Chemistry Midterm 3 Flashcards

Units 8.D-11.C (Quantum Model of the Atom, Periodictiy and Ionic Bonding, Covalent Bonding, Molecular Shape and Bonding Theories, Liquids and Solids)

Wave Function

Describes the quantum state of an electron in an atom, specifically finding the probability of the electron at a certain point in space

Shape of s orbital (ℒ = 0)

Spherically symmetric and lack directionality, meaning the probability of finding an electron is uniform on each sphere around the nucleus

Nodes

Points or surfaces where the wave function is zero; number of total nodes is n - 1

Radial Nodes

Spherical surfaces where the wave function is zero; number of radial nodes is n - ℒ - 1

Shape of 2s orbital (ℒ = 0)

Has a sphere with one radial node (the wave function changes from positive to negative)

Angular Nodes

Planar and cone-shaped nodes where the wave function is zero

Shape of p orbital (ℒ = 1)

Dumbbell-shaped; has one angular node in between the two lobes of the dumbbell

Shape of d orbital (ℒ = 2)

Mostly clover-leafed shaped; has two angular nodes 

Shape of f orbital (ℒ = 3)

Visually complicated; has three angular nodes 

Trend between energy levels (n), orbital size, and energy

As the value of n increases, the orbital size, energy, radial nodes generally increase

Trend within a shell for orbital energy and size

Within a shell, orbital energy and size increase with ℒ (s —> p —> d —> f)

Trend between number of angular nodes and ℒ

The number of angular nodes in an orbital is equal to ℒ (that’s why it’s called angular momentum quantum number)

Atomic Orbital Energy Diagram

Graphical representation of occupancies and energies of atomic orbitals for each subshell (s, p, d, f)

Orientation of Atomic Orbital Energy Diagrams

• Energy is the vertical axis

• Orbitals are represented with horizontal lines and labels

• Electrons are shown through arrows pointing up (m_s = +1/2) and down (m_s = -1/2)

Aufbau Principle

Electrons fill the lowest-energy orbitals (ground state) first before occupying higher-energy orbitals

Hund’s Rule

Every orbital within a subshell is singely occupied with one electron before any orbital is double-occupied (single-occupied orbitals must have the same spin)

Electron Configuration

Lists the arrangement of electrons within each orbital for the subshells. For ground state electron configurations, it is listed in their lowest possible energy states

• Ex. For a neutral sulfur atom: 1s²2s²2p⁶3s²3p⁴

Degenerate

Orbitals within a subshell

Highest-energy occupied subshell for s and p-block elements

Period number

Highest-energy occupied subshell for d-block elements

Period number - 1

Highest-energy occupied subshell for f-block elements

Period number - 2

Abbreviated Electron Configurations

Shorthand way of writing electron configurations using noble gases

• Ex. For a Ca atom, it would be [Ar]4s²

Valence Shell

Highest-energy occupied shell

Valence Electrons

Electrons that are within the highest-energy occupied shell involved in bonding and electron-transfer processes

• Ex. For a F atom (1s²2s²2p⁵), there are seven valence electrons in n=2 shell (2s²2p⁵)

Core Shell

Lower-energy occupied shells

Core Electrons

Electrons within lower-energy occupied shells that shield valence electrons from full positive charge of nucleus; aren’t involved in chemical reactions usually

• Ex. For a O atom ([He]2s²2p⁴), the core electrons are within the [He}

Subshells relating to Cations for Main Group (s and p-block) Elements

Main group (s and p-block) elements in groups 1,2, and 13 form cations to become more stable, specifically to the closest noble gas

Subshells relating to Cations for Transition (d-block) Elements

Transition (d-block) elements form cations by removing electrons from the highest principal energy level (n) first

• Ex. Fe has [Ar]4s²3d⁶, but when it’s Fe²⁺, it loses the 4s² electrons to become [Ar]3d⁶

Subshells relating to Anions for Main Group (s and p-block) Elements

Main group (s and p-block) elements in groups 14-17 are nonmetals that want to complete their valence shell and become more stable by configuring to nearest noble gas

Size of atoms are dictated by-

1. Electronic structure of the atom or ion

2. Interactions between the positive nucleus and the negative electrons

Interactions between the positive nucleus and the negative electrons follow-

Electrostatic Principles

Electrostatic Principles

1. Oppositely-charged particles attract each other

2. Like-charged particles repel each other

3. As charges increase, so does the forces of attraction or repulsion

4. As two charged bodies get closer to each other, so does the forces of attraction or repulsion

Effective Nuclear Charge (Z_eff)

The net positive charge from the nucleus that a valence electron experiences

Z_eff for Valence Electrons

They are lower than the actual nuclear charge (Z), thus Z_eff < Z due to the shielding effect by core electrons and increased distance

Formula for Z_eff

Z_eff = Z - S

• Z: Nuclear Charge (number of protons)

• S: Shielding Constant

Slater’s Rules

Helps determine the shielding constant (S) for a valence electron in an atom using only the number of core and valence electrons

Formula for Shielding Constant (S)

S = 0.85(N_core) + 0.35(N_val - 1)

• N_core = number of core electrons

• N_val = number of valence electrons

Atomic Radius

Measure of the size of the atom based on the distance between the nucleus and the outermost electron shell

Atomic Radius Trends

• Increases down a group

  • Principal energy levels (n) are added, thus this expands the atomic size

• Decreases from left to right across a period

  • As the number of protons increase, they pull in the electrons more tightly, thus force of attraction increases

• Applies to cations and anions as well

Atomic Radius for Cations

• All cations have smaller radii than their corresponding neutral atoms

  • With less electrons, this lowers the shielding effect, thus the remaining electrons are pulled in more tightly towards the nucleus

Atomic Radius for Anions

• All anions have larger radii than their corresponding neutral atoms

  • The addition of electrons increases electron-electron repulsion and shielding, thus the atom expands to so there is more distance between the electrons

Ionization Energy

Energy required to remove an electron from the gaseous atom to produce a gaseous cation and a free electron

Ionization Energy Trends

• Decreases going down a group

  • The outermost electrons are farther from the nucleus, thus it requires less energy to remove an electron due to less forces of attraction

• Increases going left to right across a period

  • As the number of protons increase, this leads to the electrons being pulled more tightly, thus it’s harder to remove them

Exceptions to Ionization Energy

• Outermost valence electron is in a subshell by itself

• Elements containing one set of paired electrons in the valence p subshell (Group 15 vs. Group 16)

Higher Ionization Energies:

• Involves the removal of a second electron, third electron, etc.

• Z_eff increases with each electron lost, thus the attraction between the electrons and the nucleus become stronger, thus that’s the reason for higher ionization energy as electrons are removed

Electron Affinity

Energy change when an electron is added to a gaseous neutral atom to create a gaseous anion

Electron Affinity Values

• Positive value means an increase in energy needed, thus an endothermic process

• Negative value means an decrease in energy needed, thus an exothermic process

Electron Affinity Trends

• Becomes less exothermic (less negative) moving down a group due to less energy released when an electron is added farther away from the nucleus

• Becomes more negative left to right across a period (ignoring noble gases) due to stronger pull from nucleus for incoming electron, thus more energy is released when added

Electrostatic Attraction

Forces of attraction between oppositely-charged species

Ionic Bond

The electrostatic attraction between oppositely-charged ions

Ionic Bond Characteristics

• Metal donates electron to the non-metal that accepts it to achieve the electronic structure of the nearest noble gas

• Compounds contain an infinite 3D lattice of cations and anions (formula unit)

Lattice Energy (△H_L)

The energy released when gas-phase ions combine to form 1 mole of a solid ionic compound (can be reversed where a solid ionic compound become gas-phase ions but it’s endothermic)

Lattice Energy Trends

• A decrease moving down a group because as atomic size increases, there’s a weaker electrostatic attraction between the positive and negative ions

• An increase from left to right across a period due to ionic charge increasing, which creates a stronger electrostatic attraction with nucleus and the electrons

• Trends are based on Coulomb’s Law

Coulomb’s Law

• As the charge increases, the forces of attraction increases (direct relationship)- relates to period trend

• As the distance between the center of two ions increase, the forces of attraction decreases (inverse relationship)- relates to group trend

Born-Haber Cycle

Since lattice energies can’t be directly measured, this cycle uses Hess’s Law, ionization energy, electron affinity, and the energy of other processes to calculate △H_L

Covalent Bonds

The sharing of electrons between atoms, generally between two nonmetals

Octet Rule

Achieving a stable noble gas configuration (low-energy state) with eight total valence electrons

Why Eight Electrons?

For p-block elements, this correpsonds to a full atomic valence shell (ns²np²) which hold eight electrons

Lewis Structure

Diagram of a molecule or polyatomic ion that shows the shared pairs or bonds through straight lines and the unshared or lone pairs through dots

• Total number of electrons in a Lewis structure must equal the total number of valence electrons of the atoms in the molecule

Number of Bonds

• Single Bond (weakest)

  • Holds 2 electrons

• Double Bond

  • Holds 4 electrons

• Triple Bond (strongest)

  • Holds 6 electrons

Steps for Drawing Lewis Structures

1. G

2. g

3. g

4. g

5. g

6. g

7. g

Formal Charges

Based on the count of valence electrons to obtain a theoretical charge

Formal Charge Formula

(# of valence electrons) - (# of non-bonding electrons) - (1/2 of bonding electrons)

Electronegativity

Measure of capacity or ability of an atom to attract shared electrons in a covalent bond

Electronegativity Trends

• Increases left to right across a period due to effective nuclear charge increasing, thus more protons being added to the same energy level brings the electron cloud closer for better attraction

• Decreases going down a group since electrons get further away from the nucleus and the core electrons shielding the valence electrons to weaken the nucleus pull

Characteristics of Highly-Electronegative Elements

• Have greater electron density, leading to more unequal sharing of electrons (more polar)

• Form fewer bonds due to preference of creating one strong bond rather than sharing electrons in multiple bonds 

• Placed on periphery of Lewis structure to reduce electron-electron repulsion (which destabilizes the atom)

Correlation between Electronegativity and Z_eff, Ionization Energy, and Electron Affinity

High Electronegativity means high Z_eff, high ionization energy, highly negative electron affinity (except for exceptions)

Mulliken Scale

1/2(IE + EA)

• The bigger the values, the tighter the atom holds electrons