Periodic Properties of the Elements

Question 1

eleectron config of copper (Cu)

Answer 1

1. Ar of Cu = 29
   2.Cu:[Ar]3d ^10, 4s ^1

Question 2

Aufbau principle

Answer 2

The typical order of filling orbitals is:
- 1s→2s→2p→3s→3p→4s→3d→4p→5s→4d andsoon.
- filling the orbitals according to the order of increasing energy.

1𝑠^2= (2 electrons)
2s ^2= (2 electrons)
2p^6= (6 electrons)
3s ^2= (2 electrons)
3p ^6= (6 electrons)
4s ^2= (2 electrons)
3d orbital holds 10 electrons.

= 30
Normally, you’d expect copper’s electron configuration to end with
4s^2 ,3𝑑^9
- But atoms like to be more stable,
- having a full 3d orbital (w 10 e-) = copper extra stability.
- So, 1 e- from the 4s orbital moves to the 3d orbital.

[Ar]4s ^1, 3d ^10 = shorthand version

Question 3

understanding electron con fig usuing H+

Answer 3

• The “1” in 1s^1 indicates the first energy level of atom + correlates to the first row of the periodic table.
• The “s” in 1s^1 indicates that H+ has electrons in the “s” orbital, which correlates to the s block of the periodic table.
• The “1” in 1s^1 indicates that hydrogen has one electron in the “s” orbital.

Question 4

T or F
All ionization energies are endothermic and why

Answer 4

• True
• It requires energy to remove any electron from an atom or ion.
• Successive ionization energies become more endothermic whether valence or core electrons are being removed.

Question 5

Electrons and energy levels

Answer 5

• Electrons in an atom are not randomly placed, they are found in specific regions around the nucleus called energy levels/ shells
• each energy level can hold a certain number of electrons, and they represent different distances from the nucleus
• The energy levels are labeled by principal quantum number, represented by n

1. n=1 is the first energy level (closest to the nucleus, lowest energy).
2. n=2 is the second energy level (a little farther from the nucleus, higher energy).
3. n=3, n=4, etc., represent energy levels that are progressively farther from the nucleus and have more energy.

Question 6

why do energy levels matter

Answer 6

• Lower levels (like n=1) are closer to the nucleus + are more stable because the nucleus pulls the electrons strongly.

= Higher levels (like n=7) are farther away, + electrons in these levels are less strongly attracted to the nucleus, so they have more energy and can be more easily removed or excited.

Question 7

for energy levels
The fifth row of the periodic table corresponds to what energy level in the atom?

Answer 7

• 1st row (Hydrogen and Helium) corresponds to n=1
• 2nd row corresponds to n=2.
• 3rd row corresponds to n=3.
• 4th row corresponds to n=4.
• 5th row corresponds to n=5.
• atoms in the fifth row have their outermost electrons in the n = 5 energy level.

Question 8

Ionization energy

Answer 8

• first ionization energies is to **decrease from top to bottom of columns **
• BEACUSE atoms get larger (they have more electron shells),
• electrons held loosely ,so the outer electrons are farther from the nucleus
• less tightly held.
• it’s easier to remove them/less energy to remove one,
• first ionization energies increase from left to right of row because as you move across a period,
• atoms get smaller and have a stronger pull from the nucleus (due to more protons).
• electrons are held more tightly, making it harder to remove them/more energy needed to remove
• The positively charged protons pull the electrons toward the nucleus.

Question 9

diamagnetic vs paramagnetic

Answer 9

Diamagnetic
- A substance is diamagnetic if all electrons are paired.
- do not have unpaired electrons, so they are not attracted to a magnetic field. In fact, they are slightly repelled by a magnetic field.
- the magnetic effects of paired electrons cancel each other out because paired electrons have opposite spins. This causes the atom to have no net magnetic moment.
- He, Ne, Zn

Paramagnetic
- if it has one or more unpaired electrons.
- attracted to a magnetic field because the unpaired electrons create a net magnetic moment
- Unpaired electrons generate a magnetic moment because their spins don’t cancel out.
- so O2, Fe

Question 10

True or false
Electrons closest to nuclues like 1s and 2s electrons shield the nucleus for electrons that are farther away.

Answer 10

• shielding reduces the ZEFF felt by the outer electrons, making them feel less of the nucleus’s pull.
• Inner electrons block some of the nuclear pull on the outer electrons, which is why the outer electrons feel a lower effective nuclear charge.

Question 11

Anions are _____ than their neutral parent atom, and cations are _____ than their neutral parent atom.

Answer 11

• larger
• smaller

Question 12

true or false
Elements have a series of ionization energies for removing the first and additional electrons from the atom. Each ionization energy increases for successive electrons being removed

Answer 12

• After each electron is removed = fewer electrons left to shield the nucleus, so the remaining electrons feel a stronger attraction to the nucleus.
• As more electrons are removed, the +ve charge of the nucleus becomes more dominant, making it harder to pull away additional electrons.
• each successive ionization energy increases because it takes more energy to remove an electron that is more tightly held by the nucleus after earlier ones have been removed.

Question 13

The outermost electron of boron experiences a lower effective nuclear charge than carbon. Why?

Answer 13

• The nucleus of an atom contains P++
• The more protons in the nucleus, the stronger the +Ve pull on the e- (negative charge).
  . - carbon’s extra proton creates a stronger pull on its outer electrons.
• the outermost electron in Carbon experiences a higher Zeff than in boron
  . - boron’s outer electron feels a weaker pull because its nucleus has fewer protons.

.

Question 14

how many electrosn can each orital hold

Answer 14

• s orbitals can hold 2 electrons.
• p orbitals can hold 6 electrons.
• d orbitals can hold 10 electrons

Question 15

Why Chromium Doesn’t Follow the Usual Pattern of electron config

Answer 15

• because having half-filled d orbitals (3d⁵) is more stable than having a configuration where the d orbital isn’t half-filled (like 3d⁴).

Half-filled orbitals (like 3d⁵) are more stable because of the way electrons interact with each other. In a half-filled set of orbitals, the electrons are spread out evenly, which reduces repulsion between them.

In chromium’s case, 1 electron from the 4s orbital moves to the 3d orbital so that the 3d orbital becomes half-filled (3d⁵), creating a more stable configuration.

Question 17

what does half filled mean

Answer 17

• means that exactly half of the available spots for electrons in that orbital are filled.
• A d orbital can hold a maximum of 10 electrons
• When a d orbital has 5 electrons (like in 3d⁵), it is half-filled, because 5 is half of 10
• more stable because electrons spread out evenly in the available sub-orbitals, reducing their repulsion (since like charges repel each other).

Question 18

atomic radius trends

Answer 18

• atomic radius is that the size increases moving from top to bottom in the periodic table because, a new e- shell is added = outermost electrons are further from the nucleus + more shielding
• Moving across the periodic table, from left to right, the size decreases beacuse the number of protons in the nucleus increases. This stronger +ve charge pulls the electrons closer to the nucleus.

Question 19

Which sequence lists atoms in order of increasing atomic radius (smallest to largest)?

Li < Mg < Al < Ge

O < C < Si < Al

H < He < F < Ne

B < C < N < O

Answer 19

O < C < Si < Al

herefore, O < C < Si < Al lists the elements in order of increasing atomic radius.

Li < Mg < Al < Ge is incorrect because aluminum is smaller than magnesium, and germanium is smaller than both.

B < C < N < O is incorrect because the atomic radius decreases from left to right.

H < He < F < Ne is incorrect because helium is smaller than hydrogen, and neon is smaller than fluorin; atomic radius decreases from left to right.

Question 20

For the outermost electron of an atom, effective nuclear charge XXXXX as you go from left to right on the periodic table. This is because electron shielding XXXXX due to electrons being added to the same energy level as the outermost electron.

Answer 20

• increases
• stays constant

Question 21

electron afinities

Answer 21

Electron affinity becomes more negative going from left to right but does not have a clear trend going from top to bottom for most columns.

Question 22

Identify the largest ion.

Li+
F−
Cl−
Na+

Answer 22

Chlorine is larger than fluorine, so their anions will follow the same trend, which means Cl− is larger than F−.

• (Cl) is below (F) in the periodic table.
• As you move down a group in the periodic table, atoms get larger.
• each new row (or period) adds another energy level (shell) of electrons.
• F has 9 electrons and 2 energy levels (shells).
• Cl has 17 electrons and 3 energy levels (shells).
• Since Cl has more energy levels, it has a larger atomic radius than F
• (Cl) already has more energy levels than fluorine (F), so when they both gain one electron, their anions (Cl⁻ and F⁻) still follow the same size trend as their neutral atoms.
• Cl⁻ will have 3 energy levels, while F⁻ only has 2 energy levels, so Cl⁻ is naturally larger.

Question 23

mass vs atomic number

Answer 23

• Ar is number of P
• MN is number of P+N