Transcript Chemical Periodicity Chart

Chemical Periodicity Chart
Practice Problem
Questions and Answers
Atomic Radius
• P. 178 #16, 22
• 16: How does atomic size change within
groups and across periods?
– Increases down the groups, decreases left-to-right
across periods.
• 22: Arrange in order of decreasing size: S, Cl,
Al, Na. Is there a pattern?
– Na, Al, S, Cl
– This is a period-based trend (left-to-right, n=3).
Atomic Radius
• P.181 #36
– A: Which element has a larger radius: Na or Li?
• Na
– B: Which element has a larger radius: Sr or Mg?
• Sr
– C: Which element has a larger radius: C or Ge?
• Ge
– D: Which element has a larger radius: O or Se?
• Se
Atomic Radius
• P. 182 #50
• Why does fluorine have a smaller atomic
radius than oxygen and chlorine?
– It’s further to the right in oxygen’s period, it’s
higher up than chlorine.
– In other words, “stronger nucleus” than oxygen,
fewer electrons than chlorine.
Ionization Energy
• P. 178 #17-18
• 17: When do ions form?
– When electrons are added or removed.
• 18: What happens to first ionization energy
within groups and across periods?
– Increases left-to-right across periods, decreases
down groups.
Ionization Energy
• P.178 #23
• A: Which element has the larger first
ionization energy: Na, K?
– Na
• B: Which element has the larger first
ionization energy: Mg, P?
–P
Ionization Energy
• P.181 #37, 38
• 37: Explain the difference between first and second
ionization energy:
– First i.e. = energy to remove one electron.
– Second i.e. = energy to remove a second electron
• 38: Which element has a greater first i.e.?
– Li, B
• B
– Mg, Sr
• Mg
– Cs, Al
• Al
Ionization Energy
• P. 181 #39
• Arrange the groups of elements in order of
increasing ionization energy:
– Be, Mg, Sr
• Sr, Mg, Be
– Bi, Cs, Ba
• Cs, Ba, Bi
– Na, Al, S
• Na, Al, S
Ionization Energy
• P.181 #40
• Why is there a large increase between the first
and second ionization energies of the alkali
metals?
– After removing the first electron, the second
electron is in a lower (closer) energy level (lower n
number).
Ionization Energy
• P. 182: 51, 55
• 51: Would you expect metals or nonmetals in the
same period to have higher i.e.?
– Nonmetals – they’re further right (“stronger nuclei”)
• 55: Which equation represents the first ionization
of an alkali metal atom?
–
–
–
–
A: Cl  Cl+ + eB: Ca  Ca+ + eC: K  K+ + eD: H  H+ + e-
Ionization Energy
• P. 182: 51, 55
• 51: Would you expect metals or nonmetals in the
same period to have higher i.e.?
– Nonmetals – they’re further right (“stronger nuclei”)
• 55: Which equation represents the first ionization
of an alkali metal atom?
–
–
–
–
A: Cl  Cl+ + eB: Ca  Ca+ + eC: K  K+ + eD: H  H+ + e-
Ionization Energy
• P.182 #58
• Why is there a large jump between the second
and third ionization energies of magnesium?
Why is there a large jump between the third
and fourth ionization energies of aluminum?
– Those last electrons are in closer energy shells
(lower n number).
Ionic Size
• P. 178 #19
• Compare the size of ions to the size of their
neutral forms.
– Cations lose electrons, become positively charged,
get smaller.
– Anions gain electrons, become negatively charged,
get larger.
Ionic Size
• P. 181 #41, 42
• 41: How does the ionic radius of a typical metal compare
with its atomic radius?
– Metals tend to lose electrons so their ionic radii get smaller.
• 42: Which particle has a larger radius in each atom/ion
pair?
– Na, Na+
• Na
– S, S2• S2-
– I, I• I-
– Al, Al3• Al
Ionic Size
• P. 182: #52
• In each pair, which ion is larger?
– Ca2+, Mg2+
• Ca2+
– Cl-, P3• P3-
– Cu+, Cu2+
• Cu+
Ionic Size
• P. 182 #59
• The bar graph shows the relationship between
atomic and ionic radii for Group 1A elements.
A: Describe the trend in atomic radius. B:
Explain the difference between ionic and
atomic radius size?
– A: Radius increases as you go down a group.
– B: Ions are smaller due to fewer electrons than in
the neutral atom (atomic radius).
Ionic Size
• P.183 #64, 65
• 64: The Mg2+ and Na+ ions each have ten electrons.
Which is smaller and why?
– Mg2+ is smaller because though it has ten electrons just
like Na+, it has more protons. They pull “harder” on the
electrons.
• 65: How do you expect the radii of S2-, Cl-, K+, Ca2+, and
Sc3+ to vary – they have the same total electrons as the
noble gas Argon. What about for O2-, F-, Na+, Mg2+, and
Al3+, which is the same as Neon?
– Radius decreases from left to right across a period in both
cases. Though electron # is the same, proton number goes
up.
Ionic Size
• P. 183 #68
• Atoms and ions with the same number of
electrons are isoelectronic.
– Write the symbol for a cation and anion that are
isoelectronic with Krypton:
• Br-, Rb+, Se2-, As3-, Sr2+ (each have 36 electrons)
– Can you have an isoelectric cation and anion in
the same period?
• No, cations lose electrons but anions (higher overall
number of electrons) gain them.
Electronegativity
• P. 178 #20
• How does electronegativity vary within groups
and across periods?
– Increases across period left-to-right.
– Decreases down groups.
Electronegativity
• P.181 #43
• A: Which element has a higher electronegativity value: Cl,
F?
– F
• B: Which element has a higher electronegativity value: C,
N?
– N
• C: Which element has a higher electronegativity value: Mg,
Ne?
– Mg [Ne does not react]
• D: Which element has a higher electronegativity value: As,
Ca?
– As
Electronegativity
• P.181 #44
• Why are noble gases not given
electronegativity values?
– Electronegativity only applies in compounds.
Noble gases don’t react and form compounds.