Atomic Structure, Periodicity, and Bonding

Development of a New Atomic Model

  • Rutherford's model couldn't explain heated object colors or element chemical properties.
  • Bohr proposed electrons exist in specific circular paths (orbits).

Bohr's Model of the Atom

  • Ground state: Atom's lowest energy state.
  • Excited state: Atom with higher energy than ground state.
  • Energy is released as electromagnetic radiation (colored light) when an electron:
    • Gains energy and moves to a higher energy level.
    • Drops back to its original orbital.
  • Emitted electromagnetic radiation's energy equals the energy difference between the two orbitals.
  • Elements emit specific line-emission spectra, indicating electrons exist only in specific energy states.

Electromagnetic Radiation

  • Wavelength (λ): Distance between wave peaks, measured in cm, nm, or Å.
  • Frequency (ν): Number of peaks passing a point in a specific time (usually one second).
  • Electromagnetic radiation makes up the electromagnetic spectrum and travels by waves.

Photoelectric Effect

  • The photoelectric effect provides evidence for the particle nature of light.
  • Einstein proposed light travels in energy packets called photons.
  • Light exhibits both wave-like AND particle-like properties.

Max Planck

  • Max Planck related quanta of energy to the frequency of electromagnetic radiation: E = hν
    • Where h = Planck's constant = 6.626 x 10^{-34} Js

Wave-Particle Duality of Electrons

  • Louis de Broglie proposed considering electrons to behave like waves, confined to specific frequencies around the nucleus.
  • This concept works well for particles of very small mass (quantum mechanics).

Heisenberg Uncertainty Principle

  • Looking at electrons requires adding energy, which changes the electron's position.
  • Heisenberg Uncertainty Principle: It's impossible to determine both the position and velocity of an electron.

Schrödinger Equation

  • Schrödinger created equations giving specific energies to describe electron motion.
  • The quantum mechanical model determines allowed electron energies and the probability of finding electrons in various locations around the nucleus.
  • Schrödinger’s equations specify the properties of atomic orbitals and the properties of electrons in that orbital (address of a specific electron).
  • There are 4 quantum numbers in the Quantum Mechanical Model.

Quantum Numbers and Atomic Orbitals

  • Atomic Orbitals: Region of space with a high probability of finding an electron.
    • An orbital can hold only 2 electrons.

Principle Quantum Number (n)

  • Represents the main energy level.
  • As n increases, the atom becomes larger, and the electron is further from the nucleus.
  • The principle quantum number equals the number of sublevels within that principle energy level.

Angular Momentum Quantum Number (l)

  • Indicates the shape of the orbital (sublevel within the energy level).
  • Usually represented by letters: s, p, d, and f.
  • Each energy sublevel corresponds to an orbital of a different shape, describing where the electron is likely to be found.

Magnetic Quantum Number (m)

  • m gives the 3D orientation of each orbital along the x, y, or z axis.

Spin Quantum Number (s)

  • Indicates the spin of the electron, either clockwise (CW) or counterclockwise (CCW).

Electron Configurations

  • The most stable arrangement is achieved by the electrons arranging themselves around the nucleus in specific electron configurations.

Rules for Electron Configurations

  • Aufbau Principle: Electrons occupy the orbitals of lowest energy first.
    • An electron occupies the lowest-energy orbital that can receive it.
    • Beginning in n=3, energies of sublevels start to overlap.
  • Pauli Exclusion Principle: No 2 electrons in the same atom have the same set of all 4 quantum numbers.
  • Hund’s Rule: Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin.

Electron Configurations and the Periodic Table

  • The period on the periodic table corresponds to the principle energy level (n).
  • The block on the periodic table corresponds to the sublevel being filled with electrons.

Exceptions to Electron Configuration Filling

  • Filled sublevels are very stable.
  • The next most stable configuration is one electron in each orbital (half-filled sublevel).

The Periodic Law

  • Chemists used element properties to sort them into groups
  • J.W. Dobereiner in 1829 grouped elements into triads.
  • In 1869, Mendeleev and Meyer published tables of the elements.
  • Mendeleev arranged elements based on repeating properties AND increasing atomic mass.
  • Moseley, after the discovery of protons, arranged elements by atomic number and published his work in 1912.
  • The periodic table arranges elements by atomic numbers, so elements with similar properties fall in the same column or group.
  • Arranging elements by increasing atomic number reveals a periodic repetition of physical and chemical properties.

Element Classification

  • Elements are grouped into three broad classes based on general properties:

Metals

  • Found on the left side of the periodic table.
  • 80% of elements are metals, most solid at room temperature.
  • Conduct heat and electricity, have luster, are ductile and malleable.

Nonmetals

  • Found on the right side of the periodic table.
  • Mostly gases at room temperature, do not conduct heat and electricity (insulators), are dull, brittle.

Metalloids

  • Found touching the stair-step line on the periodic table.
  • Properties are a combination of metals and nonmetals.

Specific Element Groups

  • Metals are broken down into:
    • Alkali metals in Group 1.
    • Alkaline earth metals in Group 2.
    • Transition metals in Groups 3 – 12.
    • Inner transition metals (lanthanides/actinides).
  • Nonmetals are broken down into:
    • Chalcogens in Group 16.
    • Halogens in Group 17.
    • Noble Gases in Group 18.

Hydrogen and Helium

  • Hydrogen (H, 1s^1) doesn't have similar properties to Group 1; it's unique.
  • Helium (He) is in Group 18 but has only 2 electrons (not 8).
  • Helium's outer level is full, therefore He is stable like other Noble Gases.
  • Electrons play a key role in determining the properties of elements.

Noble Gases

  • Noble Gases do not react.
  • They have a full outer energy level (s and p).
  • s and p blocks are called main group elements.
  • d block is called transition metals.
  • f block is called inner transition metals.

Periodic Trends

Factors Affecting Trends

  • Nuclear charge: More protons (p^+) = more pull on electrons (e^-).
  • Shielding effect: Inner electrons block the outer electrons from the pull of the nucleus.

Atomic Radii

  • Atomic radius is defined as one-half the distance between the nuclei of identical atoms bonded together.
  • Atoms tend to be smaller farther to the right across a period due to nuclear charge.
  • Atoms tend to be larger farther down in a group due to shielding.

Ions

  • An ion is an atom or group of bonded atoms with a positive or negative charge.
  • Valence Electrons: Electrons in the outermost energy level.
    • These electrons are lost, gained, or shared to form chemical compounds.
    • Related to the group number: in groups 1 & 2, it equals the group number; in groups 13 through 18, it equals the group # minus 10.

Ion Formation

  • Positive and negative ions form when electrons are transferred between atoms.
  • Metals tend to lose electrons, becoming positive ions (cations).
  • Nonmetals tend to gain electrons, becoming negative ions (anions).

Ionization Energy (IE)

  • Energy required to remove an electron (e^-.)
  • First ionization energy - energy required to remove the first electron (IE_1).
  • Multiple ionization energies - once one electron is removed, there are fewer electrons held by the same number of protons, making it more difficult to remove the next electron.
  • In general, ionization energies of main-group elements increase across each period because of increasing nuclear charge.
  • Among the main-group elements, ionization energies generally decrease down the groups because of increasing shielding.

Ionic Radii

  • Cations (+ ion):
    • Lost electrons but still have the same # of protons attracting electrons that are left
    • Positive ions are smaller than the atoms they come from
  • Anions (- ion):
    • Gained electrons but still have the same # of protons attracting more electrons
    • Negative ions are larger than the atoms they come from

Electron Affinity

  • The energy change that occurs when an electron is acquired by a neutral atom is called the atom’s electron affinity.
  • Electron affinity generally increases across periods due to nuclear charge.
  • Electron affinity generally decreases down groups due to the shielding effect.

Electronegativity

  • Electronegativity: Ability of an atom in a chemical compound to attract electrons from another atom in the compound.
  • Metals tend to lose electrons → low electronegativity.
  • Nonmetals tend to gain electrons → high electronegativity.
  • Across a period, electronegativity tends to increase.
  • Down a group, electronegativity tends to decrease.
  • Fluorine is the most electronegative element and the most likely to attract electrons.

Chemical Bonds and Compounds

Molecular vs Ionic Compounds

  • Chemical bond: The force that holds atoms together, the electrical attraction between the nuclei of one atom and the valence electrons of another atom.
  • Covalent bond: Sharing of electrons to hold two or more atoms together (nonmetals only).
  • Covalent compound: A discrete, neutral group of atoms held together by a covalent bond.
    • Covalent compounds are also referred to as molecules
  • Diatomic molecule: Two atoms held together by a covalent bond.
  • Ionic bonds form when oppositely charged ions attract each other (metal and a nonmetal).
  • Ionic compounds are composed of cations and anions
    • Ionic compounds do not form single units but are referred to as formula units (simplest ratio of ions).
    • Ionic compounds are also referred to as salts.
  • Since ALL compounds are neutral, the charges on the ions must cancel each other out (add up to zero).

Bond Polarity

  • The difference in the electronegativities of two bonded atoms is used to define the type of bond between the atoms.
  • The greater the difference in electronegativities, the more ionic the bond.
  • This occurs along a range and doesn’t have well-defined boundaries.
  • In general, a bond is considered:
    • Nonpolar covalent: if the difference falls between 0 and 0.3
    • Polar covalent: if the difference falls between 0.3 and 1.7
    • Ionic: if the difference falls between 1.7 and 3.3
  • Nonpolar covalent bond: Electrons in the bond are shared equally.
  • Polar covalent bond: Electrons in the bond are shared unequally.
    • More electronegative atoms attract electrons more strongly and gain a slightly negative charge (𝛅−) when bonded.
    • Less electronegative atoms gain a slightly positive charge (𝛅+) when bonded.
  • Ionic bond: Electrons are removed from one atom and transferred to another atom.

Electron Dot Diagrams

  • Write the symbol for the element to represent the nucleus and the inner electrons.
  • Write the electron configuration or use the group number to find the outer level electrons.
  • Each side around the symbol represents an orbital and can hold only 2 electrons.
  • Add dots to represent the electrons.

Covalent Bonding and Molecular Compounds

  • A molecular formula shows how many atoms of each element a molecule contains - H2O2
  • Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has eight electrons in its highest occupied energy level - Octet Rule
    • There are exceptions: molecules in which an atom has fewer or more than a complete octet of electrons.
  • Bond energy: energy required to break the bond between two covalently bonded atoms
    • Large bond energy = strong covalent bond

Multiple Bonds and Structural Formulas

  • Single covalent bond: the sharing of one pair of electrons between two atoms
  • Structural formula: One shared pair of electrons is represented by a dash between the two atoms covalently bonded
  • Unshared pair: Pair of valence electrons not shared between atoms, also called a lone pair.
  • Double covalent bond: The sharing of two pairs of electrons between two atoms
  • Triple covalent bond: The sharing of three pairs of electrons between two atoms
  • Coordinate covalent bond: Both electrons being shared come from the same atom

Example

  • Molecular formula: NH_3
  • Structural formula:
    • H - N - H \ | \ H

Naming and Writing Formulas for Binary Molecular Compounds

  • Write less electronegative atom first (C P N H S I Br Cl O F)
  • Add a prefix only if more than one atom
  • Second atom always has a numerical prefix and -ide suffix
  • “o” or “a” on a prefix is dropped when the element begins with a vowel
  • Use the prefixes in the name to tell you the subscript of each element in the formula
  • Write the correct symbols for the two elements with the appropriate subscripts

Using Chemical Formulas

  • Formula mass: The sum of the average atomic masses of a compound’s elements expressed in atomic mass units (u)
  • Molar mass: Numerically equal to the formula mass of a compound, expressed in grams per mole (g/ mol)
  • Subscripts indicate the number of atoms/ions of each element in a single unit of the compound but also the number of moles of each atom/ ion in one mole of the compound.
    1 molecule C6H{12}O6 \ 1 mole C6H{12}O6

Example Problem: Formula Mass

  • Finding the formula mass of potassium chlorate (KClO_3):
    1 K = 39.10 u \ 1 Cl = 35.45 u \ 3 O = 96.00 u
    Sum: 122.55 u
    *Find the formula mass of barium nitrate

Conversions with Molar Mass

  • Molar mass is used to convert between the mass of a substance and the moles of a substance
  • In the conversion factor, the unit mole always has a value of one (1)
  • 1 mol = molar mass in grams of a substance

Molar Mass of a Gas

  • Avogadro’s hypothesis: equal volumes of gases at the same temperature and pressure contain equal numbers of particles
  • Standard temperature and pressure (STP) = 0℃ and 1 atm
  • At STP, one mole of any gas occupies a volume of 22.4 L
  • 1 mol of a gas = 22.4 L

Mole Conversions

  • Mass to Moles to Particles and Vice Versa

Types of Chemical Reactions

5 Basic Types of Chemical Reactions

  • Combination: 2 or more elements or compounds ➔ one compound
    • A + X ➔ AX
    • metals + oxygen ➔ metal oxides
      • if metal is group 1, get M_2O
      • if metal is group 2, get MO
    • metals + halogen ➔ metal halides
      • if metal is group 1, get MX
      • if metal is group 2, get MX_2
    • active metal (groups 1 & 2) oxides react with water ➔ metal hydroxides
  • Decomposition Reactions: one compound ➔ 2 or more elements or compounds
    • AX ➔ A + X
    • binary compounds break down into the two elements
    • metal carbonates break down into the metal oxide and carbon dioxide
    • metal hydroxides break down into the metal oxide and water
    • metal chlorates break down into the metal chloride and oxygen gas
    • oxyacids break down into water and the nonmetal oxide
  • Single Displacement (Replacement): element + compound ➔ new element + new compound
    • A + BX ➔ B + AX
    • The activity series is used to predict if a reaction will actually occur
    • More reactive metals replace less reactive metals
    • More reactive nonmetals replace less reactive nonmetals
    • It also helps sometimes to view water as H^+ and OH^-
      • the most common product of H^+ and OH^- is water
  • Double Displacement (Replacement): compound + compound ➔ new compound + new compound
    • AX + BY ➔ AY + BX
    • Ions of two compounds exchange places in an aqueous solution to form two new compounds
    • For the reactions to actually occur, one of the compounds formed is either a precipitate, an insoluble gas that bubbles out of the solution, or a molecular compound (usually water)
    • Precipitates occur because the substance created will not dissolve
    • The formation of a precipitate can be predicted by using the general rules for solubility
  • Combustion: Substance combines with oxygen
    • Usually seen as: hydrocarbon + O_2 ➔ carbon dioxide and water

Describing Chemical Reactions

  • A chemical reaction is the process by which one or more substances are changed into one or more different substances.
  • A chemical equation has formulas of reactants and products with symbols for what is occurring
    • Original substances = reactants
    • Resulting substances = products

Indications of Chemical Reactions

  • Production of heat AND light
  • Production of a gas (bubbles)
  • Formation of a precipitate
  • Change in color
  • Change in odor

Balancing Chemical Equations

  • The equation must represent known facts.
  • The equation must contain the correct formulas for the reactants and products.
    • Remember the elements that exist as diatomic molecules.
  • The law of conservation of mass must be satisfied.
  • Word equation: names of reactants and products with symbols for what is occurring
  • Formula (skeleton) equation: chemical equation that does not indicate the relative amounts of the reactants and products
  • Coefficients of a chemical reaction indicate relative, not absolute, amounts of reactants and products.
  • This ratio shows the smallest possible relative amounts of the reaction’s reactants and products as either a particle ratio or mole ratio.
    • 2H2 + O2 ->2H_2O
    • 2 molecules H : 1 molecule O : 2 molecules H_2O
    • 2 moles H : 1 mole O : 2 moles H_2O
  • An equation gives no indication of whether a reaction will actually occur.

Balancing Principles

  • Must meet the Law of Conservation
    • Can NOT create or destroy matter
  • May NOT change subscripts
  • May ONLY change the coefficient
  • Balanced equation: must have the same elements and the same number of atoms of each element on both sides of the equation
  • You may balance polyatomic ions that appear on both sides as a single unit
  • Balance atoms that appear more than once on each side last (usually H & O)
  • Make sure the coefficients are in the simplest ratio

Molecular Geometry

  • The structure of a particular molecule is important in determining how it functions in a chemical reaction, in a solution, in a biological cell, or in other areas of nature.
  • Scientists use a simple model called valence shell electron pair repulsion model to predict the geometry of molecules formed by nonmetals.
  • All electrons repel each other - this will determine the shape of molecules.
  • Unshared pairs of electrons have more energy to repel than shared pairs of electrons.
  • Bonding and nonbonding pairs around a given atom will be positioned as far apart as possible.
  • Double bonds and triple bonds are counted as if they are one bond in the VESPR model.

Common Shapes

  • LINEAR: 1 shared pair of electrons, 0 to 3 unshared pairs of electrons
    • AB structure
    • 2 atoms (2 points) are linear (form a line) in shape no matter how many unshared pairs of electrons are present
  • LINEAR: 2 shared pairs of electrons, 0 unshared pairs of electrons
    • AB_2 structure
    • 180º angle
  • 3 shared pairs of electrons, 0 unshared pairs of electrons
    • AB_3 structure
    • 120º angle
  • Tetrahedral: 4 shared pairs of electrons, 0 unshared pairs of electrons
    • AB_4 structure
    • 109.5º angle
  • 3 shared pairs of electrons, 1 unshared pair of electrons
    • AB_3E structure
    • 107º angle
  • Bent: 2 shared pairs of electrons, 2 unshared pairs of electrons
    • AB2E2 structure
    • 104.5º angle

Polar Molecules

  • One end of a polar molecule is slightly negative, while the other end is slightly positive
  • Polar molecules are referred to as dipoles
  • Polar bonds do not necessarily make a molecule polar
  • If the shape of a molecule is symmetrical, the polarity of each bond has a tendency to cancel each other out

Intermolecular Forces

  • Intermolecular attractions are weaker than either ionic or covalent bonds
  • These attractions are responsible for determining whether a molecular compound is a gas, a liquid, or a solid at a given temperature
  • London dispersion forces: Occur between non-polar molecules by the motion of their electrons (also called induced dipole)
  • Dipole interactions: Dipoles are attracted to one another by their slightly opposite charges
  • Hydrogen Bonding: When hydrogen is bonded to a highly electronegative atom, it creates a uniquely strong dipole