Periodicity & Ionic Bonding

Question 1

Electron Spin

Answer 1

A form of angular momentum of electrons

Question 2

Electromagnetic Theory

Answer 2

If a charged particle spins on its axis (like the Earth), it produces a magnetic field

Question 3

Electron Spin Quantum Number (ms)

Answer 3

• (ms) is used to describe the direction of the electron’s spin
• It can take on one of two values, +½ and –½

Question 4

Stern & Gerlach

Answer 4

Did an experiment showing electron spin did exist

Question 5

Degenerate

Answer 5

Electron orbitals having the same energy levels (ex. 2p & 2s) on n = 2

Question 6

Electron Configuration

Answer 6

The distribution of the electrons among the atomic orbitals

Question 7

Orbital Diagrams

Answer 7

Use diagrams such as boxes and arrows to show the electron configuration

Question 8

Pauli Exclusion Principle

Answer 8

States that no two electrons in an atom can NOT have the same four quantum numbers (they can have the same values of n, ℓ, and mℓ, but then the values of ms must be different), therefore, each orbital can hold only two electrons, one of each spin

Question 9

Shielding Effect

Answer 9

Since n = 2 orbitals are larger than 1s orbitals, the two 1s electrons help block or “shield” the n = 2 electrons from the attractive forces of the nucleus

Question 10

2s vs. 2p Orbitals

Answer 10

• 2p electrons are closer to the nucleus than the 2s electrons
• Electron density near the nucleus is greater for the 2s electrons
• 2s orbital is more “penetrating” (less shielded) than the 2p orbital, thus the 2s orbital is lower in energy

Question 11

Orbital Penetration

Answer 11

Orbitals with lower values of ℓ are more penetrating – so “s” orbitals are more penetrating than “p” orbitals, which are more penetrating then “d” orbitals, etc

Question 12

Hund’s Rule

Answer 12

States that “most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins”

Question 13

Diamagnetism vs Paramagnetism

Answer 13

Diamagnetism - Electrons are paired
Paramagnetism - Electrons are unpaired

Question 14

Aufbau Principle

Answer 14

States that as protons are added to the nucleus to build up elements, electrons are added the same way to atomic orbitals

Question 15

Condensed Electron Configuration

Answer 15

Inner (core) electrons are given as the
noble gas element that precedes the element being considered

Question 16

Transition Metals

Answer 16

Elements that contain an unfilled “d” subshell

Question 17

F-Block Elements or Lanthanides

Answer 17

Elements that contain unfilled “f” subshells

Question 18

Lothar Meyer & Dmitri Mendeleev

Answer 18

• Created independent periodic tables
• These tables (especially Mendeleev’s) were great improvements over previous ones, mostly in the area of atomic properties
• Mendeleev left holes in his periodic table which he claimed were as-yet undiscovered elements

Question 19

Problems with the New Tables

Answer 19

• Although new tables worked well and allowed for valuable predictions, there were still some questions
• Ex. The atomic mass of argon (39.95 amu) is larger than that of potassium (39.10 amu), so some early tables had potassium as element 18 and argon as element 19
• However, that did not match the known properties of other members of their columns (e.g. argon is nothing like lithium or sodium)

Question 20

Rutherford

Answer 20

Made a few estimates of the number of positive charges in the nuclei of a few elements, but did not follow up on his work

Question 21

Henry Moseley

Answer 21

• Correlated the frequency of X-rays emitted by elements to something he called the atomic number
• Few exceptions (such as Ar / K), this atomic number increased in the same order as the atomic mass
• Allowed the correct order for argon and potassium on the periodic table

Question 22

Classification of the Elements

Answer 22

All of the elements can be classified into one of a number of categories:
1. Representative elements
2. Noble gases
3. Transition elements (metals)
4. Lanthanides
5. Actinides

Question 23

Representative Elements

Answer 23

• Found in groups 1A – 7A (1, 2, 13-17) in the periodic table
• Have unfilled s or p subshells of the highest principle quantum number (ex. C has an unfilled 2p subshell)
• Rb has an unfilled 5s subshell
• Mg has an unfilled 3p subshell (filled 3s)

Question 24

Valence Electrons

Answer 24

• All elements in a column have the same valence electron configuration, this leads to similar chemical properties
• Are the outermost electrons, and
  they are the only ones that are involved in chemical bonding

Question 25

Valence Electrons for Representative Elements

Answer 25

For representative elements, it is the outermost s and p electrons that are valence (those of the highest occupied shell – i.e. largest value of n) – e.g.) Mg:1s^22s^22p^63s^2 – 2 valence electrons (2 valence from 3s) – e.g.) Cl:[Ne]3s^23p^5 – 7 valence electrons (2 + 5 = 7) (2 valence from 3s 5 valence from 3p)

Question 26

Transition Metals

Answer 26

- Occupy groups 1B – 8B on the periodic table (in the middle) - Often group 2B (Zn, Cd, Hg) are not counted as transition metals - Have incomplete d subshells - All share some chemical properties (such as their compounds being almost always coloured)

Question 27

Physical Property

Answer 27

A property that can be observed without transforming the substance into another substance – Size – Mass – Density

Question 28

Effective Nuclear Charge

Answer 28

Is the number of elementary charges an electron experiences by the nucleus. It is denoted by Zeff.

Question 29

Effective Nuclear Charge Formulas

Answer 29

Zeff = Z – S - Z - Is the atomic number - S - Is the shielding (screening constant) (0 < S < Z)

Question 30

Electron Shielding

Answer 30

- Can be tested experimentally by measuring the ionization energies of atoms (the energy needed to remove an electron from the atom) - Electrons are only shielded by electrons in the same or lower shells, and shielding is more effective by electrons in a lower shell

Question 31

Physicist John Slate

Answer 31

Came up with some rules used to calculate the shielding of any electron in an atom

Question 32

Slater’s Rules

Answer 32

The shielding constant for an electron is calculated as the sum of: 1) An electron in the same group (s & p count as one group) contributes 0.35 2a) If it is an s or p electron, then every electron with a principal quantum number (n) one less than the electron contributes 0.85, every electron with n two less or below contributes 1. 2b) If it is a d or f electron, then every electron inside it contributes 1. This includes all electrons with n less than the electron as well as s and p electrons with the same value of n.

Question 33

Atomic Radius

Answer 33

- It is difficult to determine the size of an atom exactly - Remember the electron density of an atom extends far away from the nucleus - Generally we think of the atomic size as the volume containing about 90% of the electron density - We can also define atomic radius based on the experimental distance between atoms

Question 34

Trends in Atomic Radii

Answer 34

- Down columns = increase of n = bigger orbitals = bigger atomic radii - Atomic radii depends on attraction between valence electrons and nucleus - Across a row, nuclear charge increases by +1 for each element; however the screening effect is not enough to compensate (since the electrons in a shell do not perfectly screen each other) - The effective nuclear charge on the outermost electrons must then increase, and the orbitals contract (become smaller), reducing the atomic radius

Question 35

Isoelectronic

Answer 35

Is when two elements or ions that have the same total number of electrons (same configuration)

Question 36

Ionic Radius

Answer 36

- Ionic radius is the radius of a cation or anion - Structure of an ionic solid depends on relative sizes of the cation(s) and anion(s) - When an atom is ionized, its radius changes

Question 37

Ionic Radii

Answer 37

- When an atom becomes a cation (electrons are removed), electron screening and repulsion decreases, orbitals contract making ionic radius smaller than the atomic radius - When an atom becomes an anion (electrons are added), electron repulsion increases, orbitals expand and ionic radius is larger than the atomic radius

Question 38

Ionization Energy

Answer 38

The minimum energy required to remove an electron from a gaseous atom in its ground state There are a whole series of ionization energies for multi-electron atoms: 1. X(g) + energy X🡪 +(g) + e- ΔH = 1st ionization energy (IE1) 2. X+(g) + energy X🡪 2+(g) + e- ΔH = 2nd “ “ (IE2) 3. X2+(g) + energy X🡪 3+(g) + e- ΔH = 3rd “ “ (IE3)

Question 39

True or False: Ionization energy is always an exothermic process?

Answer 39

FALSE: Ionization is always an endothermic process (requires energy) - When an electron is removed, it decreases the shielding on the other electrons, so every ionization energy is higher than the last: IE1 < IE2 < IE3 < ...

Question 40

Ionization Energy & Periodic Table

Answer 40

- An increased size corresponds to a decreased ionization energy - As we go across a period, ionization energies increase (higher Zeff) - As we go down a group, the ionization energies decrease due to the larger size of the atoms - Valence electrons further from nucleus = decrease in ionization energy

Question 41

Electron Affinity

Answer 41

The energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion X(g) + e- --> X-(g)

Question 42

Electron Affinity & Periodic Table

Answer 42

Metals have lower electron affinities than nonmetals and the halogens (group 7A) have the largest electron affinities

Question 43

Group 1A (Alkali Metals)

Answer 43

- Have very low ionization energies, so they almost always form compounds as +1 ions * Never found in nature as the metal * React with both oxygen and water (often violently): Dog’s Nuts of the Periodic Table 2 K(s) + 2 H2O(l) --> 2 KOH(aq) + H2(g)

Question 44

Group 2A (Alkaline Earth Metals)

Answer 44

- Tend to be less reactive than the group 1A metals * Generally form +2 ions * Can react with water to form a basic solution and liberate hydrogen: Ba(s) + 2 H2O(l) --> Ba(OH)2(aq) + H2(g)

Question 45

Acid / Base Properties of Oxides

Answer 45

- All the third period (Na - Cl) elements form oxides - Oxides can either be basic or acidic - Elements on the left side of the periodic table (more metallic) tend to be basic, and the nonmetals tend to be acidic - In between are oxides that are either neutral (CO, NO) or amphoteric (acts as both acid and base)

Question 46

Properties of Oxides

Answer 46

Basic: Na2O(s) + H2O(l) --> 2 NaOH (aq) Acidic: Cl2O7(l) + H2O(l) --> 2 HClO 4(aq) Amphoteric: Al2O3(s) + 6HCl(aq) --> 2AlCl3(aq) + 3H2O(l) Al2O3(s) + 2 NaOH(aq)+ 3 H2O(l) --> 2NaAl(OH) 4(aq)