Chapter 08: Periodic Properties of the Elements

Question 1

Periodic law

Answer 1

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

Question 2

Electron configuration

Answer 2

Shorthand for showing particular orbitals that electrons occupy for that atom

e.g. 1s¹ for Hydrogen

Question 3

Pauli exclusion principle

Answer 3

No two electrons in an atom can have the same four quantum numbers

Each orbital can have a max of 2 electrons, with opposing spins

Question 4

Degenerate

Answer 4

Having the same energy

Question 5

Coulomb’s law

Answer 5

Like charges:
Potential energy is positive
Decreases as particles are farther apart
Repel each other

Opposite charges:
Potential energy is negative
Becomes more negative as particles are closer
Opposites attract

Magnitude of interaction increases as charges of particles increase
i.e. 1- and 2+ are more strongly attracted than 1- and 1+

Question 6

Shielding

Answer 6

Screening or repulsion of one electron by other electrons

Question 7

Effective nuclear charge

Answer 7

Z_eff = Z - S
= actual nuclear charge (atomic #) - charge screened (by core electrons)

Increases as you go right on periodic table
(Which is why radius gets smaller to the right)

Question 8

Penetration

Answer 8

When an outer electron penetrates into lower energy region, it experiences a greater nuclear charge, and thus a lower energy

Question 9

Orbital diagram

Answer 9

A diagram that represents the orbitals occupied by electrons in a given atom

Boxes for arrows of electrons, with orbital name under each box

e.g. 1s with a box above

Question 10

Aufbau principle

Answer 10

Lower energy orbitals fill before higher energy orbitals to minimize the energy of the atom

Question 11

Hund’s rule

Answer 11

Electrons first occuply orbitals of equal energy singly with parallel spins

Question 12

Transition elements with irregular electron configurations (10)

Answer 12

1. Cr (Chromium) 24
2. Cu (Copper) 29
3. Nb (Niobium) 41
4. Mo (Molybdenur) 42
5. Ru (Rutherium) 44
6. Rh (Rhodium) 45
7. Pd (Palladium) 46
8. Ag (Silver) 47
9. Pt (Platinum) 78
10. Au (Gold) 79

Question 13

Paramagnetic

Answer 13

Electron configurations with unpaired electrons

Have a net magnetic field

Attracted to a magnetic field

Question 14

Diamagnetic

Answer 14

All electrons are paired in electron configuration

No magnetic field of its own

Slightly repelled by a magnetic field

Question 15

Isoelectronic

Answer 15

Elements that have the same number of electrons & same ground-state electron configuration

Question 16

Van der Waals radius

Answer 16

Nonbonding radius of an atom

Question 17

Covalent radius

Answer 17

Radius of an atom when it’s bonded

Question 18

Atomic radius

Answer 18

Average radius of a bonded atom based on measuring large numbers of elements and compounds

Question 19

Periodic trends in atomic radius for transition metals

Answer 19

Remains roughly constant across the d block

Question 20

Ionic radius

Answer 20

Cation is smaller than parent atom (e.g. Li⁺ v. Li)

Anion is always larger than parent atom (e.g. Cl^- v. Cl)

Question 21

Cation v. anion radius size

Answer 21

Cations < anions

Except Rb⁺ and Cs⁺ are larger than F^- and O^2-

Question 22

Ionic radii of isoelectronic species

Answer 22

Larger positive = smaller cation

Larger negative = larger anion

Question 23

Ionization energy

Answer 23

IE₁ + X (g) → X⁺ (g) + e^-

The minimum energy required to remove an electron from a gaseous atom or ion

IE₁ < IE₂ < IE₃

(Each subsequent electron removal requires more energy, due to increasing positive charge of atom)

(Large increase in IE when core electrons are removed)

Question 24

Periodic trends in first ionization energy

Answer 24

The larger the Z_eff, the more energy it takes to remove it.
*IE₁ increases to the right
**Except: Group 2A to 3A & Group 5A to 6A

The farther the electron is from the nucleus, the less energy it takes to remove it
*IE₁ decreases toward the bottom

Question 25

Electron affinity

Answer 25

X (g) + e^- → X^- (g) + EA

The energy released (-kJ/mol) when a neutral gaseous atom gains an electron

The more energy released when electron is gained, the more negative the EA (-kJ/mol)

Question 26

Periodic trends in electron affinity (6)

Answer 26

1. EA becomes less negative going down Group 1A
2. 3rd Period becomes more negative from 2nd Period
3. EA becomes more negative going right
4. Group 5A has less negative EA than expected (extra electron pairs up)
5. Groups 2A/8A have less negative EA (extra electron goes into new level/sublevel)
6. Group 7A has most negative EA

Question 27

Periodic trends in metallic character

Answer 27

Metallic character increases going left and down

Question 28

Periodic trends in Group 1A (alkali metals) (7)

Answer 28

1. IE increases up the group (as with most groups)
2. Generally low IEs (want to lose electron)
3. EA becomes more negative down the group
4. Melting point increases up the group (unusual)
5. Generally low melting points
6. Density increases down the group (except K)
7. Reactivity increases down the group

Question 29

Periodic trends in Group 7A (halogens) (4)

Answer 29

1. IE increases up the group (as with most groups)
2. EAs are high (want to gain electrons)
3. React with H to form binary acids
4. Mass increases more than volume

Question 30

Periodic trends in Group 8A (noble gases) (3)

Answer 30

1. IE increases up the group (as with most groups)
2. IEs are high (don’t want to lose electrons)
3. Only Kr & Xe form compounds, usually with F

Question 31

Metals in Group 1A vs. 1B (2)

Answer 31

1. Similar outer electron configurations: ns¹
2. Properties different because different IEs
   1A: lower IE, more reactive (want to lose ns¹)
   1B: higher IE, less reactive

Question 32

Diagonal relationship

Answer 32

1. Li & Mg
2. Be & Al (Al is a metal like 2A)
3. B & Si (metalloids)

Similar because of charge density (ion charge / volume)

Question 33

Properties of oxides across a period (3)

Answer 33

1. Metals/Groups 1A/2A form basic oxides
2. Nonmetals form acidic oxides
3. Al form amphoteric oxides (both basic/acidic)