Chem Chapter 3

electron configuration diagram

use squares to represent orbitals and half arrow to represent electrons- half arrow pointing up = spin up and half arrow pointing down = spin down

pauli exclusion principle

no two electrons may have the same 4 quantum numbers

no orbital may have more than 2 electrons

electrons within the same orbital must have different spins

degenerate

orbitals with the same energy

periodic

exhibit a repeating pattern

periodic law (Mendeleev)

when the elements are arranged in order of increasing mass, certain properties recur periodically

periods

rows in the periodic table

modern periodic table organization

elements listed in order of increasing atomic number rather than mass

groups/families

columns in the periodic table

elements in the sae group have similar properties

electron configuration

description of orbitals occupied by electrons

multi-electron atoms

orbitals can be approximated (not exactly solved) and are found to be hydrogen-like

affected by energy splitting of sublevels and electron spin

electron spin

• fundamental property of all electrons (all e- have the same amount of spin)

• orientation is quantized, meaning it can only spin one way or the opposite direction

• m_s = ± 1/2

sublevel energy splitting

found in multi-electron atoms (single electron systems have all same energy in their energy levels)

• caused by charge interaction, shielding, and penetration

• lower L value = less sublevel energy

  • s<p<d<f

Aufbau principle

electrons enter atomic orbitals from lowest energy to highest

• 1s 2s 2p 3s 4s 3d 4p 5s 4d

d-block transition metals electron configuration

Cr, Mo, Cu, Ag, Pd ([Kr]5s⁰4d¹⁰) all have irregular electron configurations in which the ns only partially fill before the (n-1)d or don’t fill at all to pursue half filled stability (d⁵) or full stability (d¹⁰)

three element types

metals, nonmetals, and metalloids

main group elements

properties tend to be predictable based on position in table

labeled by # and letter A

transition elements/transition metals

properties tend to be less predictable based off their position

labeled by number and letter B

Coulomb’s law

describes electrons and repulsions between charged particles

ex. electrons are more attracted to a nucleus with a 2+ charge than a 1+ charge

shielding

repulsions cause a net reduction in attraction to the nucleus because the electron is shielded from the nucleus

• Electrons in lower orbitals are better shielders (s>p>d>f)

effective nuclear charge

the total amount of attraction that an electron feels for the nucleus’ protons

Zeff = atomic number- amount of inner electrons

penetration

• degree related to radial distribution function

• causes energies of sublevels in the same principle level to not be degenerate

penetration example

in the 4th and 5th principle levels, effects of penetration become so important that the s orbital lies lower in energy levels than d orbitals of previous principle level

• it takes less effort for the electron to move in a sphere shape than a peanut or clover shape

• the ordering can therefore vary among elements, causing variations in electron configurations of transition metals and their ions

Hund’s rule

when filling orbitals that have the same energy (degenerate), place one electron in each orbital before completing pairs

core (inner) electrons

electrons in lower energy shells that can shield

• n-1 electrons (ex. all of n=2 orbitals when n=3)

valence electrons

electrons in all the sublevels with the highest principle energy shell- do not shield (ex. all of n=3 when n=3)

transition metals valence electrons

the outermost d electrons are counted among valence electrons (ex. Rh = [Kr]5s¹4d⁸ so it has 9 valence electrons)

importance of valence electrons

• help determine chemical and physical behavior

• participate in bonding

• where electrons are added/lost to make ions

noble gases

have 8 valence electrons (except He)

especially nonreactive- particularly He and Ne

alkali metals

have one valence electron and occupy first column

will lost one electron in reactions and form cations (1+)

alkaline earth metals

have 2 valence electrons and occupy the 2nd column

will lose 2 electrons in reactions and form cations (2+)

halogens

nonmetals with 7 valence electrons

tend to gain one electron in reactions with metals to form an anion (1-)

transition metals and inner transition metals

transition metals occupy d block

inner transition metals occupy f block

S orbitals fill before d orbitals due to sublevel splitting

metalloids

located in p-block area between nonmetals and metals

• sit on metalloid “staircase” - also includes Ge and Sb

can exhibit metallic or nonmetallic properties during chemical reactions

can either lose electrons from p and then s orbitals to form cations or gain electrons in p orbitals to form anions

nonmetals

located in upper right hand side of table in p-block area

will gain p orbital electrons during chemical reactions to have noble gas configuration

form anions

metal ion formation

alkali metals form 1+ cations

alkaline earth metals form 2+ cations

transition, inner transition, and p-block metals form variety of charged cations

nonmetal ion formation

halogens (group 7A) usually form 1- anions

other nonmetals form a variety of charged anions

effective nuclear charge trend

increases across a period and decreases down a column

• as you go across a period, you are adding electrons in same energy level but different orbitals, whose shielding ability decreases

atomic radius

the average radius of an atom based on measuring large numbers of elements and compounds

atomic radius trend

decreases across a period and increases down a group

• pattern based on effective nuclear charge pattern

paramagnetism

electron configurations that result in unpaired electrons- atom or ion will have a net magnetic field and be attracted to magnetic fields

diamagnetism

electron configurations that result in all paired electrons = atom or ion will have no magnetic field and be slighly repelled by magnetic fields

Ions and Atomic Radius

• cation radius is smaller than its corresponding atom radius

  • loss of electrons results in remaining electrons experiencing larger effective nuclear charge + getting pulled closer to nucleus

• anions have larger radius because new electrons in valence shell experience smaller effective nuclear charge

ionization energy

the energy needed to remove an electron from an atom- endothermic process

• the larger an effective nuclear charge is, the more energy it takes to remove electrons

  • electrons further from nucleus are easier to remove

first ionization energy

the energy required to remove one electron from a neutral atom

ex. M(g)+IE₁ = M¹⁺(g)+1e-

second and subsequent ionization energies

energy required to remove an electron from a 1+ ion, 2+ ion, etc.

first ionization energy trend

increases across a period and decreases down a group

electron affinity

energy associated with the addition of an electron to the valence shell of an atom that is in the gas phase (ex. M(g) + 1e- = M^1- + EA

electron affinity energy type

defined as exothermic, but may be endothermic

• some alkali metals and all noble gases’ electron affinities are endothermic

more energy released = more electron affinity

• more negative the number, larger the EA

Electron affinity trend

increases across a period- halogens have the highest for any period

metal characteristics

ductile, malleable, shiny, conduct heat/electricity, most oxides basic and ionic, form cations in solutions, lose electrons in reactions to become oxidized

nonmetal characteristics

brittle, dull, nonreflective, electrical/thermal insulators, most oxides acidic and molecular, form anions and polyatomic anions, gain electrons in reactions to become reduced

metallic character

how closely an element’s properties match ideal metallic properties (malleability/ductility, conductors, easy to ionize)

metallic character trend

decreases from left to right and increases down a column