sends high energy photons to it

  • photons knock out electrons into the magentic field, that deflects them

• closer electrons don’t get as far

• valence electrons are the spike with the lowest binding energy(furthest from nucleus)

• if an atom has less protons than the original atom, it will have its “spikes” to the right of the original atom

isoelectronic

• two or more atoms that have the same number of electrons and the same electron configuration

• for an atom to be stable, it must be isoelectronic with a noble gas

Electron Affinity is affected by

effective nuclear charge - increased Zeff means EA is more negative because more energy is released

atomic radius - as atomic size increases, the electron is added farther away from the nucleus and the attractive force is weaker, meaning a less negative EA

electron configuration(stability) - stable electrons configurations will have low or even positive EA values because adding an electron disrupts that stabolity

To write the electron configuration for a cation/anion:

1. Add/subtract however many electrons needed

2. Find the element that matches

3. Use that element’s electron configuration