Brainscape's Knowledge GenomeTM


Top Flashcards (Certified)
All Flashcards

Fall 2017 > CHEM 105 Chapter 3 Quiz > Flashcards

CHEM 105 Chapter 3 Quiz Flashcards

1
Q

how do reoccurring patterns in the physical and chemical properties of elements aid scientists?

A

predicting undiscovered elements and understanding elemental behavior

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

periodic

A

exhibit a repeating pattern

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

density (increases/decreases) as you move down a column

A

increases

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

main-group elements

A

groups 1-2 (s orbitals), 13-18 (p orbitals)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

transition elements

A

groups 3-12 (d orbitals)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

inner transition elements

A

f orbitals

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

the orientation of an electron spin is quantized, meaning

A

it can be only in one direction or its opposite (specifically, +/- 1/2)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

by convention, and up arrow is spin (?) and a down arrow is spin (?)

A

up (+ 1/2); down (- 1/2)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

spins must ? in an orbital, meaning the two electrons in the orbital must have ? signs

A

cancel; opposite

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

paired spins are

A

diamagnetic

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

the Pauli Exclusion Principle

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

the sublevels (s, p, d, f) in each principal energy shell or hydrogen, or other single electron systems, all have (the same/different) energy

A

the same

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

orbitals with the same energy are said to be

A

degenerate

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

for multi-electron atoms, the sublevels are ? as a result of?

A

split; charge interaction, shielding, and penetration

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

for a given n value, the lower the value of the ? quantum number (the orbital quantum number), the (more/less) energy the sublevel has

A

l; less

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

aufbau principle

A

electrons enter atomic orbitals from lowest energy to highest as related to the periodic table; nature loves low energy (because it is more stable)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

Coulomb’s Law

A

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

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

in a multi-electron atom, each electron experiences

A

both an attraction to the protons in the nucleus and the repulsion by the other electrons in the atom

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

shielding

A

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)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

Zeff

A

effective nuclear charge: the total amount of attraction that an electron feels for the nucleus’s protons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

the degree of penetration is related to

A

the orbital’s radial distribution function

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

attraction between nucleus protons and orbiting electrons is related to

A

the orbital type the electron occupies (s > p > d > f)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

what causes the energies of sublevels in the same principal level to not be degenerate?

A

penetration

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

the energy separations between one set of orbitals and the next become slightly smaller beyond the ? orbital, which can cause ?

A

4s; the ordering to vary among elements, causing variations in the electron configurations of the transition metals and their ions

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Hund's rule
when filling orbitals that have the same energy (degenerate), place on electron in each orbital before completing pairs
26
core (inner) e-s
electrons in lower-energy shells
27
valence e-s
the e-s in all the sublevels with the highest principal energy shell (level)
28
one of the most important factors in the way an atom behaves, both chemically and physically, is ? because ? and ?
the number of valence electrons 1. it is the valence e-s that participate in bonding and are lost to make cations 2. the valence shell is where e-s are added to make anions
29
the periodic table is divisible into four blocks corresponding to
the filling of the four quantum sublevels (s, p, d, and f)
30
the group number of a main group element is equal to
the number of valence e-s for that element
31
the row number of a main group element is equal to
the highest principal quantum number of that element
32
because of level splitting, the 4s sublevel is lower in energy than the 3d sublevel; therefore the ? fills before the ?
4s orbital, 3d orbital
33
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
34
noble gases
1. 8 valence electrons (except for He, which only has 2) | 2. especially nonreactive (because the electron configurations are especially stable)
35
the metals (general)
make up majority of the elements in the periodic table 1. alkali metals 2. alkaline earth metals 3. transition and inner transition metals 4. p-block metals
36
alkali metals
1. have one more electron than the previous noble gas and occupy the first column 2. in their reactions, the alkali metals lose one e- and the resulting e- config is the same as that of a noble gas
37
alkaline earth metals
1. they have two or more electrons than the previous noble gas and occupy the second column 2. in their reactions, they lose two e-s and the resulting e- config is the same as that of a noble gas
38
transition and inner transition metals
1. 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
39
p-block metals
1. located in the p-block area (left-hand side of the metalloids) of the PT 2. in chemical reactions, they will lose e-s from the s and p orbitals to form cations
40
metalloids
1. located in the p-block area of the PT between the metal and nonmetal elements 2. can exhibit metallic or nonmetallic behaviors in chemical reactions 3. 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
41
nonmetals
1. located in the upper right-hand side of the PT in the p-block area 2. 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 3. form anions
42
halogens
1. nonmetals 2. have one fewer e- than the next noble gas 3. in their reactions with metals, they tend to gain an e- and attain the e- config of the next noble gas, forming an anion 4. 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
43
metals form ? and nonmetals form ?
cations, anions
44
Zeff (increases/decreases) as you go across a period and (increases/decreases) as you descend down a column
increases; decreases
45
Zeff =
Z (nuclear charge) - S (the number of e-s in lower energy levels)
46
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
47
atomic radius (increases/decreases) down a group because
increases; valence shell farther distance from nuclear; Zeff decreases down a group
48
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
49
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
50
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
51
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
52
in general, ion size (increases/decreases) down the column
increases
53
cations are smaller than anions except
Rb^+ and Cs^+, which are bigger than or the same size as F^- and O^2-
54
isoelectronic
same electron configuration
55
ionization energy
the minimum energy needed to remove an electron from an atom or ion in the gas phase; an endothermic process
56
first ionization energy
energy to remove e- from neutral atom; all atoms have this | M(g) + IE1 > M^1+(g) + 1e-
57
second ionization energy
energy to remove an e- from a 1+ ion | M^1+(g) + IE2 > M^2+(g) + 1e-
58
the larger the effective nuclear charge on the electron to be removed, the (more/less) energy it takes to remove it
more
59
the farther the electron most probably is from the nucleus, the (more/less) energy it takes to remove
less
60
first ionization energies trend
first IE decreases down the group; first IE increases across the period
61
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)
62
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
63
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
64
the more energy that is released (the more ? the number), the ? the EA
negative; larger
65
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
66
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
67
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
68
metallic character
how closely an element's properties match the ideal properties of a metal
69
metallic character trends
decreases L to R across a period; increases down the column
70
metallic elements
ionization energy decreasing down the column; very low IEs

Fall 2017 (24 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2025 Bold Learning Solutions. Terms and Conditions