Chemistry Q1 Pt. 2

Multi-Electron Atoms

- Energies of orbitals depend on n and l, no degenerate states
- Treat multi-electron state as a sum of single-electron states

Shielding

Inner electrons screen/shield outer electrons from the full Coulombic effect of the nucleus, increasing the energy of the orbital where the electron is found

Penetration of the subshells

s > p > d > f

Relative Orbital Energies

- As l increases, the energy of the subshell increases
- All orbitals with the same value of l have the same energy

Electron Configuration

- Electrons occupy lowest orbitals first (Aufbau Principle)
- Each orbital can hold a max of 2 electrons w/ opposite spin (Pauli Exclusion Principle)
- If two or more orbitals have the same energy, electrons will occupy orbitals singly before pairing to reduce electron-electron repulsion and lower atom's energy (Hund's Rule)

Two Types of Representations

- spdf notation: write shell/subshell in increasing energy, # of electrons in superscript (e. 1s^2 2s^2)
- Orbital Diagrams

Electron Configurations - Period 1

- Hydrogen, Z = 1, 1s^1
- Helium, Z = 2, Fill the first shell to create a CLOSED SHELL, 1s^2

Electron Configurations - Period 2

- Lithium, Z = 3, [He]2s^1
And so on and so forth

Noble Gases in Notation

- Chemically inert closed shells that are full

Electron Configuration - Period 4

- 4s subshell is lower in energy than the 3d subshell so it fills first
- 4s < 3d < 4p < 5s < 4d < 5p

The (n + l) Rule

The ordering of the energies of the orbitals in a multi-electron atom increases with the value of n + 1

Valence Shell

Outermost, highest energy shell, equal to the row # of the element

Valence Electrons

- Determine element's physical/chemical properties
- # of valence electrons = last digit of group #

Excited State

Atom absorbs energy, promoting an electron to a higher energy orbital
- Ground state of C: [He]2s^(2) 2p^(2)
Excited state of C: [He]2s^(2) 2p^(3)

Effective Nuclear Charge (Zeff)

- Net positive charge attracting an electron in an atom
- Zeff is always less than Z in a multi-electron atom because of the shielding effect

Zeff Equation

Zeff = Z (actual nuclear charge) - S (core electrons)

Across a period, Zeff...

Increases as electrons join the same valence shell

Down a group, Zeff...

Decreases

Atomic Radius

- Estimated as half the distance (d/2) between the center of neighboring nuclei
- 2 metal atoms: Metallic radius
- 2 nonmetal atoms: Covalent radius
- 2 gaseous atoms: Van der Waals radius

Ionic Radius

Effective radius of an ion in an ionic solid

Across a period, atomic radii...

Decreases as more electrons are attracted to more protons int he nucleus

Down a group, atomic radii...

Increases

Cation Size

Always SMALLER than the atoms they form from

Isoelectronic Cations

The more positive the ionic charge, the smaller the cation

Anion Size

Always LARGER than the atoms they form from

Isoelectronic Anions

The more negative the ionic charge, the larger the anion

Ionization Energy (IE)

- Minimum energy required to remove one electron from a single atom in the gaseous state
- The larger the IE, the more stable the electronic configuration of the atom

Across a period, IE...

Increases due to increased force of attraction between nucleus and electrons

Down a group, IE...

Decreases due to an increasing n

IE1 < IE2 < IE3

Atoms are less stable after electron loss, more energy is required to remove more electrons for higher IEs

Electron Affinity (EA)

- Energy released on adding one electron to an atom in the gaseous state
- The larger the EA, the more stable the anion that is formed

Across a period, EA...

Increases

Down a group, EA...

Decreases

Electron Affinity (EA1 vs. EA2)

ANS: Adding a second electron (EA2) always requires energy input because of the electrostatic repulsions so EA2 will be negative

Down a group, reactivity...

increases as shielding lowers nuclear pull on electrons

Across a period, reactivity...

decreases

Down a group, lattice energy...

Decreases for molecules

Diagonal Relationship in Main Group

- Some pairs of elements diagonally-adjacent to each other have similar chemical properties
- Atomic radii decreases from left to right, increase top to bottom
- IE increases from left to right but decreases from top to bottom

Diamagnetism

- Atoms/ions which possess only spin-paired electrons
- Magnetic effects cancel out and they are weakly repelled by magnetic field

Paramagnetism

- Atoms/ions which possess one or more orbitally unpaired electrons
- Attraction to magnets

Chemical Bond

Result of attraction between two or more atoms/ions that involves sharing or transferring electrons

Chemical Bond Determinants

Ionization Energy and Electron Affinity: How readily an atom gives up/accepts an electron

Why do they form?

- Reduction in the Total Energy
- Lowering of the potential energy (Potential Energy of bonded atoms must be lower than that of non-bonded atoms)

Maximum Stability for an Atom

- Atom is isoelectronic with a noble gas atom, reaches a stable electron configuration

Lewis Dot Structure

- Dots (denote valence electrons) are distributed around the four sides of an element symbol, first singularly and then in pairs
- Combine Lewis structures to represent transfer/sharing of electrons in an ionic or covalent bond
- Most atoms gain/lose atoms to complete octet or duplet

Chemical Formula

- Composition: What numbers/kinds of atoms present?
- Constitution: How are the atoms connected?
- Configuration: How are the atoms positioned in 3-D space?

Empirical Formula

Shows relative # and kinds of atoms present
Ex. CH2O

Molecular Formula

Shows actual # and kinds of atoms present
Ex. C2H4O2

Structural Formula

Shows the atom connectivity with #/kinds of atoms present

Ionic Bonding

- Occurs between ions (a cation and anion)
- Cation is a metal, Anion is a non-metal
- Compounds represented by empirical formula ONLY, termed "formula unit"

Step 1 of Ionic Bonding

- Form Ions
- Metals tend to lose electrons, Non-metals tend to gain them
- Main Group Elements (Groups 14/15) don't have a strong preference for either gaining or losing electrons
- Transition Metal Cations firs lose valence s electrons, may lose d electrons (2+ is common)
- Cations have electron configuration of preceding Noble Gas
- Anions have electron configuration of next Noble Gas

Inert-Pair Effect in Lower P-Block Elements

- Cation formed by loss of ONLY the p electrons
- p-Block of periods 4, 5, or 6 formt wo types of cations depending on if they lose only the p subshell or both p and s subshell

Step 2 of Ionic Bonding

- Combine Ions
- Ionic Bond is complete transfer of one or more electrons, held together by electrostatic forces
- Ionic Compound has lowest ratio of ions

Chemical Formula of Ionic Compounds

- Cations listed first,t hen anions second
- Eliminate common factor in subscripts
- Simplest ratio of anions and cations equal zero charge

Formula Mass (amu)

Sum of atomic mass of each atom x # of atoms of that element int he compound

How do we name Ions and Ionic Compounds?

- Depends on if the ions are monoatomic or polyatomic

Metal cations (Main Groups 1, 2, and 13)

Identify metal and add "ion" to the end
Ex. K+ = Potassium Ion

Transition Metals/Metals (Groups 13-15)

Use Roman numerals to indicate ion's charge
Ex. Mercury = Hg2^2+ = Mercury (I)
Mercury = Hg2+ = Mercury (II)

Anions

End with "ide" and "ion"
Ex. Bromide = Br- = Bromide ion

Polyatomic Ions

Ions composed of more than one atom (covalently bonded)
- Positive Molecular Ions: End in "ium" or "onium"
Ex. NH4^1+ = Ammonium
- Negative Molecular Ions: End in "ide", "ite", or "ate"
Ex. ClO^1- = Chloride ClO2^1- = Chlorite ClO3^1- = Chlorate

3-D Lattices of Ions

- Ion sizes/charges determine crystalline structures

Ionic Bonding: The Energetics

- Ion-Pair of from gas phase
- Step 1: Ionization of Sodium (loses electron); Ionization Energy is given
- Step 2: Electron Gain by Chlorine; Electron Affinity is given
- Step 3: Formation of the ion-pair into NaCl; Apply Coulomb's Law by multiplying charges of two ions and dividing by distance (radii of ions given, add together to get total distance)
Net Charge = E = (IE - EA) - NaCl Energy from Coulomb's law

Born-Haber Cycle

Features a series of (hypothetical) steps starting from the elements in their standard states, used to calculate Lattice Energies

Lattice Energy Trends

The greatest stabilization in the solid occurs when ions are small and highly charged

How Can We Explain the Conductivity of Salt Solutions?

Solid = No conductivity
Molten = Some conductivity
Dissolved in water = Lots of conductivity

What Import do Ionic Compounds Have in Our Lives?

Real-world uses and examples of Ionic Compounds

Stalagmites/Stalactites

Calcium carbonate/Magnesium Carbonate

Bones and Cement Matrices

Calcium Phosphate used to replace damaged/diseased bone

Fertilizers

Ammonium Nitrate

Cleaners/Disinfectants

Solution of Sodium Hypochlorite

Baking and Food Additives

Sodium Bicarbonate (creates CO2 gas in baking)
Sodium Nitrite (food additive)

Multivitamin Supplements

Magnesium oxide, Potassium Chloride, Calcium Carbonate, Ferrous Fumarate, Calcium Stearate, Manganese Sulfate, Nickelous Sulfate, Potassium Iodide, Sodium Ascorbate, Sodium Benzoate, Sodium Citrate, Titanium Dioxide, Zinc Oxide

Covalent Bonding

- Create molecules, electrically neutral groups held together by covalent bonds
- Occurs between non-metals/metalloids

Lewis Structure of Covalent Bond

- Sharing electrons between atoms
- Two dots or a line represent a single bond, or shared pair of electrons

How do Covalent Bonds Form?

- As two atoms with incomplete valence shells interact, their electron clouds overlap and concentrates the electron density, forming a covalent bond
- If atoms are too close, strong repulsions occur
- If atoms are too far apart, attractions are weak and no bonding occurs
- When atoms are optimally separated, the energy is at a minimum and the distance between the nuclei is the bond length

Diatomic Elements: Period 2

- OCTET RULE applies for C, N, O and F
- Multiple Bonds, or more than one shared pair of electrons, can form
Ex. Single, Double or Triple Covalent Bonds
- Bond Order (B.O.) indicates # of shared electron pairs between atoms

Chemical Formulae: Covalent Compounds

- Element of lowest IE is written first
- Write names with numerical prefixes (don't add mono- to first element if there is only one present in compound)
- Ex. N2O5 = Dinitrogen Pentoxide

Numerical Prefixes (1-10)

Mono
Di
Tri
Tetra
Penta
Hexa
Hepta
Octa
Nona
Deca

Formal Charge

FC = # of valence electrons - # of bonds made - # of unshared electrons

Formal Charge Rules

- Sum of Formal Charges should equal the overall charge of the atom/ion
- Negative Formal Charges go to atoms with highest IE
- Positive Formal Charges go to atoms with lowest IE

Resonance: Equivalent Structures

- When certain molecules/ions have 2 or more valid but different Lewis structures

Resonance Hybrid

Combines structures to make a more accurate description of the real molecular structure

Bond Order

B.O. = (Add up bond types)/# of bonds made
Ex. O=O-O (2 + 1)/2 = 1.5

Resonance: Inequivalent Structures

- When there are several (inequivalent) resonance structures that do no contribute equally to resonance hybrid
- Sometimes one structure is the main contributor
- Consider if there is a single bond between central and terminal atoms
- Consider having low formal charges, positive FC on low IE and negative FC on high IE

Odd-Electron Molecules (Radicals)

- If a molecule has an odd # of electrons, it's not possible to write a Lewis structure in which there is an octet of electrons about each atom
- Valid Lewis structures have: Maximized the # of atoms which possess an octet and place the odd electron on the atom w/ lowest IE (minimize formal charges)

Hypervalent (Valence Shell Expanded) Molecules

- Some atoms in the 3rd period and beyond appear to accommodate more than 4 pairs of electrons, suggesting expanded valence shells
- Valid Lewis Structures first assign lone pairs to the outer atoms to give them octets, then assigning remaining valence electrons to the central atom as lone pairs

Hypovalent (Octet-Deficient) Molecules

- If molecule possess less than the # of required valence electrons (electron deficient), it won't be possible to write a Lewis structure in which there is an octet of electrons about each atom
- Valid Lewis Structures assign octets to the atoms with highest IE, then incomplete octet occurs on the least electronegative atom

Properties of the Covalent Bond: Bond Strength

- Measured by Dissociation Energy (D), kJ mol^-1
- Energy required to separate gaseous bonded atoms; deeper the potential energy well, the stronger the bond

Properties of the Covalent Bond: Bond Length

- Distance between atomic nuclei in a given covalent bond
- Measured experimentally; as bond length decreases, bond strength increases

Trends: Bond Orders

When more electrons are shared between two nuclei, bond length decreases and bond strength increases

Non-Polar Covalent Bond

- Typically connect identical atoms, electrons are shared equally

Polar Covalent Bond

- Connect atoms of different elements
- Nucleus of one atom attracts electrons more strongly than the nucleus of the other atom, acquiring a partial negative charge and leaving the other atom with a partial positive charge
- Electric Dipole Moment: Arrow points from - to +
- Calculated by Q (charge) x r (distance m)

Correcting the Covalent Model: Electronegativity

- Measures the ability/tendency of an atom to attract electrons from another atom to which it is bonded to
- The element with the higher E.N. value will attract electrons more strongly in the bond (greater electron-pulling power)

Electronegativity: Mulliken Definition

- Use Ionization Energy and Electron Affinity provide indication of how readily an atom may give up or accept an electron
- Higher the IE, the greater the value of E.N.
- E.N. = 1/2(IE + EA)

Electronegativity: Pauling Definition

- Created an electronegativity scale based on dissociation energies, D in eV, of the homonuclear and heteronuclear bonds
- Described the difference in electronegativities between two elements
- Excess bond energy is a result of the ionic component of the bond caused by partial charges

Electronegativity Differences as a Measure

- Polarity of a covalent bond is directly proportional to the difference in electronegativity of the bonded atoms
- As E.N. increases, so do the partial charges on each atom, resulting in a bond with a greater ionic character

Covalent E.N. Range

0 - 0.3

Polar Covalent E.N. Range

0.4 - 2.0

Ionic E.N. Range

2.1+

3 Different Ways of Depicting Molecular Shape

- Dashed-Wedged Line Formula
- Ball-and-Stick Model
- Space-Filling Model