Ch. 9: Electrons in Atoms and Ch. 9: Electrons in Atoms and the Periodic Tablethe Periodic Table

Dr. Namphol Sinkaset

Chem 152: Introduction to General Chemistry

I. Chapter OutlineI. Chapter Outline

I. Introduction

II. Electromagnetic Radiation

III. The Bohr Model of the Atom

IV. The Quantum-Mechanical Atom

V. Electrons and the Periodic Table

VI. Periodic Trends

I. Hydrogen vs. HeliumI. Hydrogen vs. Helium

I. Electrons and ChemistryI. Electrons and Chemistry

• Chemistry is all about electrons.

• Therefore, how electrons are organized in the atom is an important concept.

• We will see that reactivity and the arrangement of the periodic table are both related to electrons.

I. Stranger Than Anyone I. Stranger Than Anyone ThoughtThought

• Theories used to explain how electrons are organized in the atom were devised by scientists like Bohr, Schrödinger, Planck, and Einstein.

• “I don’t like it, and I am sorry I ever had anything to do with it.”

• “God does not play dice with the universe.”

II. It Started With LightII. It Started With Light

• Light can interact with atoms, and studying the interactions led to understanding how electrons are organized in the atom.

• Light is not a form of matter; it is electromagnetic radiation.

• EM radiation is a type of energy that travels at 3.0 x 108 m/s.

II. Classical WaveII. Classical Wave

• Light was considered a classical wave phenomenon – like ripples in water or a moving rope.

• But that wasn’t quite right…

II. Describing WavesII. Describing Waves

• EM waves are characterized by their wavelength (λ) and their frequency (ν).

II. White LightII. White Light

• White light can be separated.

• Note that short wavelength = high frequency and vice versa.

II. Light As ParticlesII. Light As Particles

• In the early 20th century, scientists like Einstein saw that light was not a classical wave.

• They explained that light acted like particles, which were called photons.

• Thus, when we are under a light, we are being showered with light particles.

II. The EM SpectrumII. The EM Spectrum

III. Atoms Can Emit LightIII. Atoms Can Emit Light

III. Atomic Emission SpectraIII. Atomic Emission Spectra

III. Atomic LINE SpectraIII. Atomic LINE Spectra

• The individual lines were key to formulation of Bohr’s atomic model.

• The movement of electrons were the reason for atoms emitting light.

• But why lines? Lines meant only specific wavelengths (colors) were allowed.

III. The Bohr ModelIII. The Bohr Model

• Bohr reasoned that electrons were only allowed to have certain energies.

• Strange – it’s like someone telling you were only allowed to have certain amounts of money.

III. Bohr EnergiesIII. Bohr Energies

• The electron orbits in the Bohr model are like rungs on a ladder.

• You can stand on one step or another, but never in between.

III. Bohr OrbitsIII. Bohr Orbits

• Each orbit in the Bohr model has a specific energy that is specified by a quantum number, n.

• When an electron moves to a higher orbit, it must absorb a quantum of energy.

• When an electron moves to a lower orbit, it must emit a quantum of energy.

III. Moving Between OrbitsIII. Moving Between Orbits

III. Line Spectra Correspond III. Line Spectra Correspond to Electron Transitionsto Electron Transitions

III. Works Great for HydrogenIII. Works Great for Hydrogen

• Originally, Bohr set out to model only the hydrogen atom.

• When people tried to extend it atoms with more than one electron, it didn’t work!

• A new model that worked for all atoms was needed.

IV. Wavy ElectronsIV. Wavy Electrons

• Experiments found that electrons don’t always act like particles – they sometimes act like waves!

• Electrons phase in and out (similar to how waves oscillate), so we don’t know exactly where they are.

• The best we can do is plot probability maps.

IV. Baseballs vs. ElectronsIV. Baseballs vs. Electrons

IV. Orbits to OrbitalsIV. Orbits to Orbitals

• In the Bohr model, electrons were in well-defined orbits like planets around the sun.

• In the new quantum-mechanical model, orbits are replaced by orbitals, which are probability maps of where an electron could be found.

IV. OrbitalsIV. Orbitals

• In the Bohr model, each orbit was labeled with a single quantum number.

• For orbitals, it’s more complex, so we need something else.

• Orbitals are labeled with a principal quantum number (n)and a subshell letter designation (s, p, d, f).

IV. Principal Quantum NumberIV. Principal Quantum Number

• The principal quantum number specifies the principal shell of the orbital.

• Higher n means higher energy.

IV. SubshellsIV. Subshells

• Each principal shell has one or more subshells.

• Each subshell has a different “shape.”

IV. The IV. The ss Orbitals Orbitals

• s orbitals are spherical.

• The 1s orbital is the lowest possible energy for an electron; it is the ground state.

IV. IV. 1s1s vs. vs. 2s2s

• A 2s orbital is bigger and has more energy than a 1s orbital.

• If a 1s electron in hydrogen transitions to 2s, then the hydrogen atom is now in an excited state.

IV. The IV. The pp Orbitals Orbitals

• There are three p orbitals, each with a different orientation.

IV. The IV. The dd Orbitals Orbitals

• There are five d orbitals.

IV. The IV. The ff Orbitals Orbitals

• There are seven f orbitals.

IV. Energy Order of OrbitalsIV. Energy Order of Orbitals

IV. Electrons and OrbitalsIV. Electrons and Orbitals

• Electrons have an intrinsic spin property. They can spin up or down.

• According to the Pauli exclusion principle, only two electrons with opposite spin can “occupy” an orbital.

• We use electron configurations or orbital diagrams to show electrons in atoms.

IV. The Hydrogen AtomIV. The Hydrogen Atom

IV. Energy Order of OrbitalsIV. Energy Order of Orbitals

• Order of the orbitals can be obtained from the periodic table.

• But, if you don’t have one, you can remember the orbitals with a simple diagram.

IV. Hund’s RuleIV. Hund’s Rule

• We know there are only two electrons w/ opposite spin per orbital.

• What about when more than one of the same type is available (for > s)?

IV. Sample ProblemIV. Sample Problem

• Write electron configurations and orbital diagrams for the following. Write condensed forms for the last two. Al Mn Sr Br

V. Blocks on the Periodic TableV. Blocks on the Periodic Table

• We can organize the Periodic Table into blocks in which s, p, d, or f orbitals are being filled.

• This allows us to easily write electron configurations or orbital diagrams based on an element’s location.

• Note there are 2 columns for s-block, 6 columns for p-block, 10 columns for d-block, and 14 columns for f-block. Why?

V. e- Config Based on LocationV. e- Config Based on Location

V. Using the Periodic TableV. Using the Periodic Table

V. Two Types of ElectronsV. Two Types of Electrons

• Core electrons are those that are not in the outermost principal shell.

• Valence electrons are those in the outermost principal shell.

V. Valence ElectronsV. Valence Electrons

• The reactivity of an atom is determined by its valence electrons.

• The valence electrons are loosely held by the atom because they are the furthest away from the nucleus.

• Thus, they can be easily lost or gained, which leads to chemical reactions/properties.

V. Valence Electrons and V. Valence Electrons and Element FamiliesElement Families

• Elements in same family have similar reactivity because they have the same valence electron configuration.

V. The Periodic TableV. The Periodic Table

V. Why Group V. Why Group 17’s Form 1- 17’s Form 1-

AnionsAnions

V. Why Group 1’s V. Why Group 1’s Form 1+ CationsForm 1+ Cations

VI. Periodic TrendsVI. Periodic Trends

• The quantum mechanical model of the atom allows prediction of some periodic trends.

• We will examine the trends of atomic size, ionization energy, and metallic character.

VI. Shells and SubshellsVI. Shells and Subshells

• Shells are like layers of groups of electrons around the nucleus. More shells = larger size.

• Subshells rest inside a shell and don’t add any thickness to the shell.

• As we go across a period, we add electrons AND protons.

• This knowledge will help us understand the trends.

VI. Main Group Atomic SizeVI. Main Group Atomic Size

VI. Explaining the TrendVI. Explaining the Trend

• Down a family is easy: opening more shells, so atomic size must increase.

• Across a period, the principal shell stays the same; electrons are just filling subshells. However, each additional proton pulls

everything in closer.

VI. Ionization EnergyVI. Ionization Energy

• Ionization energy is the energy needed to take away an electron from an atom in the gas phase. Na + ionization energy Na+ + 1e-

• The ionization energy follows a clear trend.

VI. Ionization Energy TrendVI. Ionization Energy Trend

VI. Explaining the TrendVI. Explaining the Trend

• The trend can be correlated with atomic size.

• An electron is easier to remove if it is further away from the nucleus.

• Thus, larger atoms will have LOWER ionization energies than smaller atoms.

VI. Metallic CharacterVI. Metallic Character• One characteristic of metals is that they

tend to lose electrons.

• If we use this as a criteria for metallic character, then atoms with low ionization energies are more metallic than those with high ionization energies.

• Thus, the trend in metallic character is the opposite of the trend in ionization energy.

VI. Metallic Character TrendVI. Metallic Character Trend

VI. Sample ProblemVI. Sample Problem

• Choose the appropriate atom in each pair. Larger atom: Pb or Po Larger atom: Cl or Se Higher ionization energy: Mg or Sr Higher ionization energy: Cu or P More metallic: Au or Cu More metallic: N or S