Chapter 3: Atomic Structure: Explaining the Properties of Elements

Atomic radius (size) ________________ across a period or row.

decreases; Effective nuclear charge (Zeff) increases for valence shell electrons.

Zeff _______________ somewhat upon going down a column

decreases; valence orbitals larger and further from the nucleus

Electron removal is an _____________________ reaction

endothermic

Across a period, Zeff increases ______________ than shielding.

faster; results in contraction of atomic radius. Zeff = Zactual - shielding

The more negative EA, the ______________ the tendency to accept an electron and the ______________ the electron affinity.

greater; greater

The ______________ corresponds to the number of valence electrons

group number

Zeff ________________ upon moving across a row (period)

increases; nuclear charge (Z) increases faster than shielding

Heisenberg Uncertainty Principle

it is impossible to know precisely where an electron is and what path it follows. The act of determining the position of an electron causes it to move. by observing the electron, we change its energy/speed.

The ______________ the Zeff, the more energy required to remove an electron.

larger; The larger n, the less energy required.

The length of each "block" is the ______________ number of electrons the __________________ can hold

maximum; subshell

Removal of each successive electron costs ___________ energy.

more IE2 > IE1; endothermic; Zeff increases for valence electrons as electrons removed. Regular increase in energy for each successive valence electron.

The more _______________ EA, the more stable the resulting atom

negative

How many unpaired electrons would be predicted for the oxygen atom?

2. Full 1s shell, full 2s shell, but only 4 electrons in the 2p shell.

Speed of light

3.00 x 10^8 m/s

First Row Transition Metals

4s subshell fills before the 3d subshell (4s, 3d, 4p)

Second Row Transition Metals

5s subshell fills before 4d subshell

Third Row Transition Metals

6s subshell fills before 4f subshell (6s, 4f, 5d, 6p)

Wave Behavior of Electrons

A beam of electrons diffracts giving an interference pattern. Wave behavior and particles don't diffract. If electrons behave like particles, there should only be two bright spots on the target. However, an interference pattern is seen.

Wave Function

A function of the coordinates of an electron's position in three-dimensional space that describes the properties of the electron. Has no easily visualized meaning. Shows the probability of finding an electron in an atom; area/space.

Photon

A particle of electromagnetic radiation with no mass that carries a quantum of energy. A packet of energy. One photon = one electron emitted.

Wave

A periodic oscillation that transmits energy through space. It repeats in space and time.

Subshell

A subdivision of an energy level in an atom. They are divided into orbitals. Same orbital type in a level (all p orbitals in n=3 level).

Orbital Filling Diagram

A visual way to represent the arrangement of all the electrons in a particular atom. Follow Aufbau Principle

For many elements, EA is negative, but for some elements EA is positive.

Alkaline Earth Metals and Noble Gases

Schrodinger's Equation

Allows us to calculate the probability of finding an electron with a particular amount of energy at a particular location in the atom. There are multiple solutions. More than one orbital present in an atom - different energies and positional probability distributions. Each wave function or orbital characterized by a unique combination of three Quantum numbers - each is an integer. n, l, and m.

Aufbau Principle

An electron occupies the lowest-energy orbital that can receive it first.

Aufbau Principle Rule 3

An orbital can hold two electrons of opposite spin - Pauli Exclusion Principle

Aufbau Principle Rule 4

For degenerate orbitals (same energy), the lowest energy is attained when the number of electrons with the same spin is maximized - Hund's Rule. The spin up goes in the orbital first, only filling it half way.

Electromagnetic Spectrum

Frequency decreases from left to right. High frequency/short wavelength waves are most energetic. Gamma rays are most energetic. Radio waves are least energetic. Visible light is only a small portion of the EM spectrum.

Covalent Atomic Radium

Half the distance between the nuclei of two identical atoms covalently bonded together. d/2 is the radius.

Quantized

Having values restricted to whole-number multiples of a specific base value.

Amplitude

Height from the center line of the wave to the peak (or trough). Proportional to intensity.

Frequency (v)

How fast wave peaks are going by a fixed point. Units of hertz (Hz) or reciprocal seconds (s-1). Frequency and wavelength are inversely related.

Atomic Spectra Explained

In order to transition to a higher energy state (absorption), the electron must gain the correct amount of energy corresponding to the difference in energy between the final and initial states. Electrons in high energy states are unstable and tend to lose energy and transition to lower energy states. Energy is released as a photon of light. Each line on the emission spectrum corresponds to the difference in energy between two energy states.

Group Number

Indicates number of valence electrons in the neutral atom and the charge of the cation

Shielding

Inner shell electrons reduce the nucleus' attractive force felt by outer shell electrons. Reduce the attractive force between an electron and the nucleus. Valence electrons cannot shield other valence electrons.

Core Electrons

Inner shells. Filled orbitals. Provide shielding.

Emission Spectrum

Light emitted from an excited atom. Can be separated into individual wavelengths using a prism or diffraction grating. The spectra is non-continuous. Each element has a unique spectrum. Colored line on a black background shows emission.

Nature of Light

Light is a form of electromagnetic radiation. Travels as a wave. Has a magnetic field component and electric field component. Moves through space at the same constant speed.

Main Group Metals

Loss of an electron(s) from the highest energy occupied atomic orbital. Isoelectric to Noble Gas. K: 1s2 2s2 2p6 3s2 3p6 4s1 (lose) or [Ar] 4s1 K+: 1s2 2s2 2p6 3s2 3p6 or [Ar] 4s0 Same equation as the Noble Gas

Absorption Spectrum

Missing wavelengths when white light passes through a sample. Black line on a colored background shows absorption.

Anions

Negatively charged ions. Form from Non-metals. Fill valence electrons. Group # - 8 gives the charge on the anion.

Discrete Levels

Not continuous. Designated by n

EA2

Not favored. Adding 2nd electron is not ideal.

Emission

Occurs when an electron falls to a lower energy level. Excess energy is emitted as light.

Absorption

Occurs when light of a certain wavelength forces an electron to "jump" to a higher energy level.

Aufbau Principle Rule 2

Orbitals are filled in order of increasing energy, lower energy orbitals first. Lowest energy shells fill first (n=1 before n=2, etc). Lost energy orbitals in a shell fill first (s, p, d, f).

Ions

Positively and negatively charged atoms. Result from the loss or gain of electrons.

Cations

Positively charged ions. Form from Main Group Metals. Lose ALL valence electrons.

de Broglie

Proposed that electrons could ave wave-like character. Applicable to all matter. Wavelength is inversely proportional to velocity and mass (in kg).

Atomic radius (size) _______________ down a group or column.

Quantum number n increases down a group - valence electrons occupy larger orbitals.

Positional Probability Distribution

Region of space with a high likelihood for finding the electron. Squared.

Anomalous Electron Configuration

Results from the unusual stability of half-filled and full-filled subshells. Occur where the subshell energy differences are small. Half filed plus filled orbitals are more stable than partially filled plus filled orbitals. Ex: Chromium and Copper

Eight Valence Electrons

Results in a very unreactive atom (or ion) that is very stable. All Noble Gases have 8 valence electrons and are very stable. Helium is the exception with 2 valence electrons (very stable).

Isoelectric

Same number of electrons

Electron Configuration Notation

Show principle quantum number (n), subshell designation (l), and number of electrons in subshell.

Ground State Configuration

Take the number of electrons of the given element and subtract the number of electrons in the nearest previous Noble Gas. When writing the numbers and letters after the Noble Gas, you need to pick up where the Noble Gas letters and numbers left off. For example, the corresponding Noble Gas to Ge is Ar. The condensed equation for Ar is 3s2 3p6. However, Ge is in the 4th row and 2nd column in the p group, making the condensed equation [Ar] 4s2 3d10 4p2 because the Ar element does not hold the electrons in the shells listed but Ge does so we have to write in the extras after.

Effective Nuclear Charge (Zeff)

The actual nuclear charge experienced by an electron, defined as the charge of the nucleus minus the charge of the shielding electrons. Effective = Zactual - shielding Zeff: Attractive force experienced by electron Zactual: Full attractive force of the nucleus. Based on the number of protons. Li atom: 1s2 2s1 2s electron Nuclear charge (Z) = +3 Shielding by 1s electron: -2 Zeff = +1

Electron Configuration

The arrangement of electrons in the orbitals of an atom.

Electrostatic Attraction

The attraction between positive and negative charges. This is what holds an atom together. In multi-electron atoms, the electrons do not experience the total attractive force of the nucleus. Effective nuclear charge, Zeff.

Diffraction Through 2 Slits

The diffraction of light between two slits separated by a distance comparable to the wavelength results in an interference pattern of the diffracted waves. An interference pattern is a characteristic of all light waves.

Ionic Radius

The distance from the nucleus to the outer edge of the highest occupied atomic orbital. The radius of a cation is smaller than the radius of the parent atom. The radius of an anion is larger than the radius of the parent atom.

Condensed Electron Configuration

The electron configuration of the nearest noble-gas element of lower atomic number is represented by its chemical symbol in brackets

Photoelectric Effect

The emission of electrons from a metal when light shines on the metal. The energy of a photon of light is directly proportional to its frequency. inversely proportional to its wavelength.

Hydrogen Atom Emission Spectrum

The emission spectrum of the hydrogen atom can be modeled mathematically using the Rydberg equation, where n2 is greater than n1. Gives wavelength of emitted photon when a electron moves from a larger n value to a smaller n value. Discrete energy levels. Equation limited to Hydrogen atoms.

EA

The energy released due to the addition of an electron to a gas phase. H(g) + e- -> H-(g). Releases energy - EA is negative. Exceptions where EA is positive. Exothermic.

Electron Affinity

The energy released when a neutral atom in the gas phase gains an electron. Can be exothermic. Eea. Measure of the ability of an atom to accept an electron.

Ei

The energy required to remove an electron from a gas phase. H(g) -> H+(g) + e- . Requires energy - IE is positive. Endothermic.

Ionization Energy

The energy required to remove an electron from the ground state of a neutral atom in the gas phase. Measure of how strongly an atom holds its electrons.

Interference

The interaction between waves

Ground State

The lowest energy state of an atom

Binding Energy

The minimum energy required for an electron to be emitted. Also called the metal's work function. At frequencies above the threshold, the electron absorbs more energy than is necessary to escape. Excess energy become kinetic energy of the emitted electron.

Threshold Frequency

The minimum frequency of light required to produce the photoelectric effect. The greater the intensity, the more electrons emitted. When the frequency is less than the threshold, no electrons are emitted, even at very high intensity. At frequencies above the threshold, the electron absorbs more energy than is necessary to escape. Excess energy become kinetic energy of the emitted electron.

Quantum Mechanical Model

The modern description, primarily mathematical, of the behavior of electrons in atoms. Schrodinger (1926). Focuses on the wave-light properties of the electron. Wave equation, a mathematical model of the atom.

Non-metals

Typically gain electrons to form anions which are negatively charged. Cl-, O2-, P3-. Negative charge indicates how many electrons were gained.

Metals

Typically lose electrons to form cations which are positively charged. Na+, Ca2+, Al3+, Mn4+. Positive charge indicates how many electrons were lost.

Diffraction

When traveling EM waves encounter an obstacle or opening in a barrier that is about the same size as the wavelength, they bend around it.

Constructive Interference

When waves interact so that they add to make a larger wave. Waves are in-phase.

Destructive Interference

When waves interact so they cancel each other. Waves are out-of-phase.

Families

Elements in the same family have similar chemical and physical properties because they have the same number of valence electrons in the same kinds of orbitals.

Max Planck

Emission of electromagnetic radiation occurs in discrete units - not continuous. Tiny packets of energy called photons.

Endothermic

Energy absorbed

Classic Wave Theory

Energy from the light is transferred to electrons in the metal atoms. When sufficient energy is transferred the electrons can escape from the atoms. Does not match what is observed in the photoelectric effect.

Particle Nature of Light

Energy in light travels in bundles called photons which have a specific amount of energy and behaves like particles. Light treated purely as waves. Photoelectric effect best describes this.

Exothermic

Energy released

Quantum Numbers

Every atom has a unique set of quantum numbers, which describe which orbitals are occupied by electrons in an atom (2 electrons in each orbital).

Albert Einstein

Explained the photoelectric effect. Proposed that light behaves as a stream of particles called photons or quanta.

Exceptions to Ionization Energy Trend

Be & B; Mg & Al; N & O; P & S

Hydrogen Atom Energy

Calculation of energy levels in the hydrogen atom. Highest energy at n=1. Lowest energy at n=infinity. Electron completely removed from the atom at E=0. As n increases levels, spacing between levels becomes smaller.

Transition Metals

Can also form cations by loss of electrons. Lose valence shell s electrons before losing their d electrons. Different from Aufbau Principle. Fe: [Ar] 4s2 3d6 Fe2+: [Ar] 4s0 3d6 Fe3+: [Ar] 4s0 3d5

Non-metal Anions

Can be formed by accepting an electron(s) into the lowest energy level. Isoelectric to Noble Gas. Cl: 1s2 2s2 2p6 3s2 3p5 or [Ne] 3s2 3p5 Cl-: 1s2 2s2 2p6 3s2 3p6 (add) or [Ne] 3s2 3p6 Same equation as the Noble Gas

Aufbau Principle Rule 1

Determine the number of electrons in an atom of ion, loos at neutral atoms; periodic table atomic number

Wavelength (lambda)

Distance from one wave peak to the next. Units of meters (m) or nanometers (nm). Frequency and wavelength are inversely related.

Electron Affinity Trends

EA generally decreases (less negative) on going down a Group. Electron added to larger orbitals - further from nucleus. Increased electron - electron repulsion. EA generally increase (more negative) on going across period. Smallest for Noble Gasses (8A). Positive values - not favorable. Small for Alkaline Earths (2A). least negative EA. Largest for Halogens (7A). Most negative EA.

EA Trend Exceptions

EA generally decreases on going down a Group. F and O out of line. Small atoms, small volume for electrons, electron-electron repulsion large. Elements with stable electron configurations have small EA. He = 1s2; Be = [He] 2s2; N = [He] 2s2 2p6; Ne = [He] 2s2 2p6

Duality

Electromagnetic radiation can behave both as a wave and as small particles. Light diffracts. Energy delivered in quantized photons.

Valence Electrons

Electrons on the outermost energy level of an atom. Largest n. Most loosely held. Involved in bonding. Determine properties and chemical behavior of element. Similar configuration in same group or family.

Bohr Model of the Atom

Electrons treated as particles only. Electrons located at different energy levels in the atoms. Energy is proportional to the distance from the nucleus. Lowest energy level closest to the nucleus (n=1).

The _________________ corresponds to the principle energy level (n) of the valence electrons

period number

Large increase in energy when we start _________________ core electrons.

removing