Chapter 08: Periodic Properties of the Elements Flashcards Preview

Gen Chem 01 > Chapter 08: Periodic Properties of the Elements > Flashcards

Flashcards in Chapter 08: Periodic Properties of the Elements Deck (33):
1

Periodic law

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

2

Electron configuration

Shorthand for showing particular orbitals that electrons occupy for that atom

e.g. 1s1 for Hydrogen

3

Pauli exclusion principle

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

4

Degenerate

Having the same energy

5

Coulomb's law

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+

6

Shielding

Screening or repulsion of one electron by other electrons

7

Effective nuclear charge

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)

8

Penetration

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

9

Orbital diagram

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

10

Aufbau principle

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

11

Hund's rule

Electrons first occuply orbitals of equal energy singly with parallel spins

12

Transition elements with irregular electron configurations (10)

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

13

Paramagnetic

Electron configurations with unpaired electrons

Have a net magnetic field

Attracted to a magnetic field

14

Diamagnetic

All electrons are paired in electron configuration

No magnetic field of its own

Slightly repelled by a magnetic field

15

Isoelectronic

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

16

Van der Waals radius

Nonbonding radius of an atom

17

Covalent radius

Radius of an atom when it's bonded

18

Atomic radius

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

19

Periodic trends in atomic radius for transition metals

Remains roughly constant across the d block

20

Ionic radius

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

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

21

Cation v. anion radius size

Cations < anions

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

22

Ionic radii of isoelectronic species

Larger positive = smaller cation

Larger negative = larger anion

23

Ionization energy

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

Periodic trends in first ionization energy

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: ns1

2. Properties different because different IEs
1A: lower IE, more reactive (want to lose ns1)
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)