CHEM 101 - Midterm Exam #2

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Formation of Ions

• What do METALS typically form and how?

• Metals typically form cations (Na → Na⁺) by losing electrons (radii in pm)

• Metals are generally electropositive

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Formation of Ions

• What do NONMETALS typically form and how?

• Nonmetals typically form anions by gaining electrons (F → F^-)

• Nonmetals are generally electronegative

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Practice Questions

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Stable Ion Configurations for Metals

• How do you form CATIONS from elements?

• To form cations from elements, remove 1 or more e^- from subshell of highest n [or highest [(n + l)].

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Practice Questions

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Stable Ion Configurations for Nonmetals

• How do you form ANIONS from elements?

• To form anions from elements, add 1 or more e^- to subshell of highest n [or highest (n+l)].

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Practice Questions

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Stable Ion Configurations for Transition Metals

• What do you do for transition metals?

• For transition metals, remove n s electrons FIRST and then (n-1) d electrons.

  • This is because the s electrons are farther from the nucleus on average and easier to remove.

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Stable Ion Configurations for Transition Metals

• Why does iron lose the 4s electrons first?

• Larger d for 4s electrons means less negative PE; less E needed to remove an electron

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Effective Nuclear Charge, Z_eff Shielding by Core Electrons

• What is EFFECTIVE NUCLEAR CHARGE (Z_eff)?

• The magnitude of the nuclear charge that is actually felt by the valence electrons

• Z_eff = Z - shielding (~ Z - # core electrons)

  • assume valence electrons (highest n) don’t shield each other

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Exam-Type Question

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Reference: Definitions of Element Radius - Atomic, Metallic, Ionic

• Draw ATOMIC RADIUS.

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Reference: Definitions of Element Radius - Atomic, Metallic, Ionic

• Draw METALLIC RADIUS.

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Reference: Definitions of Element Radius - Atomic, Metallic, Ionic

• Draw IONIC RADIUS.

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Knowledge Check

• Is the following statement true or false? The radius of Ar is expected to be larger than that of Ne because Ar has more protons in its nucleus.

• False! → The radius of Ar is larger than that of Ne because it has a larger valence shell of n = 3 compared to Ne’s valence shell of n = 2.

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Electrostatic Potential Energy (E_el)

• How do you explain an interaction between a valence electron and the nucleus?

• For electrostatic potential energy, Q₁ and Q₂ are charged particles (often opposite in sign)

  • Q⁺ represents the effective nuclear charge (Z_eff)

  • Q^- represents the charge of an electorn (-1)

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) General Size Trends

• What happens down a group?

• valence n increases

• average d from nucleus increases

• electrons held less tightly

• atom size increases

• Z_eff is essentially the same

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) General Size Trends

• What happens across a row/period?

• Z_eff increases

• atom size decreases

• electrons held more tightly

• same n / d from nucleus

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Practice Questions

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Practice Questions

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Ion Sizes: Forming a Cation

• What happens to the ION SIZE when forming a CATION?

• CATIONS are SMALLER than their parent atoms because the electron-electron repulsions have been reduced, and so size DECREASES.

• Especially evident when all electrons from outer shell are lost (as distance between nucleus and electrons decreases, as n decreases).

• Trends in cation sizes down a group are the same as atom sizes.

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Ion Sizes: Forming an Anion

• What happens to the ION SIZE when forming a ANION?

• ANIONS are LARGER than their parent atoms because the electron-electron repulsions increase, and so size INCREASES (even though additional electrons do into same-size subshells).

• Trends in anion sizes down a group are the same as atom sizes.

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Practice Questions

• As the number of protons increase, Z_eff increases, and the radius becomes smaller, decreasing the size of the atom.

• As the number of protons decrease, Z_eff decreases, and the radius becomes larger, increasing the size of the atom.

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Sizes of Ions: Isoelectronic Series

• What does ISOELECTRONIC mean?

• Two or more species have the same number of electorns.

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Sizes of Ions: Isoelectronic Series

• What happens to ION SIZE with an INCREASING Z_eff? Why?

• Ion size decreases with an increasing effective nuclear charge, Z_eff, because electrons are held more tightly and closer to nucleus (and shielding is the same!)

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Exam-Type Question

(Lesson 11: Electron Configurations of Ions; Sizes of Atoms and Ions) Exam-Type Question

(Lesson 12: Ionization Energies; Electron Affinities) Knowledge Check

• Which equation correctly represents the second ionization of aluminum?

• Al⁺ (g) → Al ²⁺ (g) + e^-

(Lesson 12: Ionization Energies; Electron Affinities) Ionization Energy: Definitions

• What is IONIZATION ENERGY?

• The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion.

(Lesson 12: Ionization Energies; Electron Affinities) Ionization Energy: Definitions

• What is FIRST IONIZATION ENERGY (IE₁)?

• First ionization energy (IE₁) is the energy required to remove the first electron from a neutral atom.

(Lesson 12: Ionization Energies; Electron Affinities) Trends in Ionization Energies

• What happens to IE₁ down a column?

• As one goes down a column, less energy is required to remove the first electron, so IE₁ decreases.

(Lesson 12: Ionization Energies; Electron Affinities) Trends in Ionization Energies

• What happens to IE₁ across a row?

• Generally, as one goes across a row, it gets harder to remove an electron, so IE₁ increases.

(Lesson 12: Ionization Energies; Electron Affinities) Practice Questions

(Lesson 12: Ionization Energies; Electron Affinities) (Exceptions to) Trends in First Ionization Energies

• Explain the first exception.

• The first occurs between Groups 2 & 13:

  • Electron removed from p-orbital rather than s-orbital

  • Outermost electron is in slightly higher energy orbital

(Lesson 12: Ionization Energies; Electron Affinities) Practice Questions

(Lesson 12: Ionization Energies; Electron Affinities) (Exceptions to) Trends in First Ionization Energies

• Explain the second exception.

• The second occurs between Groups 15 & 16:

  • Electron removed comes from doubly occupied orbital.

  • Repulsion from other electron in orbital helps in its removal.

(Lesson 12: Ionization Energies; Electron Affinities) Ionization Energy: Successive Electrons

• What happens with successive electrons?

• More energy is required to remove each successive electron.

• When all valence electrons have been removed, the ionization energy is much greater.

(Lesson 12: Ionization Energies; Electron Affinities) Exam Type Question

(Lesson 12: Ionization Energies; Electron Affinities) Electron Affinity: Definition

• What is ELECTRON AFFINITY?

• Electron Affinity is the amount of energy absorbed or released when adding an electron to the ground state of a gaseous atom or ion.

(Lesson 12: Ionization Energies; Electron Affinities) Practice Questions

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Periodic Trends

• Atomic Radius

  • (same period) Increasing Z_eff _________ radius (stronger attraction between core and valence electrons)

  • (same group) Increasing n (principal quantum number) of valence shell _________ radius (electrons are farther away)

  • (same element) Increasing number of valence electrons __________ radius because of increased electron-electron repulsion (e.g., going from F to F^-)

1. decreases

2. increases

3. increases

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Periodic Trends

• IE₁

  • (same period) Increasing Z_eff _________ IE₁ (stronger attraction between core and valence electrons)

  • (same group) Increasing n of valence shell _________ IE₁ (electrons are farther away)

  • (same element) Increasing number of valence electrons __________ IE (e.g., going from F to F^-)

1. increases

2. decreases

3. decreases

(Lesson 12: Ionization Energies; Electron Affinities) Exam-Type Question

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Carbonate

• CO₃^2-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Hydrogen Carbonate / Bicarbonate

• HCO₃^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Acetate

• CH₃COO^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Cyanide

• CN^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Thiocyanate

• SCN^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Hypochlorite

• ClO^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Chlorite

• ClO₂^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Chlorate

• ClO₃^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Perchlorate

• ClO₄^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Chromate

• CrO₄^2-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Dichromate

• Cr₂O₇^2-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Permanganate

• MnO₄^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Azide

• N₃^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Ammonium

• NH₄⁺

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Nitrite

• NO₂^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Nitrate

• NO₃^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Hydroxide

• OH^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Peroxide

• O₂^2-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Phosphate

• PO₄^3-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Hydrogen Phosphate

• HPO₄^2-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Dihydrogen Phosphate

• H₂PO₄^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Disulfide

• S₂^2-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Sulfite

• SO₃^2-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Sulfate

• SO₄^2-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Hydrogen Sulfite / Bisulfite

• HSO₃^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6)

• Hydrogen Sulfate / Bisulfate

• HSO₄^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6) → Oxoanions

• Hypobromite

• BrO^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6) → Oxoanions

• Bromite

• BrO₂^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6) → Oxoanions

• Bromate

• BrO₃^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6) → Oxoanions

• Perbromate

• BrO₄^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6) → Oxoanions

• Hypoiodite

• IO^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6) → Oxoanions

• Iodite

• IO₂^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6) → Oxoanions

• Iodate

• IO₃^-

(Lesson 12: Ionization Energies; Electron Affinities) Reference: Naming Polyatomic Ions (Tables 4.5 and 4.6) → Oxoanions

• Periodate

• IO₄^-

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Ionic Bond vs. Covalent Bond

• What is an IONIC BOND?

• electrostatic attraction between oppositely charged ions (no electron sharing)

• elements are far apart on periodic table, usually between metal and nonmetal

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Ionic Bond vs. Covalent Bond

• What is a COVALENT BOND?

• electrons are shared in overlapping valence orbitals on neutral atoms

• elements are close together on periodic table

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Practice Questions

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas)

• Does breaking bonds absorb energy or release energy?

• absorbs energy

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas)

• Does building bonds absorb energy or release energy?

• release energy

• Describe this graph.

• When the ions are very far apart (right side of graph)…

  • very little electrostatic attraction

  • PE approaches 0

  • ions barely interact

• When the ions move closer together (middle of graph)…

  • opposite charges attract

  • PE decreases (becomes more negative)

  • system becomes more stable

• When the ions are too close together (left side of graph)…

  • the electron clouds overlap excessively

  • strong repulsion occurs

  • PE increases

  • system becomes less stable

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Ionic Bonds: Electrostatic Interaction

• Ionization energy and electron affinity do not adequately describe why energy is released during ionic bond formation, so we need to quantify electrostatic interactions. What is LATTICE ENERGY?

• The amount of energy released when gas-phase ions combine to form one mole of a crystalline solid.

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Practice Questions

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Reference: Lattice Energy

• Lattice energy becomes ____ _______ (more energy is released) as the charge magnitude on the ions increases.

1. more negative

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Reference: Lattice Energy

• Lattice energy becomes ____ ________ (less energy is released) as ion size increases, because the ions are further apart.

1. less negative

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Exam Type Question

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Electronegativity

• What is ELECTRONEGATIVITY?

• X, an atom’s tendency to attract electrons toward itself within a covalent bond

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Electronegativity

• Fluorine (F)

• 4.0

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Electronegativity

• Oxygen (O)

• 3.5

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Electronegativity

• Nitrogen (N)

• 3.0

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Electronegativity

• Chlorine (Cl)

• 3.0

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Electronegativity

• Carbon (C)

• 2.5

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Electronegativity

• Sulfur (S)

• 2.5

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Electronegativity

• Iodine (I)

• 2.5

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Electronegativity

• Hydrogen (H)

• 2.0

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Electronegativity

• Phosphorus (P)

2.0

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Electronegativity

• Bromine (Br)

• 2.8

(Lesson 13: Chemical Bonds; Electronegativity and Bonding; Naming Compounds and Writing Formulas) Electronegativity

• Selenium (Se)

• 2.4