Chapter 08: Periodic Properties of the Elements Flashcards

1
Q

Periodic law

A

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

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

Electron configuration

A

Shorthand for showing particular orbitals that electrons occupy for that atom

e.g. 1s1 for Hydrogen

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

Pauli exclusion principle

A

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

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

Degenerate

A

Having the same energy

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

Coulomb’s law

A

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+

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

Shielding

A

Screening or repulsion of one electron by other electrons

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

Effective nuclear charge

A

Zeff = 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)

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

Penetration

A

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

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

Orbital diagram

A

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

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

Aufbau principle

A

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

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

Hund’s rule

A

Electrons first occuply orbitals of equal energy singly with parallel spins

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

Transition elements with irregular electron configurations (10)

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

Paramagnetic

A

Electron configurations with unpaired electrons

Have a net magnetic field

Attracted to a magnetic field

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

Diamagnetic

A

All electrons are paired in electron configuration

No magnetic field of its own

Slightly repelled by a magnetic field

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

Isoelectronic

A

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

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

Van der Waals radius

A

Nonbonding radius of an atom

17
Q

Covalent radius

A

Radius of an atom when it’s bonded

18
Q

Atomic radius

A

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

19
Q

Periodic trends in atomic radius for transition metals

A

Remains roughly constant across the d block

20
Q

Ionic radius

A

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

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

21
Q

Cation v. anion radius size

A

Cations < anions

Except Rb+ and Cs+ are larger than F- and O2-

22
Q

Ionic radii of isoelectronic species

A

Larger positive = smaller cation

Larger negative = larger anion

23
Q

Ionization energy

A

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)

24
Q

Periodic trends in first ionization energy

A

The larger the Zeff, 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

25
Electron affinity
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)
26
Periodic trends in electron affinity (6)
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
27
Periodic trends in metallic character
Metallic character **increases** going **left** and **down**
28
Periodic trends in Group 1A (alkali metals) (7)
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
29
Periodic trends in Group 7A (halogens) (4)
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
30
Periodic trends in Group 8A (noble gases) (3)
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**
31
Metals in Group 1A vs. 1B (2)
1. Similar outer electron configurations: ns**1** 2. Properties different because different IEs 1A: lower IE, more reactive (want to lose ns**1**) 1B: higher IE, less reactive
32
Diagonal relationship
1. Li & Mg 2. Be & Al (Al is a metal like 2A) 3. B & Si (metalloids) Similar because of **charge** **density** (ion charge / volume)
33
Properties of oxides across a period (3)
1. **Metals**/Groups **1A**/**2A** form **basic** oxides 2. **Nonmetals** form **acidic** oxides 3. **Al** form **amphoteric** oxides (**both** basic/acidic)