CHEM 105 Chapter 3 Quiz Flashcards
how do reoccurring patterns in the physical and chemical properties of elements aid scientists?
predicting undiscovered elements and understanding elemental behavior
periodic
exhibit a repeating pattern
density (increases/decreases) as you move down a column
increases
main-group elements
groups 1-2 (s orbitals), 13-18 (p orbitals)
transition elements
groups 3-12 (d orbitals)
inner transition elements
f orbitals
the orientation of an electron spin is quantized, meaning
it can be only in one direction or its opposite (specifically, +/- 1/2)
by convention, and up arrow is spin (?) and a down arrow is spin (?)
up (+ 1/2); down (- 1/2)
spins must ? in an orbital, meaning the two electrons in the orbital must have ? signs
cancel; opposite
paired spins are
diamagnetic
the Pauli Exclusion Principle
no two electrons in an atom may have the same set of four quantum numbers
the sublevels (s, p, d, f) in each principal energy shell or hydrogen, or other single electron systems, all have (the same/different) energy
the same
orbitals with the same energy are said to be
degenerate
for multi-electron atoms, the sublevels are ? as a result of?
split; charge interaction, shielding, and penetration
for a given n value, the lower the value of the ? quantum number (the orbital quantum number), the (more/less) energy the sublevel has
l; less
aufbau principle
electrons enter atomic orbitals from lowest energy to highest as related to the periodic table; nature loves low energy (because it is more stable)
Coulomb’s Law
describes the attractions and repulsions between charged particles:
for like charges, the potential energy is positive and decreases as the particles get farther apart as r increases
for opposite charges, the potential energy is negative and becomes more negative as the particles get closer together
the strength of the interaction increases as the size of the charges increases
E is proportional to (q1q2)/r
in a multi-electron atom, each electron experiences
both an attraction to the protons in the nucleus and the repulsion by the other electrons in the atom
shielding
the repulsions by other electrons in an atom cause an electron to experience a net reduction in attraction to the nucleus (the electron does not experience the full attraction by protons in the nucleus because the other electrons in the atom are interfering/blocking the attractive forces)
Zeff
effective nuclear charge: the total amount of attraction that an electron feels for the nucleus’s protons
the degree of penetration is related to
the orbital’s radial distribution function
attraction between nucleus protons and orbiting electrons is related to
the orbital type the electron occupies (s > p > d > f)
what causes the energies of sublevels in the same principal level to not be degenerate?
penetration
the energy separations between one set of orbitals and the next become slightly smaller beyond the ? orbital, which can cause ?
4s; the ordering to vary among elements, causing variations in the electron configurations of the transition metals and their ions
Hund’s rule
when filling orbitals that have the same energy (degenerate), place on electron in each orbital before completing pairs
core (inner) e-s
electrons in lower-energy shells
valence e-s
the e-s in all the sublevels with the highest principal energy shell (level)
one of the most important factors in the way an atom behaves, both chemically and physically, is ? because ? and ?
the number of valence electrons
- it is the valence e-s that participate in bonding and are lost to make cations
- the valence shell is where e-s are added to make anions
the periodic table is divisible into four blocks corresponding to
the filling of the four quantum sublevels (s, p, d, and f)
the group number of a main group element is equal to
the number of valence e-s for that element
the row number of a main group element is equal to
the highest principal quantum number of that element
because of level splitting, the 4s sublevel is lower in energy than the 3d sublevel; therefore the ? fills before the ?
4s orbital, 3d orbital
some transition metals have irregular e- configs (anomalies) in which
the ns only partially fills before the (n-1)d or doesn’t fill at all (ex: Cr, Mo, Cu, Ag); these anomalies happen when an s electron jumps to a d orbital to created a half-filled or completely filled sublevel
noble gases
- 8 valence electrons (except for He, which only has 2)
2. especially nonreactive (because the electron configurations are especially stable)
the metals (general)
make up majority of the elements in the periodic table
- alkali metals
- alkaline earth metals
- transition and inner transition metals
- p-block metals
alkali metals
- have one more electron than the previous noble gas and occupy the first column
- in their reactions, the alkali metals lose one e- and the resulting e- config is the same as that of a noble gas
alkaline earth metals
- they have two or more electrons than the previous noble gas and occupy the second column
- in their reactions, they lose two e-s and the resulting e- config is the same as that of a noble gas
transition and inner transition metals
- located in the d-block area of the PT
2. in chemical reactions, they will lose e-s from s and then d orbitals to form cations
p-block metals
- located in the p-block area (left-hand side of the metalloids) of the PT
- in chemical reactions, they will lose e-s from the s and p orbitals to form cations
metalloids
- located in the p-block area of the PT between the metal and nonmetal elements
- can exhibit metallic or nonmetallic behaviors in chemical reactions
- can either lose e-s from p and then s orbitals to form cations or gain e-s in their p orbitals to form anions
nonmetals
- located in the upper right-hand side of the PT in the p-block area
- in chemical reactions, nonmetal elements will gain electrons in the p orbitals, resulting in their ions having the same electron configuration as a noble gas at the same end of their period/row
- form anions
halogens
- nonmetals
- have one fewer e- than the next noble gas
- in their reactions with metals, they tend to gain an e- and attain the e- config of the next noble gas, forming an anion
- in their reactions w/ nonmetals, they tend to share e-s with the other nonmetal so that each attains the e- config of a noble gas
metals form ? and nonmetals form ?
cations, anions
Zeff (increases/decreases) as you go across a period and (increases/decreases) as you descend down a column
increases; decreases
Zeff =
Z (nuclear charge) - S (the number of e-s in lower energy levels)
atomic radius (increases/decreases) across a period because
decreases; adding e-s to the same valence shell, Zeff increases across a period, valence shell held closer
atomic radius (increases/decreases) down a group because
increases; valence shell farther distance from nuclear; Zeff decreases down a group
paramagnetism
electron configurations that result in unpaired electrons means that the atom or ion will have a net magnetic field; will be attracted to a magnetic field
diamagnetism
electron configurations that result in all paired electrons mean that the atom or ion will have no magnetic field; will be slightly repelled by a magnetic field
cation radius is (bigger/smaller) than its corresponding atom radius because
smaller; the loss of e-s results in the remaining e-s in the atom experiencing a larger Zeff than the neutral atom; traversing down a group increases the (n-1) level, causing the cations to get larger; traversing to the right across a period increases the effective nuclear charge for isoelectronic cations, causing the cations to get smaller
anion radius is (larger/smaller) than its corresponding atom radius because
larger; when atoms form anions, e-s are added to the valence shell, and these new valence e-s experience a smaller Zeff than the old valence e-s; traversing down a group increases the n level, causing the anions to get larger; traversing to the right across a period increases the Zeff for isoelectronic anions, causing the anions to get smaller
in general, ion size (increases/decreases) down the column
increases
cations are smaller than anions except
Rb^+ and Cs^+, which are bigger than or the same size as F^- and O^2-
isoelectronic
same electron configuration
ionization energy
the minimum energy needed to remove an electron from an atom or ion in the gas phase; an endothermic process
first ionization energy
energy to remove e- from neutral atom; all atoms have this
M(g) + IE1 > M^1+(g) + 1e-
second ionization energy
energy to remove an e- from a 1+ ion
M^1+(g) + IE2 > M^2+(g) + 1e-
the larger the effective nuclear charge on the electron to be removed, the (more/less) energy it takes to remove it
more
the farther the electron most probably is from the nucleus, the (more/less) energy it takes to remove
less
first ionization energies trend
first IE decreases down the group; first IE increases across the period
exceptions to the first ionization energy trend are ? and are usually a result of ?
2A to 3A and 5A to 6A; the type of orbital (s, p, d, f) and its shielding ability, repulsion factors associated with the electrons occupying degenerate orbitals (i.e., p orbitals)
trends in second and successive ionization energies
they depend on the number of valence electrons an element has; the removal of each successive electron costs more energy; regular increase in energy for each successive valence e-; large increase in energy when core e-s are removed
electron affinity
the energy associated with the addition of an electron to the valence shell of an atom that is in the gas phase; defined as exothermic but may actually be endothermic
M(g) + 1e- > M^-1(g) + EA
the more energy that is released (the more ? the number), the ? the EA
negative; larger
electron affinities general trend for main group elements
EA increases as across a period (becomes more negative from left to right); Halogens have the highest EA for any period
characteristics of metals
malleable, ductile; shiny, lustrous, reflect light; conduct heat and electricity; most oxides basic and ionic; form cations in solution; lose electrons in reactions - oxidized
characteristics of nonmetals
brittle in solid state; dull, nonreflective, solid surface; electrical and thermal insulators; most oxides acidic and molecular; form anions and polyatomic anions; gain e-s in reactions - reduced
metallic character
how closely an element’s properties match the ideal properties of a metal
metallic character trends
decreases L to R across a period; increases down the column
metallic elements
ionization energy decreasing down the column; very low IEs
