Chem exam 4

Traveling wave

A wave that moves through space, transferring energy from one location to another (like light or sound waves)

Standing waves

A wave that remains fixed in position and is formed by the interference of two waves traveling in opposite directions

Energy oscillates in place instead of propagating like in electorn wave patterns in atoms

Equation relating the speed of light, wavelength, and frequency

c = λ x v

Line spectra

A series of distinct wavelengths (lines) of light emitted or absorbed by atoms, corresponding to specific electronic transitions between energy levels

Continuous spectra

A seamless range of all wavelengths or colors of light with no gaps between them produced by things like incandescent solids

Photoelectric effect

It observed that electrons could be ejected from a metal's surface if the FREQUENCY was greater than some threshold frequency

However, the KE of the ejected electrons DID NOT depend on the brightness of the light but increased with increasing frequency of the light

Classical wave theory

A wave's energy depends on its intensity (which depends on its amplitude), not frequency----now a new light model was needed

Explain how the photoelectric effect supports the particle nature of light

A wave increases in energy when intensity increases, BUT a particle increases in energy when its frequency increases.

Ejecting electrons takes a certain amount of energy that was not satisfied with light with increased intensity but it was with increased frequency that demands light behave like a particle

New model of light

Light behaves like a wave and a particle

Blackbody radiation

The light emitted by an object that absorbs all radiation and re-emits energy based only on its temperature

An example is a metal oven that can be heated to high temperatures and is a conventient and ideal emitter for study

Describe how blackbody radiation, the photoelectric effect, and line spectra led to thequantum model of the atom.

Classical physics thought energy output would increase without limit at short wavelengths but experiments showed that it peaked and decreased (known as the UV catastrope).

Max Planck proposed that energy is quantized, it can only be emitted/absorbed in quanta with energy proportional to its frequency E = hv

This idea introduced the concept of energy quantization and quantum theory

Compare emission and absorption in the Bohr model—what determines if a photon isemitted or absorbed

When energy is released, light is emitted

When energy is absorbed, a photon is absorbed

Wave-particle duality

The idea that light and matter exhibit wave and particle like behavior

Significant because it shows that

Ionic compound properties

Ionic solids are rigid and brittle crystalline structures

Poor conductors of electricity

Poor conductors of electricity

Excellent conductors of heat and electricity in solution or melted

Covalent compound properties

Electrically neutral with a weaker attraction between them

They have lower melting and boiling points

Likely water INSOLUBLE

Poor conductors of heat and electricity at any state

Form softer solids than ionic compounds

Many are gases or liquids at room temperature

Hund's rule

The orbitals within a subshell must first be filled spin up to maximize the number of unpaired electrons before pairing electrons

Pauli Exclusion Principle

No two electrons in the same atom can have exactly the same set of all the 4 quantum numbers

What is the relationship between valence-electron configuration and periodic trends?

The number of valence electrons increases as you go from left to right

Johann Balmer

He derived an equation based on the four visible line spectra for hydrogen

Quantum mechanics

Describing matter using quantization of energy, wave-particle duality, and the Heisenberg uncertainty principle

Louis de Broglie

He extended the wave-particle duality of light that Einstein used to resolve the photoelectric-effect paradox to material particles

He modeled electrons as circular standing waves required to have only integer wavelengths

Heisenberg uncertainty principle

It is fundamentally impossible to determine simultaneously and exactly both the position and the momentum of a particle

Principle quantum number (n)

Describes the location (shell) of the orbital

-Allowed values of n = 1, 2, 3, ….

-Orbitals with the same n value are in the same shell

Defines the general size/energy of orbital

Higher n values are larger, higher in energy and farther from nucleus

Angular momentum quantum number (l)

Defines the shape of the orbital

-Allowed values depend on the shell: l = 0, 1, 2, …, n-1

-Higher the l value, higher the angular momentum of an electron in the orbital

Orbitals with the same l value in same subshell

Magnetic quantum number (ml)

Defines the z component of the angula rmomentum

-Allowed values of ml = -l, …, 0, …., l for a total of 2l + 1 values for a given value of l

Spin quantum number (ms)

Defines the electron spin

-It has nothing ot do with position or the Schrodinger equation

-It describes the behavior of an electron's magnetism

-Allowed values of ms = ± 1/2 with spin up +1/2 and spin down -1/2

Valence electrons

The electrons in the outermost shell of a ground-state atom; determines how an element reacts

Core electrons

The electrons occupying the inner shell orbitals

Valence shell

Outermost shell of electrons in a ground-state atom

Periodic properties that govern the chemical behavior of elements

1. Size (radius) of atoms and ions

2. Ionization energies

3. Electron affinities

Covalent radius

Defined as 1/2 the distance between the nuclei of two identical atoms when they are joined by a covalent bond

The size of the atom (and its covalent radius) ____ down a group

Increases

The size of the atom (and its covalent radius) ___ across a period

Decreases

Effective nuclear charge

Zeff is the pull exerted on a specific electron by the nucleus, taking into account any electron-electron repulsions. Zeff = Z - shielding

Isoelectronic species

Atoms and ions that have the same electron configuration

First Ionization energy

The amount of energy required to remove the most loosely bound electron from a gaseous atom in its ground state

Electron affinity

The energy change for the process of adding an electron to a gaseous atom to form an anion

Bond length

The internuclear distance at the lowest potential energy

Polar covalent bond

It forms when the electrons are pulled toward one atom causing a partial negative chrage and a corresponding partial positive charge on the other atom. This separation of charge (dipole) acts as tiny electric field

Electronegativity

The measure of the tendency of an atom to attract shared electrons to itself

Lewis symbol

A depiction of an atom or ion that represents the valence electrons with dots around the elemental symbol

Free radicals

Molecules that contain odd number of electrons

Electron-deficient molecules

seldom seen but in Group 2 Be compounds and Group 13 B compound

Lewis structure

A diagram of a molecule or an ion that shows lone pairs and bonding pairs of electrons

Pure covalent bond

Occurs whenever the same type of atoms are involved

Isaac Newton

White light composed of a mixture of colors streams of high speed "corpuscular" particles.

Christiaan Huygens

Light as waves that can be felected and refracted

Thomas Young

Interference of light during slits suggest waves, not corpuscles

James Clerk Maxwell

Theory of electromagnetic radiation; classical thermodyamics vs classical mechancis (Newtonian motion)

Modern understanding of light

It is an electromagnetic radiation wave and a particle

Rayleigh-Jeans Law

The maxima in the blackbdoy curves λ max shift to shorter wavelengths as the temperature increases

Max planck

-Planck derived a theoretical expression for blackbody radaition that fit the experimental observations exactly within experimental error over all wavelengths.

-The atoms vibrated at an increasing frequencies or defeated wavelength as the temperature increased

-Planck assumed that the vibrating atoms required quantized energies which he could not explain.

E = hv where n = 1, 2, 3, 4

Niels Bohr Model

-Bohr used Planck's quantization of energy and Einstein's photon explanation with his own assumptions that emission/absoprtion of photons resulted from electrons moving between orbits in an atom

-As a result, he coudl reproduce Rydberg's empirical equation from fundamental constants

Electron configuration subshell labels include:

-Principal quantum shell number (n)

-Number of electrons in subshell

-Letter of subshell (l)

What is the periodic trend for electronegativity

EA increases from left to right across a period, bottom to top within a group

If n = 1, what can l be?

l = 0

If n = 4, what can l be?

l = 0, 1, 2, 3

If l = 2, what is the lowest value n can be? What values can ml be?

n = 3

ml = -2, -1, 0, 1, 2

Davisson and Germer

They demonstrated that electrons could be diffracted, with interference patterns only possible with waves

Erwin Schrodinger

-Built upon de Broglie's relation to develop a fundamental equation for quantum mechanics that determines the energies of electrons: Ĥψ = Eψ

Ĥ = Hamiltonian operator (set of mathematical functions) 

ψ = wavefunction of the particle (a mathematical equation)

E = total energy

Born Interpretation

The square of the magnitude of the wave function (|ψ|²) is the probabilityi of finding the electron in a region of space

What are quantum numbers used to characterize?

The atomic orbitals

When can 2 electrons occupy the same orbital?

IF and only IF they have opposite spins

Ionization processes are ____ and IE values are always ___

Endothermic, positive

Hypervalent molecules

often see, especially in the third period or below with the central atom having up to 12 electrons

Inert pair effect as an explanation for unexpected ion charges

• The inert pair effect is the reluctance of the outermost s-electrons in heavier p-block elements to participate in bonding or ionization, leading to unexpectedly low oxidation states. 

  • The inert pair effect explains why heavier p-block elements form lower-than-expected positive ion charges —because their ns^2 electrons resist ionization and remain nonbonding. 

Difference between electron affinity and electronegativity

• Electronegativity is a calculated value that has no dimensions and is arbitrarily set with a 4.0 maximum

• Electron affinity is a measurable quantity of energy (kJ/mol) released or absorbed when a gas-phase atom gains an electron 

Octet Rule for Lewis Structures

• This rule guides the construction of Lewis Structures. Main group atoms tend to form enough covalent bonds to obtain eight valence electrons around them (valence filled).

EN difference = 0

Pure covalent

EN difference between 0.4-1.8

Polar covalentE

EN difference greater than 1.8

Ionic

Identifying elements exhibiting exceptions to the Octet Rule

• Hydrogen can only form one bond because its valence has only 2 electrons

• Free radicals: molecules that contain an odd number of electrons

• Electron-deficient molecule: Group 2 Be compounds, Group 13 B compounds

• Hypervalent molecules: elements in the third period and below, the central atom may have up to 12 electrons

Number of bonds expected for p-block elements based on valence electrons

• For p-block elements, the number of bonds an atom tends to form equals the number of unpaired electrons in its valence shell, which allows it to reach an octet

• Group 14 → 4 bonds

• Group 15 → 3, Group 16 →, Group 17 → 1 bond

Describe the wave nature of light

• Light behaves as an electromagnetic wave consisting of oscillating electric/magnetic fields that move perpendicular to each other and to the direction of travel

• These waves can travel through a vacuum and are characterized by wavelength, frequency, amplitude, and speed of light

Describe the particle nature of light and the findings about the photoelectric effect

• The photoelectric effect showed that light behaves as particles (photons) each carrying energy proportional to its frequency, and only photons with enough energy can eject electrons from a metal surface, proving that energy transfer from light to matter is quantized 

Describe the Bohr model of the hydrogen atom

• Bohr combined Planck’s quantization of energy and Einstein’s photon theory with the idea that electrons move in specific quantized orbits around the nucleus

• Each orbit/energy level is described by a quantum number (n), and the energy of each level increases with distance from the nucleus and the levels get closer together at higher energies

• Electrons naturally occupy the lowest available energy level

• When an electron absorbs energy, it moves to a higher orbit, and when it falls back down, it emits a photon whose energy equals the difference between the two levels 

How is the periodicity of elemental properties related to valence electron configurations?

• Elements in the same group have the same valence electron configuration, which determines their chemical behavior

• As the atomic number increases, the valence configurations repeat periodically, producing the periodic patterns in properties such as size, ionization energy, and electron affinity. 

How is the Pauli exclusion principle related to quantum numbers and/or atomic orbitals?

• It is related to orbitals because each orbital can hold a maximum of 2 electrons, and they must have opposite spins, one with +½ and -½, ensuring that each electron has a unique quantum state. 

How is Hund’s rule related to quantum numbers and/or atomic orbitals?

• It maximizes the number of unpaired electrons and minimizes electron–electron repulsion, leading to greater stability. It applies to the way electrons are arranged in orbital diagrams within the same subshell

How did the discovery of blackbody radiation shed light on atomic theory?

• Classical physics, like Rayleigh-Jeans law, failed to explain why radiation didn’t increase infinitely at short wavelengths (UV catastrophe)

• Planck proposed that energy is emitted in quantized packets, and this idea of quantized energy became the first step to understanding that electrons could only occupy discrete energy levels in atoms

• Blackbody radiation showed that energy is not continuous but quantized, which laid the foundation for later atomic models.

How did the discovery of the photoelectric effect shed light on atomic theory?

• Classical wave theory predicted that any bright light should eject electrons, which did not happen

• Einstein used Planck’s quantization to propose that light consists of photons, and the photon’s energy has to be high enough to knock the electron loose 

• The Photoelectric effect confirmed that energy transfer between light and matter occurs in quantized amounts, supporting the idea that atomic electrons absorb or emit energy in discrete packets (wave-particle duality)

How did the discovery of line spectra shed light on atomic theory?

• It was observed that excited gases emit line spectra, not continuous spectra, with only specific wavelengths of light appearing for each element. Each line corresponded to a photon with a specific energy difference between electron levels. Balmer derived an equation describing visible hydrogen lines, and Rydberg extended it to all hydrogen spectral lines.

  • Line spectra shed light on atomic theory by providing experimental proof of quantized energy levels in atoms, saying that electrons can only exist in specific energy states

How did the Bohr Model shed light on atomic theory?

• Combined Planck's and Einstein’s ideas, but proposed that electrons move in quantized orbits around the nucleus

  • When electrons jump between orbits, photons are emitted/absorbed with energy equal to the difference between levels

  • Bohr’s model explained the hydrogen line spectrum and introduced the concept of energy levels and electron transitions.

    • The Bohr model contributed to atomic theory by demonstrating that atomic structure and light emission are interconnected through quantized electron orbits, marking a significant step toward modern quantum mechanics.

Define electronegativity and assess the polarity of covalent Electronegativity is the measure of the tendency of an atom to attract shared electrons to itself 

• Electronegativity is the measure of the tendency of an atom to attract shared electrons to itself 

If the atoms are the same, they share electrons equally and are pure covalent

If the atoms are different, electrons are shared unequally and are polar covalent

• The difference in electronegativity between two atoms indicates the type and polarity of the bond. Bonds with the largest electronegativity value have greater polarity as electrons are drawn more strongly toward the electronegative atom, producing partial charges and a dipole 

Explain the formation of cations, anions, and ionic compounds

• Metals have low ionization energies and lose valence electrons readily to form CATIONS

• Nonmetals have relatively high electron affinities and gain electrons to form ANIONS

• Ionic compounds consist of electrically balanced combinations of anions and cations. The formula for an ionic compound lists the cation, then the anion, with ion ratios that have the simplest combination of charges adding up to zero net charge